SOLUBILITY OF BASES AND SALTS IN WATER AT 18°.
K
Na
Li
Ag
Tl
Ba
Sr
Ca
Mg
Zn
Pb
01
32.95
3.9
35.86
5.42
77.79
13.3
0.0316
0.0410
0.3
0.013
37.24
1.7
51.09
3.0
73.19
5.4
55.81
5.1
203.9
9.2
1.49
0.05
Br
65.86
4.6
88.76
6.9
168.7
12.6
0.041
0.066
0.04
0.0215
103.6
2.9
96.52
3.4
143.3
5.2
103.1
4.6
478.2
9.8
0.598
0.02
I
137.5
6.0
177.9
8.1
161.5
8.5
0.063
0.071
0.006
0.0317
201.4
3.8
169.2
3.9
200
4.8
148.2
4.1
419
6.9
0.08 1
0.0,2
F
92.56
12.4
4.44
1.06
0.27
0.11
195.4
13.5
72.05
3
0.16
0.0292
0.012
0.001
00016
0-0,2
0.0076
0.0214
0.005
0.035
0.07
0.003
NOS
30.34
2.6
83.97
7.4
71.43
7.3
213.4
8.4
8.91
0.35
8.74
0.33
66.27
2.7
121.8
5.2
74.31
4.0
117.8
4.7
51.66
1.4 .
CIO.,
6.6
0.52
97.16
6.4
313.4
15.3
12.25
0.6
3.69
0.13
35.42
1.1
174.9
4.6
179.3
5.3
126.4
4.7
183.9
5.3
150.6
3.16
Br03
6.38
0.38
36.67
2.2
152.5
8.20
0.59
0.025
0.30
0.009
0.8
0.02
30.0
0.9
85.17
2.3
42.86
1.5
58.43
1.8
1.3 (
0.03 «
10,
7.62
0.35
8.33
0.4
80.43
3.84
0.004
0.0314
0.059
0.0216
0.05
0.001
0.25
0.0.j57
0.25
0.007
6.87
0.26
0.83
0.02
0.002
0.043
OH
142.9
18
116.4
21.
12.04
5.0
0.01
0.001
40.04
1.76
3.7
0.22
0.77
0.063
0.17
0.02
0.001
0.0.,2
0.0,5
0.045
0.01
0.0.,4
SO4
11.11
0.62
16.83
1.15
35.64
2.8
0.55
0.020
4.74
0.09
0.0,23
0.0410
0.011
0.036
0.20
0.015
35.43
2.8
53.12
3.1
0.0041
0.0313
Cr04
63.1
2.7
61.21
3.30
111.6
6.5
0.0025
0.0.,15
0.006
0.031
0.0,38
0.0415
0.12
0.006
0.4
0.03
73.0
4.3
0.042
0.0C5
C204
30.27
1.6
3.34
0.24
7.22
0.69
0.0035
0.032
1.48
0.030
0.0086
0.0338
0.0046
0.0326
0.0,56
0.0443
0.03
0.0027
0.0,6
0.044
0.0,15
0.055
C03
108.0
5.9
19.39
1.8
1.3
0.17
0.003
0.031
4.95
0.10
0.0023
0.0311
0.0011
0.047
0.0013
0.0313
0.1
0.01
0.004?
0.033?
0.0,1
0.043
The upper number in each square gives the number of grams of the
anhydrous salt held in solution by 100 c.c. of water. The loiver number if
the molar solubility, i.e., the number of moles contained in one liter of the
saturated solution. For some other solubilities, see pages 157 and 644.
tbe Same Hutbor
Second Edition. Hevised.
With Eighteen Figures. Interleaved.
Price 75 cents.
CHICAGO
THE UNIVERSITY OF CHICAGO
PRESS
c
5
INTRODUCTION TO
ALEXANDER SMITH
PROFESSOR OF CHEMISTRY AND DIRECTOR OF GENERAL AND PHYSICAL CHEMISTRY
IN THE UNIVERSITY OF CHICAGO
507021
NEW YORK
THE CENTURY CO.
1906
COPTSIGHT, 1905, 1906,
BY
THE CENTURY CO.
Stanbope press
F.H.GILSON COMPANY
BOSTON. U.B.A.
PREFACE
THIS book, the first draft of which was written six years ago, is the
outgrowth of the introductory course in chemistry which the author
has given for the past fifteen years. A subject undergoing the per-
sistent, though unconscious criticism of keen minds should gain in
self-consistency and coherence as it is presented year after year. For
example, an answer must be found for the common question, " Why
does the chemistry of the laboratory differ from the chemistry of the
text-book and the lecture to such an extent that they seem to be different
sciences ? " The chemistry of the laboratory is, of course, the only
real chemistry, and that of the lecture must be somewhere at fault.
The student neither sees nor weighs atoms, for instance, and so the
details of the laboratory experiment, which are seen and studied, become
the basis of the whole treatment. The atom and the ion assume the
role of merely figurative aids in the description of the facts. Gradually
the conception of chemical equilibrium comes to contribute the major
part of the explanation which is essential to the evolution of a system
of chemistry founded upon experiment.
In the choice and arrangement of the material, several principles
have served as guides :
The book is intended primarily for students beginning the study of
chemistry in a college, university, or professional school. It is assumed
that use of the book goes hand in hand with systematically arranged
laboratory work in general chemistry. The first four chapters, for ex-
ample, contain a discussion of a few typical experiments. They appeal
directly to experience derived from the performance and observation
of these and other similar experiments in the laboratory and in the
class-room. In these chapters some of the features which are charac-
teristic of every chemical phenomenon are sought out, put into words,
and illustrated.
vi PREFACE
No conception is defined, and no generalization or law is developed,
until such a point has been reached that applications of the conception
and experimental illustrations, later to be related in the law, have
already been encountered, and there is about to be occasion for further
applications and illustrations of the same things in the chapters
immediately succeeding. In these chapters the applications are
frequent and explicit. Later, page references in parentheses con-
tinue to indicate the recurrence of examples which might other-
wise fail to be noticed. It is one thing to come to know a principle
of the science, and quite another thing to have acquired, by constant
repetition of the process, a confirmed habit of using the prin-
ciple on every appropriate occasion. To assist still further in the
attainment of this end, an attempt is made in the last six chapters
again to mention and illustrate, by way of review, the most important
principles of the science.
No conception or principle is given at all, unless, in its most elemen-
tary aspects, it can be made clear to a beginner ; and unless it is
capable of numerous applications in elementary work; and, finally,
unless a knowledge of it is of material use in organizing and unifying
the result of such elementary work.
An attempt has been made to state the laws and to define the con-
ceptions of the science in terms of experimental facts. The figurative
language of hypothesis has been employed only in explanations.
Familiarity with physical conceptions and facts is so indispensable
to the chemist that no apology is needed for the rather full treatment
which some of them have received.
No two chemists would agree perfectly in regard to the apportion-
ment of space. The processes of chemical industry, and the every -day
applications of chemical science, cannot all be mentioned. These
fields, and that of mineralogy, can be represented by examples, with-
out the incompleteness of the result being in any way a detriment to a
work of a general character. Again, a dense array of descriptive
material, unillumined by explanation, is a positive injury to an intro-
ductory treatise. All reference to historical matters cannot be omitted,
but a logical display of the subject can be achieved with comparatively
PREFACE Vli
little of the history. Of all the aspects of the science, the theoretical
is thus the one whose treatment is susceptible of least abbreviation.
The principles of chemical equilibrium are (and have been for the
past half-century) fully as much required for intelligent consideration
of the simplest experiment, as is the theory of combining proportions
itself. Important parts of the theories of solutions and of the battery
are much more recent, but each is equally indispensable to the under-
standing of matters which cannot long be withheld from the notice of
the beginner. Surely space ought not to be saved by entire omission of
essential parts of the chief thing that makes chemistry at all worthy
of a place amongst the sciences. Nor may we attain brevity, no matter
how great the temptation, by condensing the passages on theory until
they reach the limit of comprehensibility by an expert. Without clear
exposition, full illustration, and frequent application, laws and princi-
ples simply repel, or worse still, mislead the beginner.
We reach the same conclusion from another view-point. Every
student should have access to, and should use, reference books devoted
especially to descriptive, industrial, historical, and physical chemistry,
and to mineralogy and crystallography — at least one good book in
each of these five subjects. With the help of the index, the veriest tyro
can find in a few moments, almost anything he wants in four out of
five of these branches. But just the opposite is the case with the
theory. Only an expert realizes what information he is in need of,
and knows under what titles to look for it. And often even the
expert would fail to understand the isolated sentence or paragraph
when found. In a large proportion of connections the beginner sim-
ply cannot use such a book for rapid reference at all. In many lines,
therefore, much may be left to outside reading, but for theory almost
no dependence can be placed on reference work in other books.
For the reasons enumerated above, an unusually large proportion
of space has been given to theoretical matters. The actual amount of
theory is no greater than is usual in books of the same class, but the
explanations are often fuller. Even so, the beginner will probably
find that some parts form reading as stiff as any he is accustomed to
undertake, without complaint, in physics or mathematics. It can only
Vlll PREFACE
be said that easily read modes of presenting the science of chemistry-
are apt to delude the beginner into thinking he has mastered the sub-
ject, when in reality he has simply been steered clear of the chief diffi-
culties.
The order of topics was determined by many considerations,
jointly. For example, in the first week of his work, a student may
encounter experiments, in connection with which, almost every part of
chemical theory might usefully be discussed. But mastery of the
theory must necessarily come bit by bit, and the theory is therefore
distributed through the book. Instead of being introduced as soon as
a fragment offers a chance for explanation, the treatment of each of the
various theoretical subjects, as far as possible, has been postponed until
a whole chapter could be devoted to it. The result makes subsequent
reference easier, and facilitates alterations in the order of study.
Thus, the hypothesis of ions is not mentioned as soon as it well might
be, because satisfactory treatment of it must follow the molecular and
atomic hypotheses of which it is an extension, and because the full
explanation of this hypothesis must be preceded by some account of the
phenomena of electrolysis and of the essential properties of solutions,
and, also, by a discussion of chemical equilibrium, a subject which of
necessity presupposes two or three months' work in chemistry. There
is another disadvantage which arises from a premature explanation of
the hypothesis of ionization. When it appears at an early stage, too
long an interval separates this subject from the study of the metallic
elements, and the details are largely forgotten before the field for their
chief employment is reached.
The paragraphs in smaller type are not intended for beginners, but
for advanced students and teachers, which accounts for the fact that
reference will frequently be found in them to subjects treated system-
atically in later chapters.
The exercises and problems are simply samples of some of the
various kinds of questions which might be raised in dozens at the end
of every chapter.
Recent works on general chemistry have been consulted during the
revision of the manuscript. Of these A. A. Noyes' admirable General
PREFACE ix
Principles, Ostwald's Grundlinien — a veritable tour de force, and Blox-
am's Chemistry may be mentioned as having proved most suggestive.
The author owes special thanks to several friends who have undertaken
the toilsome work of reading part or all of the book in manuscript or in
proof, and in particular to his colleagues Messrs. Julius Stieglitz, H. N.
McCoy, L. W. Jones, and E. S. Hall, to Dr. J. B. Tingle of Johns
Hopkins University, to Mr. C. M. Wirick of the R. T. Crane High
School, Chicago, and to Mr. Maurice Pincoffs of Chicago. The author
alone is responsible for any defects which may be inherent in the plan
of the book and for errors which may have escaped detection, but must
gratefully acknowledge the very great benefit the book has derived
from the friendly criticism of these gentlemen. Other corrections or
suggestions will be gladly received by the author.
CHICAGO, January, 1906. ALEXANDER SMITH.
CONTENTS*
CHAPTERS PAGES
I. Introductory 1 3
The Scientific Method — Three Illustrative Chemical Phenom-
ena — Two Characteristics of Chemical Phenomena.
II. Introductory II 17
Conservation of Mass — Energy — Applications of the Con-
ception of Energy in Chemistry — The Progress of Chem-
ical Actions Spontaneous or Forced — Chemical Activity
and its Cause — Of Simple and Compound Substances —
Some of the Fundamental Ideas Used by Chemists and the
Corresponding Terms — Methods of Work and Observation
in Chemistry.
III. Introductory III 41
The Law of Definite Proportions — The Law of Multiple Pro-
portions — The Measurement of Combining Weights —
The Law of Combining Weights — Equivalents — Atomic
Weights.
IV. Introductory IV •. 53
Symbols, Formulae, and Equations — Units of Measurement
in Chemical Work — Calculations in Chemistry.
V. Oxygen 61
Preparation of Simple Substances — Preparation of Oxygen —
Physical Properties of Oxygen — Specific Chemical Properties
— Chemical Properties of Oxygen — The Making of Equations
Again — Oxides — - Combustion — Oxidation — Means of Alter-
ing the Speed of a Given Chemical Action : By Change of
Temperature — Eapid Self-sustaining Chemical Action and
Means of Initiating It — Other Means of Altering the Speed
of a Given Chemical Change : By Change in Concentration ;
By Catalysis ; By Solution — Thermochemistry.
VI. The Measurement of Quantity in Gases 80
The Variable Concentration of Gases — The Method of Allow-
ing for Varying Concentration in Measuring Quantity in
Gases — The Relation of the Volume of a Gas to Temperature
— Mixed Gases — Densities of Gases — Vapor Densities of
Liquids and Solids.
* Tne titles of most of the descriptive paragraphs have been omitted from this table
because it is easier to find matters of this nature by consulting the index. The titles of all
the theoretical paragraphs, however, have been included.
XI
xii CONTENTS
CHAPTERS PAGES
VII. Hydrogen 92
Acids — Preparation of Hydrogen — Tests — Displacement —
Valence — The Valence of Radicals — How to Ascertain the
Valence of an Element or Radical — Physical Properties of
Hydrogen — Diffusion — Chemical Properties of Hydrogen —
A False Use of the Word "Affinity" in Explanation of
Chemical Actions — The Speed of Chemical Actions : A
Means of Measuring Activity.
VHI. Water 113
Natural Waters — Physical Properties of Water — Ice —
Steam and Aqueous Tension — Chemical Properties of
Water — Hydrates — Composition of Water — Gay-Lussac's
Law of Combining Volumes.
IX. The Kinetic-Molecular Hypothesis 128
Kinetic-Molecular Hypothesis Applied to Gases — Critical
Phenomena — Kinetic Hypothesis Applied to Liquids — Dif-
fusion in Liquids — Kinetic Hypothesis Applied to Solids —
Crystal Forms — Crystal Structure — Molecular Magnitudes
— Formulative and Stochastic Hypotheses.
X. Solution 146
General Properties of Solutions — The Scope of the Word —
Limits of Solubility — Recognition and Measurement of Solu-
bility — Terminology — Solution one of the Physical States of
Aggregation of Matter — Kinetic- Molecular Hypothesis Ap-
plied to the State of Solution — Kinetic-Molecular Hypothesis
Applied to the Process of Solution — Independent Solubility
— Solution of a Gas in a Liquid — Two Immiscible Solvents :
Law of Partition — Influence of Temperature on Solubility —
Phases — Equilibrium in a Saturated Solution — Metastable
Condition — Saturation — Properties of Solutions Propor-
tional to Concentration : Vapor Tension — • Freezing-Points
of Solutions — Densities of Solutions — Heat of Solution —
Definition of a Solution — Application in Chemical Work.
XI. Chlorine and Hydrogen Chloride 168
Chlorine — Chemical Relations of the Element — Hydrogen
Chloride • — The Kinetic Hypothesis Applied to the Interaction
of Sulphuric Acid and Salt — Classification of Chemical In-
teractions and Exercises Thereon — SUMMARY OF PRINCIPLES.
XII. Molecular Weights and Atomic Weights 190
Meaning of Avogadro's Hypothesis — Two Kinds of Laws in
Science — The Relative Weights of the Molecules — The Re-
lation between the Weights of Molecules and the Combining
Weights : Atoms — The Choice of the Value 16 for the
Atomic Weight of Oxygen — The Chemical Unit of Molecular
Weight — The Chemical Molecule — The Gram-Molecular
CONTENTS Xlll
CHAPTERS PAGES
(Molar) Volume — Determination of the Atomic Weight of
Each Element — The Accepted Atomic Weights of the Ele-
ments — Molecular Formulae of Compounds — Molecular
Formulae of Simple Substances — Molecular Weights of the
Elements at Various Temperatures — Dissociation of Com-
pounds— Chemical Interactions Studied by Observation of
Volumes : Molecular Equations — Other Methods of Deter-
mining Molecular Weights — Properties Possessed by Atomic
Weights, but not by Equivalents — Dulong and Petit's Law —
Periodic System : Molecular Structure — Application to the
Case of a New Element.
XIII. The Atomic Hypothesis .^ 217
XIV. The Halogen Family 226
The Chemical Relations of Elements — The Chemical Rela-
tions of the Halogens — Bromine — A Plan for Making Very
Complex Equations — Hydrogen Bromide — Iodine — Hydro-
gen Iodide — Fluorine — Hydrogen Fluoride — Association —
The Halogens as a Family — Compounds of the Halogens
with Each Other.
XV. Chemical Equilibrium 246
Reversible Actions — Kinetic Explanation — The Influence
of Homogeneous Mixture — The Influence of Molecular Con-
centration — Law Connecting Molar Concentration and Speed
of Reaction— The Condition for Chemical Equilibrium — The
Effect of Changes of Volume on Chemical Equilibrium
— Heterogeneous Equilibrium — Applications in Chemistry :
Displacement of Equilibria by Changes Affecting Concentra-
tion — Affinity vs. Solubility — Displacement of Equilibria by
Changes in Temperature : van 't Hoff's Law and Le Chate-
lier's Law — SUMMARY or PRINCIPLES.
XVI. Oxides and Oxygen Acids of the Halogens 263
Compounds of Chlorine Containing Oxygen — Nomenclature
of Acids and Salts — Salts and Double Decomposition —
Hypochlorites and Hypochlorous Acid — Hypochlorous
Anhydride — Hypochlorous Acid as an Oxidizing Agent :
Bleaching — Thermochemistry of Hypochlorous Acid —
Simultaneous, Independent Chemical Changes in the Same
Substance — Chlorates — The Separation of Substances by
their Solubility — Chloric Acid — Making of Equations
Again — Chlorine Dioxide : Chlorous Acid — Perchlorates,
Perchloric Acid, and Perchloric Anhydride — Oxygen Acids
of Bromine — Oxides and Oxygen Acids of Iodine — Todates
and lodic Acid — Various Acids Derived from One Anhy-
dride — Periodates and Periodic Acid — Chemical Rela-
tions.
xiv CONTENTS
CHAPTERS PAGES
XVII. Dissociation in Solution 281
Some Characteristic Properties of Acids, Bases, and Salts,
Shown in Aqueous Solution.
OSMOTIC PRESSURE : Phenomena Produced by Osmotic Pressure
— The Phenomena a Logical Consequence of Semi-permea-
bility— Measurement of Osmotic Pressure — Osmotic Pres-
sure and Concentration — Osmotic Pressure and Temperature
— An Analogue of Avogadro's Hypothesis — Determination
of Molecular Weights — Osmotic Pressure and Dissociation
in Solution.
DEPRESSION IN THE FREEZING-POINT OF A SOLVENT : Measure-
ment of Freezing-Points — Laws of Freezing-point Depres-
sion— Determination of Molecular Weights — Freezing-
Points and Dissociation in Solution — Boiling-Points and
Dissociation in Solution — Comparison of the Results of
the Three Methods.
THE APPLICATION OF THESE CONCLUSIONS IN CHEMISTRY :
The Constitution of Solutions of Acids, Bases, and Salts —
Nomenclature : The Ionic Hypothesis — Ionic Equilibrium —
Appendix : Recent Measurements of Osmotic Pressure.
XVIII. Ozone and Hydrogen Peroxide 300
Ozone — Hydrogen Peroxide — The Interaction of Barium
Peroxide and Sulphuric Acid — Thermochemistry of Hydro-
gen Peroxide — Peroxides : Chemical Constitution and
Molecular Structure.
XIX. Electrolysis 310
Introductory — Chemical Changes Connected with Electro-
lysis — Ionic Migration — Relative Speed of Migration of
Different Ions — Faraday's Laws — The Ionic Hypothesis —
Difficulties Presented by this Hypothesis — Amounts of
Electricity on the Ions — Re"sume" and Nomenclature —
Actual Quantities of Electricity Concerned in Electrolysis
— The Electrical Energy Required to Decompose Different
Compounds — Polarization — Conductivity for Electricity —
Degrees of lonization of Common Substances : Acids, Bases,
Salts — Degree of lonization of Water — General Remarks
on these Values — Comparison with the Results Obtained
by Other Methods.
XX. The Chemical Behavior of Ionic Substances 334
Solutions of lonogens are Mixtures — Each Kind of Ion in a
Mixture Acts Independently.
SALTS, IONIC DOUBLE DECOMPOSITION, PRECIPITATION : Salts
— Double Decomposition of Salts in Solution — Solution
and Precipitation of Salts — Individual, Specific Chemical
Properties of Each Ionic Material — Application in Chemical
Analysis — Hydrolysis of Salts.
CONTENTS XV
CHAPTERS PAGES
ACIDS AND BASES AND THEIR DOUBLE DECOMPOSITION WITH
SALTS : Hydrogen Salts — Hydrion — Modes of lonization of
Acids — Activity of Acids — Salts of Hydroxyl — Hydrox-
idion — Ionic Double Decomposition and Precipitation of
Acids and Bases — Ionic Double Decomposition and Activity.
NEUTRALIZATION : Acidimetry and Alkalimetry — Theory of
Neutralization — Indicators — Neutralization of Little-ion-
ized Substances — Thermochemistry of Neutralization —
Volume Change in Neutralization.
MIXED IONOGENS AND DOUBLE SALTS : Acid Salts — Basic Salts
— Mixed Salts — Double Salts.
KINDS OF IONIC CHEMICAL CHANGES : Disunion and Combi-
nation of Ions — Displacement of One Ion : Electromotive
Series of the Metals — Destruction or Formation of a Com-
pound Ion — Change in the Charges of Two Ions — Charge
and Discharge of Two Ions, Electrically — Two or More
Kinds of Ionic Action Simultaneously.
NON-IONIC MODES OF FORMING IONOGENS : Acids and Bases —
Salts.
XXI. Sulphur and Hydrogen Sulphide 367
Sulphur — Hydrogen Sulphide — Sulphides — Polysulphtdes
— The Chemical Relations of the Element Sulphur.
XXII. The Oxides and Oxygen Acids of Sulphur 378
Sulphur Dioxide — Sulphur Trioxide — Oxygen Acids of
Sulphur — Sulphuric Acid — Sulphates — Constitution of
Hydrogen Sulphate — Hyposulphurous Acid — Sulphurous
Acid — Illustration of the Effect of Concentration on Speed
of Interaction — Sulphites — Thiosulphuric Acid — Persul-
phuric Acid — Polythionic Acids — Sulphur Monochloride —
Thionyl Chloride — Sulphuryl Chloride.
XXIII. Selenium and Tellurium — The Periodic System 401
Selenium — Tellurium — The Chemical Relations of the Sul-
phur Family — Metallic and Non-metallic Elements — Class-
ification by Atomic Weights — Mendelejeff' s Scheme —
General Relations in the System — Applications of the Peri-
odic System.
XXIV. Nitrogen and its Compounds with Hydrogen 415
Nitrogen — Ammonia — Ammonium Compounds — Hydra-
zine — Hydrazoic Acid — Hydroxylamine — An Active State
of Hydrogen (Nascent Hydrogen) — Halogen Compounds
of Nitrogen.
XVI
CONTENTS
CHAPTERS PAGES
XXV. The Atmosphere. The Helium Family 426
Components of the Atmosphere — Air a Mixture — Liquefac-
tion of Gases — Liquid Air — Argon — Helium — Neon,
Krypton, and Xenon.
XXVI. Oxides, and Oxygen Acids of Nitrogen 438
Nitric Acid — Nitrates — Nitric Oxide — Molecular Com-
pounds — Nitrogen ^Tetroxide — Oxidizing Actions of Nitric
Acid — Nitrous Acid — Nitrites — Nitrous Oxide — Explo-
sives— The Principle of Transformation by Steps.
XXVII. Phosphorus 455
Phosphorus — The Electric Furnace — Phosphine — Phos-
phonium Compounds — Halides of Phosphorus — Oxides of
Phosphorus — The Phosphoric Acids — Phosphorous Acid —
Hypophosphorous Acid — Structural Formulae of Salts of
Hydrogen — Sulphides of Phosphorus — Comparison of Phos-
phorus with Nitrogen and with Sulphur.
XXVIII. Carbon and the Oxides of Carbon 473
Carbon — Calcium Carbide — Carbon Dioxide and Carbonic
Acid — Carbonates — R61e of Chlorophyl-bearing Plants in
Storing Energy — Photochemical Action — Carbon Monox-
ide— Carbonyl Chloride and Urea — Carbon Bisulphide.
XXIX. Some Carbon Compounds 490
The Hydrocarbons — Petroleum — Fractional Distillation —
Methane — Organic Radicals — Ethylene — Acetylene —
Benzene — Formic Acid — Acetic Acid — Oxalic Acid —
Carbohydrates and Fermentation — Alcohols, Esters, and
Ethers — Soap — Drying Oils — Cyanogen — Hydrocyanic
Acid — Cyanates — Thiocyanates.
XXX. Flame
XXXI. Silicon and Boron
Silicon — Silicon Hydride — Carbide of Silicon — Silicon
Tetrachloride — Silicon Tetrafluoride — Hydrofluosilicic Acid
— Silicon Dioxide — Silicic Acid — Colloidal Solution — Sili-
cates— Boron — Hydrides and Halides of Boron — Boric
Acid — Borates — Boron Trioxide — Nitride and Carbide.
XXXII. The Base-forming Elements
Physical Properties of the Metals — General Chemical Rela-
tions of the Metallic Elements — Hydrolysis of Halogen Com-
pounds, Used to Distinguish Metallic from Non-metallic
Elements — Salts of Complex Acids — Classification of the
Metallic Elements by their Chemical Relations — Occurrence
of the Metallic Elements in Nature — Methods of Extraction
from the Ores — Compounds of the Metals : Oxides and
509
518
530
CONTENTS
XVII
CHAPTERS
PAGES
Hydroxides — Compounds of the Metals : Salts — Solubilities
of Bases and Salts — Hydrated Founs of Salts Commonly
Used — Isomorphism .
XXXIII. The Metals of the Alkalies : Potassium and Ammonium. 548
The Chemical Relations of the Metallic Elements of the Alka-
lies — Potassium — the Hydride — Chloride — Iodide — Bro-
mide and Fluorides — Hydroxide — Oxides — Chlorate —
Bromate and lodate — Nitrate — Carbonate — Cyanide —
Sulphate and Bisulphate — Sulphides — Properties of Kal-
ion : Analytical Reactions — the Spectroscope — Rubidium
and Caesium — Ammonium Chloride — Ammonium Hydrox-
ide — Nitrate — Carbonate — Sulphate — Sulphides — Micro-
cosmic Salt — Ammonium Amalgam — Ammonion: Ana-
lytical Reactions.
XXXIV. Sodium and Lithium. -Ionic Equilibrium Considered Quanti-
tatively 569
SODIUM : Sodium Hydride — - Chloride — Hydroxide and Oxides
— Nitrate and Nitrite — Carbonate — Bicarbonate — Sulphate
— Thiosulphate — Phosphates — Tetraborate — Silicate —
Properties of Natrion : Analytical Reactions — Lithium.
IONIC EQUILIBRIUM CONSIDERED QUANTITATIVELY : The Simple
Case — Excess of One Ion — Special Case of Saturated
Solutions — An Illustration : Potassium Chlorate — Other
Illustrations — Direction of a Chemical Change Deter-
mined by Ionic Concentrations. Rule for Precipitation or
Solution of Insoluble Substances.
XXXV. The Metals of the Alkaline Earths 588
The Chemical Relations of the Elements — Calcium and its
Compounds — The Phase Rule, a Method of Classifying All
Systems in Equilibrium — Hard Water — Mortar and Cement
— Interaction of Insoluble Salts with Acids — Precipitation of
Insoluble Salts in Presence of Acids — Dissolving of Insoluble
Salts by Solutions of Salts — Glass — Calcion : Analytical
Reactions — Strontium and its Compounds — Barium and its
Compounds — Analytical Reactions of the Calcium Family.
XXXVI. Copper, Silver, Gold 614
The Chemical Relations of the Copper Family —
Copper and its Compounds — The Solution of In-
soluble Salts when Complex Ions are Formed — Silver and
its Compounds — Electro-plating — Photography — Gold and
its Compounds.
XXXVII. Glucinum, Magnesium, Zinc, Cadmium, Mercury. The
Recognition of Cations in Qualitative Analysis 641
XV111
CONTENTS
CHAPTERS
XXXVIII.
PAGES
Electromotive Chemistry 664
Factors and Units of Electrical Energy — Displacement Cells
— Potential Differences Produced by the Metals Singly
— Application to Cells — Other Applications : Couples :
Concentration Cells — Electrolysis: Discharging Potentials
— The Factors of Energy — Methods of Measuring Chemical
Activity.
XXXIX.
Aluminium and the Metals of the Earths
The Rare Elements of This Family — Aluminium and its
Compounds — Dyeing : Mordanting — Kaolin and Clay :
Earthenware and Porcelain.
XL. Germanium, Tin, Lead
XLI. Arsenic, Antimony, Bismuth.
681
692
707
XLII. The Chromium Family. Radium 722
The Chemical Relations of the Family — Chromium —
Derivatives of Chromic Acid — Chromic Compounds —
Chromous Compounds — Analytical Reactions — Molyb-
denum — Tungsten — Uranium — Radium — The Discovery
of the Element — Properties of Radium Compounds —
The Radiations of Radium Salts — The Decay of an Element.
XLIII. Manganese 737
XLIV. Iron, Cobalt, Nickel 746
XLV. The Platinum Metals 763
Ruthenium and Osmium — Rhodium and Iridium — Palla-
dium and Platinum.
INTRODUCTION TO
GENERAL INORGANIC CHEMISTRY
INTRODUCTION TO
GENERAL INORGANIC CHEMISTRY
CHAPTER I
INTRODUCTORY I
HUMAN knowledge has become, in recent times, so extensive and
complex that the truths ascertained have had to be divided, more or
less arbitrarily, into groups. Thus, the study of animals, their classi-
fication by structure, life-history, distribution, and so forth, form the
group known as zoology. Such a group is called a science, and in-
cludes a more or less distinct body of knowledge. That the boun-
daries of these groups are purely arbitrary, and do not exist in the
subject-matter itself, is seen at once in our own treatment of them.
Thus, for convenience, we take the structure of animals by itself and
style it anatomy ; but we include in the science of physiology the
study of the way in which the parts of both plants and animals per-
form their functions ; and we assemble cognate parts of two groups
in sciences like astro-physics and physical chemistry. The sciences,
therefore, are not mutually exclusive, and their boundaries overlap in
every direction.
The difficulty in deciding what are the most convenient boundaries
is as great with chemistry as with the other groups. At the one
extreme we have the abstract sciences, logic and mathematics. At
the other extreme lie the concrete sciences like geology, zoology, and
astronomy. .The former are not concerned primarily with the study
of matter at all, but with that of abstract conceptions. The latter
deal with definite aggregates of matter, such as the nature and his-
tory of a particular deposit of sulphur and their relation to those
of other deposits of sulphur, or the structure and history of a partic-
ular collection of organic material known as a pike, and their relation
to the structure and history of a mass of similar material called a
salmon. Between these two sets of sciences are the regions occupied
by the abstract-concrete sciences, physics and chemistry. These
sciences deal in part with the same portions of matter, but in a mor§
4 INORGANIC CHEMISTRY
abstract way than do geology or biology. To them, all specimens of
pure sulphur are alike, whether they have been formed by volcanic
action, or have been deposited by bacteria in an entirely different
manner. In particular, the line which divides physics and chemistry
from one another is often difficult to draw. It is assumed, however,
that the reader is already familiar with the elements of physics, and
so, in place of entering upon an academic discussion of the nature of
this line, we shall allow its location to emerge as we proceed.
The same principle of grouping is pursued within each field.
Thus the preparation and properties of chemical compounds is called
descriptive chemistry, and the content of this portion of the subject
is in turn classified according to a plan involving the consideration of
the constituents of each compound. The study of the proportions of
the constituents in compounds, of the conditions under which chem-
ical action occurs, and related matters of a more abstract character,
are grouped together in theoretical chemistry, which is likewise sub-
divided. Again, the means that have been devised for recognizing
the components of mixtures or compounds and measuring their quan-
tities constitute the several branches of analysis. The subdivisions
of chemistry of this kind are numerous.
The ideal in view in thus classifying the content of a science is to
convert it into an organized body of knowledge. The various ways
used to organize the facts of a science will be presented in detail as
opportunity offers. These ways constitute what is called the scientific
method.
It is only by following intelligently the way in which the science
is manufactured, step by step, out of the raw material furnished by
observation and experiment, that the student can gain a sound foun-
dation for more advanced work in the same science, or a mental train-
ing broad enough in its tendency to add measurably to his efficiency
in every other task. The chief object of useful thought, no matter
whether the problem is one of language, history, business, or life, is
to organize isolated facts into knowledge, and the means of success-
fully accomplishing this is the use of the scientific method.
Chemical Phenomena : Their Most Obvious Character-
istic. — Chemistry being primarily an experimental, and not a purely
abstract science, we can best make our first approach to it through the
consideration of some familiar chemical phenomena which shall serve
as, illustrations,
INTRODUCTORY I 5
When clean iron is exposed to air and moisture it becomes rusty,
first on the surface, and finally throughout its whole mass. The iron
with which we start is a dark-gray, metallic-looking substance, which
has a high specific gravity (about 7.5), is ductile, has great tenacity,
and is readily attracted by a magnet. Rust, on the other hand, is a
reddish or brownish substance which has an earthy appearance, is
much lighter specifically than iron (sp. gr. about 4.5), is brittle, and
is not attracted by a magnet. The phenomenon seems to have con-
sisted in the gradual substitution of the second of these substances
for the first.
The chemical fact here is that iron, when in contact with air and
moisture, gives rust. Consideration will show, however, that we do
not perceive this fact except by noting the physical qualities of the
original and of the final materials. The data describing a chemical
phenomenon consist in an enumeration of physical observations, like
the above physical properties of iron and rust. It is, indeed, an in-
variable characteristic of every chemical phenomenon that our informa-
tion is all derived from physical study of the materials. Thus, when
a candle burns, it undergoes a chemical change, and we observe a
solid, waxy substance disappearing progressively from view. A closer
study of the facts shows that a mixture of gases is rising from the
flame, and we find that the material of the candle is contained in this.
Here, again, we use physical means of observation.
It is further true of these, as of all chemical phenomena, that the
physical properties of the original and those of the final substances
are invariably entirely different. We observe also that this sort of
transformation is always abrupt. As the change spreads, the minute
fragment of iron of one moment becomes the minute fragment of rust
of the next, and is endowed at once with all the new properties in full
measure. No part of the mass loses its magnetic qualities completely,
for example, before it has reached the limit of change in color or ten-
acity. Since we are compelled to define the species of matter by
their constant physical properties (see p. 35), the possession by two
or more bodies of permanently different properties constitutes them
ipso facto samples of different materials. We may, therefore, say,
more briefly, that the products of a chemical change are recognized
at once to be new and different substances.
Thus the most obvious characteristic of a chemical phenomenon is
that all the physical properties of the substances alter, that this alter-
ation is abrupt, that, in factt the products are different substances, and
6 INORGANIC CHEMISTRY
that the recognition and study of such a phenomenon is accomplished
entirely by observations of a physical nature. Other characteristics,
less obvious, but equally constant, will presently be brought to light.
The chemical phenomena which are familiar are innumerable.
The burning of illuminating-gas or of kerosene, the coagulation of
the white of an egg, the decay of animal and vegetable matter, the
action of hard water on soap, the " burning " of limestone, and the
slaking of quicklime, are examples. The reader should analyze these
changes, applying the above criteria, and see for himself why they
are classed as chemical.
So far as this first characteristic is concerned, some merely physi-
cal phenomena will be recalled, which are not unlike the rusting of
iron. Thus, when water is raised in temperature, its properties begin
to alter : it becomes more mobile, its specific gravity changes, its
refracting power for light is modified, and so forth. At first, how-
ever, these changes proceed by imperceptible gradations and lack the
abruptness of chemical change. But when 100° C. is reached the
water passes into steam, and the whole of the properties of this gas
are different from those of water. Conversely, by cooling, the water
may be converted into ice, and again there is an abrupt and profound
change in properties. On this account, some chemists would classify
"changes of state" as chemical phenomena. More usually, however,
the absence of the other characteristics yet to be mentioned is held to
justify the more common view, and ice, water, and steam are regarded
as identical, and not different substances.
The very language we use bears testimony to the universal ac-
ceptance of the view that a physical change does not constitute a
fundamental alteration. Thus, we have solid lead and molten lead,
air (the gas) and liquid air. In the case of water alone has it been
found convenient to distinguish ice, water, and steam by separate names.
Fact and Generalization or Law. — If the preceding paragraph
be reexamined, it will be seen that a number of facts are mentioned
in it. We begin the organization of knowledge according to the
scientific method by trying to determine the facts. Thus, we find
some specimens of iron are variously colored, and some are brittle.
Examination shows, hoAvever, that the former peculiarities are due to
paint, for example, and the latter to the presence of carbon and other
foreign materials in the iron (cast iron), Finally, we ascertain the
facts that iron itself is gray and tough.
I1TTRODUCTOKY I 7
Facts are the ultimate units of the structure of a science. Thus, a single iso-
lated observation, no matter how accurately made, does not furnish us with the
sort of fact that can receive a permanent place in our collection. It is only after
much research and thought that we can ascertain which are the fixed elements in
the variety of experience in any line, and so determine what are the facts, in the
sense in which we have used the term.
Putting together a statement like that appearing in heavy type
above, is the second step. We examine facts of a like kind, or per-
taining to like phenomena, to see whether any general statement can
be made that will cover some feature common to the whole of them*.
In the above illustrations, after settling the intrinsic properties of
several substances, and then determining the facts about the way in
which these properties are affected under certain circumstances, Ave
decide that, when all the properties change abruptly in a permanent
way, the cases in which this takes place shall constitute a distinct
class, and we call them cases of chemical change. The statement
which in a few words or phrases sums up those features of all the
phenomena of like kind which are constant, is called a generalization.
Often it is called a law: there is no decided difference in the usage
of the two words. Sometimes it is described as a principle of the
science. A generalization or law in chemistry, therefore, is a brief
statement describing some constant mode of behavior.
Of course, the same set of facts may be viewed in many ways. Picking out
the relationships which are most comprehensive and, at the same time, are best
fitted to form part of still broader generalizations, or to take their places alongside
of equally broad ones, requires the highest ability. The most important laws,
like that describing the behavior of gases when compressed (Boyle's law), are
usually connected with the names of the men by whom the relationships were dis-
covered and the generalizations formulated.
The word "formulate," as applied to a law, is preferable to the word "dis-
cover." The latter is ambiguous and suggests that the law existed before it was
found. Even the relation which it puts into words, did not, properly speaking,
exist, because relations are picked out of a complex by the mind, and the particu-
lar relation selected is a property of the mind and not of the constituents of the
complex itself.
The reader must beware of the misconception that a law enforces behavior in
accordance with its tenor. The mere statement that every piece of matter attracts
every other, cannot compel a stone to fall ; nor can the mode of behavior of one
falling stone persuade any other to do likewise. The law is simply a record
of what is invariably observed to happen.
It is, therefore, also very misleading if we permit ourselves to say that Boyle's
law "acts" so as to "cause" gases to behave in a certain way, or that the law
"operates" to "produce" a certain behavior, or that other behavior is "impos-
sible" to the gas, or that the law of cohesion " intervenes " when the gas is under
8 INORGANIC CHEMISTRY
low pressure, and causes its behavior to "diverge from that required" by Boyle's
law, or to say that a gas "disobeys" Boyle's law. In scientilic discussions, such
figures of speech are alone permissible as throw light on the subject. Phrases
like the above, common as their use is, have been selected apparently with a view
to introducing a maximum of distortion and obscurity. It is the gas that " acts "
and gives rise to the making of Boyle's law, and the latter is only an epitome of
the way it acts. Everything that may be conceived may also be possible in nature,
although there are many things which have not yet been known to occur. There-
fore, we may not say that it is " impossible " for a gas to alter its volume otherwise
than inversely with the pressure. The progress of science would be almost com-
pletely arrested, if, every time we succeeded in formulating a seemingly satisfac-
tory statement of truth in regard to some set of phenomena, the exhibition by
nature of any behavior which was in conflict with our statement became forthwith
"impossible." It is not the gas which "diverges" from our statement or "dis-
obeys " our law, but our statement which is proved by the behavior of the gas to
be inaccurate. Our procedure, in such cases, is always more logical than our lan-
guage, for we never attempt to cure the gas of its error, but always the law itself
by suitable modification in its phraseology.
The Explanation of Rusting. — In considering the complete
disguise which is assumed by a substance that has undergone chemi-
cal change, the first question which arises is : Can we in any way
account for this great and permanent change in properties which very
generally distinguishes the chemical phenomenon from the physical ?
Some additional facts will be required before we can. answer this
question. Many attempts have been made, from the earliest times, to
learn the exact nature of the change known as rusting. It was found
that many metals besides iron underwent a somewhat similar altera-
tion, and that, where the transformation did not occur spontaneously,
heating prompted it. The first fact which seemed to throw light on
the subject was the observation that a piece of metal becomes heavier
when it rusts. This was noticed as early as 1630, by a French physi-
cian, Jean Rey. His work was done chiefly with lead and tin, the
former of which gives a dirty yellow, and the latter, a white powder,
on rusting. He inferred correctly from his experiments that contact
with the air had something to do with this chemical change. Other
investigations on the same subject were made by Boyle (1627-1691),
Mayow (1645-1679), and Hooke (1635-1703), in England, and they led
to the same conclusion. The increase in weight was to be accounted
for, therefore, by the supposition that some material from the air had
been added to the metal during the process. In other words, iron, for
example, was one substance composed of -iron only, and rust was
another substance composed of iron and some other material taken
INTRODUCTORY I 9
from the atmosphere ; and, the two substances being different in
composition, their properties might naturally be expected to differ.
The substance taken from the air was subsequently named oxygen.
A rough imitation of the rusting of iron may be shown to be accompanied by
an increase in weight. Iron powder is suspended by means of a magnet over one
pan of a balance, and the equipoise is restored by placing small shot on the other
pan. When the iron is heated, union with oxygen begins, and, after a time, the
pointer inclines markedly to one side. The product here is magnetic oxide of
iron (Fe3O4), however, and not rust (Fe203, icH2O).
An Older View: Phlogiston. — Simple, and to a certain extent
satisfactory, as this explanation appears to us, it must be said that it
gained little sympathy from the contemporaries of these men. The
other investigators had been prejudiced by a remarkable hypothesis
which was supposed to explain both rusting and combustion. Starting
with a suggestion of Plato's, that during combustion some material
escaped from the burning body, and that the flames and heat repre-
sented the vigor with which this substance rushed out, Stahl and
Becher invented the idea that a substance, which they called " phlo-
giston," was contained in all materials capable of rusting or burning.
Its escape accounted for the phenomena of combustion, and its absence
for the alteration in properties of the residual substance. They were
perfectly aware that the material was heavier after rusting than
before, but refused to sacrifice their hypothesis to a mere fact like
this. So they ingeniously appended the suggestion that phlogiston
was a substance which not only was not subject to gravitation, but
possessed the opposite property of levity. Thus, its escape rendered
the material from which it issued heavier than before. Instead of
demanding the preparation and examination of phlogiston itself, and
the demonstration that it weighed less than nothing, the generality of
chemists of that age accepted the idea without proof. It is not sur-
prising, therefore, that many of their attempts to explain chemical
phenomena on the basis of an arbitrary assumption like this should
have proved confusing and infertile. The fact that foreign matter
was actually gained by a body during the process of rusting was not
generally accepted until it was demonstrated anew by Lavoisier
(1774). He showed that a portion of the air really disappeared when
tin rusted, and that the increase in weight of the tin corresponded
with the loss in matter of the air. The whole development of chem-
istry was stunted by the general belief in the conception of phlogiston
10 INOHGANIC CHEMISTRY
and, during the one hundred and fifty years which passed between
Jean Key's discovery and Lavoisier's, relatively little progress was
made.
The introduction of the balance into the chemical laboratory, and the first
use of measurements of weight as a means of exploring and explaining chemical
changes, is frequently ascribed to Lavoisier. As a matter of fact, it is difficult to
state when measurement of weight first became the chief ally of the chemist in his
work. Important and conclusive results were obtained by its means, however,
before the time of Lavoisier, for example, by Jean Key, Boyle, and Black (1728-
1799).
Explanation in Science. — The word explanation, which was
employed repeatedly in the last two sections, is used in science as the
name of a definite process. It stands simply for a description in
greater detail. Thus, when, to the acquaintance witli the outward
manifestations of rusting, we are able to add the further description
that it is produced by the union of oxygen from the air with iron, we
feel increased satisfaction, and we say that an explanation has been
found. Sometimes we get this satisfaction by an explanation which
is a description, not of additional facts, but of some imaginary con-
dition which takes their place in our thought. The hypothesis (or
supposition) of phlogiston was somewhat of this nature (see Formula-
tive hypothesis).
There is no such thing as a final explanation. At the very next
step, when we ask why the union of oxygen and iron produces a body
that is red and non-magnetic, we are compelled to say we do not know.
Even a supposition in regard to how this happens is as yet lacking.
An explanation in science never professes for a moment to give the
reasons for any occurrence. We simply do not know why behavior in
nature is as it is. Hence statements like that in regard to union with
oxygen, constitute all the answer that we can give to the question with
which a recent section (p. 8) opened.
Three Illustrative Chemical Phenomena. — Since oxygen is
an invisible gas, the demonstration that rusting consists in the union
of this gas with a metal, requires somewhat complicated apparatus.
The next illustration, while lacking historical interest, is simpler, be-
cause both substances are visible, and are easily handled.
Iron unites, not only with oxygen from the air, but also, with almost
equal ease, with sulphur. To study the behavior of the materials we
must know their properties, The properties of iron we have already
INTRODUCTORY I
11
FIG. l.
enumerated (p. 5). Sulphur is a pale-yellow substance of low spe-
cific gravity (sp. gr. 2). It is easily melted (m.-p. 115° 0). and, al-
though, it does not dissolve in water, it may be dissolved completely in
carbon disulphide (41 : 100 at 18°).* It crystallizes in rhombic forms
(Fig. 1). Now, when iron filings and powdered sulphur are rubbed
together in a mortar, the product, although it has different color from
either of the constituents, is still really com-
posed of the two original kinds of matter, side
by side. With the help of a microscope and
a needle, they can be picked apart completely.
By manipulation of the mixture with a mag-
net, we may remove some of the iron without
much difficulty.
Again, using the solubility of sulphur in
carbon disulphide and the insolubility of iron,
we may shake the mixture with this liquid in
a test-tube, and so dissolve out the sulphur (Fig. 2). When the con-
tents of the tube are poured on to a filter (Fig. 3), the liquid (the
filtrate) runs through, carrying all dissolved matter with it, and leav-
ing undissolved matter behind. The
former may be allowed to evaporate
(Fig. 4), when yellow crystals of rhom-
bic outline will be found to be the sole
residue. The dark material remaining
on the filter, when dry, is wholly at-
tracted by a magnet. All these facts
convince us that the properties of the
components are still unaltered, and that,
therefore, no chemical change has oc-
curred.
If we now put some of the original
mixture into a test-tube and warm it,
we soon notice a rather violent develop-
ment of heat taking place, the contents
begin to glow, and what appears to be a form of combustion spreads
through the mass. The heating employed at the start fa^ls far
short of accounting for the much greater heat produced. When
these phenomena have ceased, and the test-tube has been allowed to
* This expresses the fact that 41 parts, by weight, of sulphur dissolve in 100
parts, by weight, of carbon disulphide at 18° C.
FIG. 2.
12
INORGANIC CHEMISTRY
cool, we find that it now contains a somewhat porous-looking, black
solid. This material is brittle ; it is not magnetic ; it does not dissolve
in carbon disulphide ; and close examination, even under a microscope,
does not reveal the presence of different kinds of matter. This sub-
stance is known to chemists
as ferrous sulphide, and, as
we see, its properties are
entirely different from those
of the constituents.
A substance formed in
this way by the union of
other materials is called a
compound. The rusts given
by various metals are there-
fore compounds also.
The second illustration
is selected on account of its
historical interest. One of
the earliest chemical changes
FlG 3 in which a gas, recognized
to be distinct from air, was
observed among the products, was noticed by Priestley (1774).* The
observation was made with mercuric oxide, a bright red, rather
heavy powder. When this substance is heated (Fig. 5), we find that
a gas is given off, which is easily shown to be different from air, since
a glowing splinter of wood is instantly relighted on being immersed
in it. The gas is pure oxygen.f
We notice also during the heating
that a sort of mirror appears on
the sides of the tube. As this
shining substance accumulates it
takes the form of globules, which
may be scraped together. It is,
in fact, the metal mercury, or quicksilver. If the heating continues
long enough, the whole of the powder eventually disappears, and is
converted into these products.
* An English nonconformist minister who occupied his leisure time with exper-
iments in chemistry. He afterwards moved to the United States, and died in
Northumberland, Pa.
t Air (q.v.) is a mixture of which only one-fifth is oxygen.
FIG. 4.
INTRODUCTORY I
The extensive nature of the change in properties in this case is evi-
dent. It should also be observed that continuous heating is required
to maintain this change in operation. It differs markedly from the iron
and sulphur case in this respect. When the flame is removed, the evo-
lution of oxygen ceases. The significance of this will appear shortly.
The third and last example is taken purposely in order to illustrate
the variety of ways in which chem-
ical change may be carried out. It
is the interaction of silver nitrate
and sodium chloride (common salt).
The substances may be recognized
by the form of the crystals of which
they consist. The latter is composed
of small cubes (Fig. 6), while the
former presents a less familiar form
geometrically (Fig. 7). Both sub-
stances are capable of being dis-
solved in water and, for this experi- FIG. 5.
ment, portions of each substance are
shaken in separate vessels with water, until none of the solid remains.
When the solutions are now poured together, we observe that the
clear liquids at once become opaque, and that a dense
mass of white, solid material appears suspended in
the mixture (Fig. 8). This white substance consists
of an extremely fine powder without any observable
crystalline form. We know at once that it must
represent a new substance, since it would not have
appeared had it been soluble in water like the two
materials from which it was
made. We continue adding the one liquid gradu-
ally to the other until no further formation of
this solid takes place, and then stop. By filtra-
tion (Fig. 3), we obtain the insoluble material
(the precipitate) upon the filter paper, and the
clear liquid (the filtrate) passes through and is caught in the vessel
below.
We are confronted with two possibilities : either both the original
materials have come together to form one white insoluble material, or
some other product (or products) may be present in addition to it. In
the latter case, search must evidently be made in the liquid. By evap-
FiG. 7.
14
INORGANIC CHEMISTRY
orating the filtrate in a suitable vessel (Fig. 4), we find that the second
assumption represents the fact, for a considerable quantity of a white
crystalline substance remains. The homogeneous character of this
shows that there was but one product in solution, while the same prop-
erty of the precipitate shows that there are but two products altogether.
The insoluble material is composed of silver
and chlorine, and it is known as silver chloride.
Like some other compounds of silver, it darkens
on exposure to light, turning first purple and then
brown, and being decomposed by this agency into
its constituents. The soluble solid obtained from
the filtrate, we recognize as identical with a
mineral, sodium nitrate, which is found in Peru.
Its crystals are rhombohedral (Fig. 9). They
resemble cubes which have been slightly dis-
torted by pressing inwards two opposite corners.
A strict investigation of all four substances
shows that the following statement presents the
facts in regard to the nature of the change : — Nitrate of silver,
consisting of silver, nitrogen, and oxygen, with sodium chloride, con-
sisting of sodium and chlorine, have changed
into silver chloride, consisting of silver and chlo-
rine, and sodium nitrate, consisting of sodium,
nitrogen, and oxygen. This change presents
several features which distinguish it from the
previous ones : it is much more complex ; it takes
place in the presence of water ; it requires no
heating for its promotion ; and the change is complete the instant the
materials have been mixed, while the others required a good deal erf
time for their accomplishment.
The Kinds of Chemical Change. — Having these three cases
before us, — and they are types of most chemical phenomena, — we now
proceed to analyze them, and so take another step towards explaining
(i.e., describing in detail) the nature of a chemical phenomenon. In
the iron and sulphur experiment two materials were used, and a dif-
ferent one with new properties was produced. Here, in chemical
language, the first two substances united or underwent combination.
The rusting of iron, lead, and tin belongs also to this class. In the
second illustration one material was used, and it was driven apart by
INTRODUCTORY I 15
heating so that two new ones arose. A chemical phenomenon of this
kind is called a decomposition. The last was the most complex, but
a little examination shows that it does not present any novel featiires.
The statement we gave (p. 14) about the way in which the constit-
uents were distributed, before and after the change, shows that the
original materials were altered by decomposition, and that the products
were formed by recombination of these on a different plan. A third
variety of chemical change consists, therefore, in the concurrence of
both of the first two kinds in the same action. As several different
varieties of this sort of complex redistribution of material can be dis-
tinguished, a distinct name is given to each. The present is called a
double decomposition or metathesis.
When we encounter other chemical changes, we shall find the
extent of the stride we have taken in this critical analysis of a
few examples. It will then appear that the chemical changes of
matter are not nearly so various as we might have anticipated. In
fact, there are few changes which cannot be placed in one of the
above categories. Those that cannot be so placed belong to a fourth
sort of change which, for the sake of completeness, may now be.
mentioned.
It occasionally happens, especially in the case of compounds of
carbon (see Urea), that one single kind of material turns into another
single kind of material. Nothing is added and nothing removed, yet
the new substance has different properties in every respect from the
old. Most of those substances whose transformation is definitely
assigned to this class contain several constituents, but a rough notion
of this sort of chemical change may be obtained by considering the
two forms of phosphorus. One of them is pale yellow in appearance,
easily melted, and very easily ignited ; the other is red, does not melt
on being heated, and is difficult to ignite. The latter is made from
the former by heating it continuously for some hours in a closed ves-
sel at about 300°. As no material is taken up, the weight of the sub-
stance is unchanged, and yet, when the vessel is opened, the common
phosphorus is found to have turned into the red variety. As the
foregoing paragraphs show, we have good reason to believe that the
properties of a substance are intimately related to its constituents.
Carrying out this idea, the hypothesis (or suggestion) has been made
that, since here also the properties change, there must be some read-
justment of the material, even in cases like this. Hence, we designate
changes of this kind internal rearrangements.
16 INORGANIC CHEMISTRY
In the completeness of the transformation, and in the fact that only one origi-
nal substance and one product are required, the physical changes of the nature of
melting a solid and vaporizing a liquid resemble the fourth variety of chemical
change. Where, then, is the line between chemistry and physics to be drawn? It
is in this fourth group only that the difficulty is encountered. All agree that
warm solid phosphorus is chemically identical with cold phosphorus. Nearly all
scientific men at present assign the study of the melting and the vaporization of
phosphorus and other substances to physics. Some chemists consider the solu-
tion of phosphorus in carbon disulphide or some other solvent (it is practically
insoluble in water), although the material is recovered unchanged by evaporation
of the liquid, as a chemical change. But the great majority regard this as physi-
cal also. If not chemically different, the solid, liquid, gaseous, and dissolved forms
of a substance must be classed as mere physical states of aggregation. On the
other hand, red phosphorus is held by most chemists to be chemically different
from yellow phosphorus. Many kinds of matter show as much variety in form as
phosphorus, and some show more.
In solving an ordinary puzzle, we work knowing that some simple solution
exists. When studying nature, we are saved much embarrassment by remember-
ing that, in cases like this, we are not seeking for a clear-cut distinction of whose
existence we are assured in advance. Nature is under no obligation to furnish
easily classifiable facts at all. The limitations of our minds compel us to classify
as far as we are able to do so. On this plan alone can we master the infinity of
detail. But the resulting system exists in our own minds, where it originated, and
not in nature. In the present instance, we are seeking, as always, to make a
distinction for our own convenience. But, at present, the effort to make a final
distinction that can be used consistently causes more inconvenience than the alter-
native of letting the matter rest. The history of previous experiences leads us to
hope that, with fuller knowledge, we shall be able, sooner or later, to construct a
system of classification even for this obscure region.
The Second Characteristic of Chemical Phenomena. — We
are now able to make another generalization (or condensed statement
of fact) : In chemical phenomena substances enter into combination,
come out of combination, or change their associates in combination or
their state in the compound. In other words, the material changes its
composition or its constitution.
This generalization furnishes an explanation of the former (p. 5)
to a certain extent. Each individual substance has its own composi-
tion and its own set of physical properties. Hence, when, by chemical
transformation, substances of new and different composition are pro-
duced, the materials must simultaneously acquire the specific proper-
ties of the new substances.
Summary. — Thus far, we have learned that chemistry deals with
the changes in composition and constitution which substances undergo,
and with the alteration in properties which accompanies and gives
evidence of those changes.
CHAPTER II
INTRODUCTORY II
IF we now return to the three illustrations of chemical phenomena
which we have been studying (pp. 10-14), we shall find a new question
arising naturally out of them. This is, whether the mass of the ma-
terials is altered, as are the other attributes, in these chemical changes.
Third Characteristic of Chemical Phenomena : Conserva-
tion of Mass. — The most painstaking chemical work seems to show
that, if all the substances concerned in the chemical change are
weighed before and after the change, there is no evidence of any
alteration in the quantity of matter. The two weights, representing
the sums of the constituents and of the products respectively, are,
indeed, never absolutely identical, but the more careful the work and
the more delicate the instrument used in weighing, the more nearly
do the values approach identity. We are able to state, therefore, as
a third characteristic of all chemical phenomena, that the mass of a
system is not affected by any chemical change -within the system.
This statement simply means that the great law of the conserva-
tion of mass holds true in chemistry as it does in physics. Chemical
changes, thoroughgoing as they are in respect to all other properties,
do not affect the mass ; an element carries with it its weight, entirely
unchanged, through the most complicated chemical transformations.
This is the only attribute which persists.
A law, as we have seen (p. 7), is a condensed statement describing some con-
stant mode of behavior. It is simply a summary of our experience. As such, it
is subject to modification when a fact is discovered with which it conflicts. Thus,
it is perfectly possible that we may yet find cases of demonstrable changes in
weight accompanying other physical or chemical changes in a limited system.
Indeed, it has more than once been alleged that such changes have been observed.
It used to be a law that the earth was flat. It is now more correct to say that a
limited area of perfectly level ground is very nearly flat.
It will be observed that the phrasing of the above law carefully limits its
scope to amounts of matter such as are dealt with in laboratory experience. We
have no evidence on which to make any statements about the mass of matter in
more extensive chemical changes. A common form of the law, to the effect that
17
18
INORGANIC CHEMISTRY
" the mass of matter in the universe is unchangeable in amount," is not a law at
all, in the only sense in which the word is used in science. It is a statement in
regard to supposed facts which are almost entirely beyond our experience. It is,
therefore, a proposition of a transcendental (that is, transcending experience)
nature, and has its proper place in metaphysics. Astronomical observation, it is
true, has as yet furnished no evidence of changes in the mass of our own or other
celestial systems. But, absence of evidence to the contrary, especially considering
the relatively limited scope of our knowledge, both in respect to space and time, is
far from being proof of the correctness of the proposition.
Superficial observation, as of a growing tree, might seem to give
evidence of the very opposite of conservation of matter. But here
the carbon dioxide gas in
the air, the most impor-
tant source of nourishment
for plants, is overlooked.
Similarly the gradual dis-
appearance of a candle by
combustion seems to illus-
trate the destruction of
matter. But if we insert
sticks of sodium hydroxide
in a U-tube (Fig. 10) to
catch the gases which rise
through the flame, we find
that the gases weigh even
more than the part of the
candle which has been sac-
rificed in making them.
When we take account of
the weight of the oxygen
obtained from the air which
sustains the combustion,
we find that there is really
neither loss nor gain in
weight. If we carry out chemical changes in closed vessels (Tig. 11),
which permit neither escape nor access of material, we find that the
weight does not alter.
One way of stating the difference between chemistry and physics is to say that
changes in which both the mass and the identity of the substance are conserved
belong to the latter, while those in which the mass alone is conserved belong to
chemistry.
FIG. 10.
INTRODUCTORY II
19
Physical Concomitants of Change in Composition. — Study
of the three typical chemical changes described in the last chapter
may now be resumed, in order to see whether anything further of a
general nature is characteristic of such phenomena. We recall at
once that a prominent feature of the union of iron and sulphur was
the heat which, as shown by the glow
spreading through the mass, seemed to
be developed after the action was once
started. It is found that many chemi-
cal changes are like this one, in exhibit-
ing simultaneously the production of
very perceptible amounts of heat. On
the other hand, the decomposition of
mercuric oxide, as was pointed out
(p. 13), owed its continuance to the per-
sistent application of heat, and ceased
so soon as the source of heat was with-
drawn. Here, apparently, heat was
consumed during the progress of the
change, and the chemical action was
limited by the amount of heat supplied.
The production or consumption of heat
may, therefore, be a feature of chemical
change.
In 'the iron and sulphur case, as in other chemical actions where
the heat developed is great, light also was given out. In the last of
the three actions, on the other hand, we obtained a substance (silver
chloride), which may be kept for any length of time in the dark, but,
by the action of sunlight is broken up into its constituents (p. 14). It
would appear, therefore, that light may be given out or used in con-
nection with chemical change. Noting these facts stimulates us to
look for other similar concomitants of changes in composition.
If we dip two wires from a battery or dynamo into a solution of
nitrate of silver (Fig. 12), such as was used in the third experiment,
we observe the instant production of a coating of silver on the nega-
tive wire. By preparing the solution properly and allowing the
electricity to flow through it for a sufficient length of time, all of
the compound can be decomposed and all its silver deposited. It is
needless to say that this release of the silver from chemical combina-
tion and liberation of the metal at the electrode, goes on only so long
FIG. 11.
20
INORGANIC CHEMISTRY
FIG. 12.
as the current of electricity is employed, and that electricity is con-
sumed in the process. Very many substances can be decomposed in
this way.
The inverse of this is
likewise familiar. If we
place in dilute sulphuric
acid a stick of the metal
zinc, we find that a gas
is given off rapidly (Fig.
13), that the zinc gradu-
ally dissolves, and that a
large amount of heat is
developed. Under favor-
able circumstances, the
liquid may even rise spon-
taneously to the boiling-
point. This form of the
action produces heat. If,
however, we attach the same stick of zinc to a copper wire, and,
having provided a plate of platinum also connected with a wire,
immerse the two simultaneously in the acid
(Fig. 14), then a galvanometer, with which the
wires are connected, shows at once the passage
of a current of electricity round the circuit.
Exactly the same chemical change goes on as
before. The sole difference is that the gas
appears to arise from the surface of the plati-
num. It is easy to show, however, that the plat-
inum by itself is not acted upon by dilute acids,
and, in this case, undergoes no change whatever ;
it serves simply as a suitable conductor for the
electricity. Here, then, in place of the heat
which the first plan produced, we get electricity.
The arrangement is, in fact, a battery, for a bat-
tery is a system in which a chemical action which
would otherwise give heat furnishes electricity in-
stead. Thus, electricity may be consumed or pro-
duced in connection with a change in composition.
Even violent rubbing in a mortar, in the case of some substances,
can effect an appreciable amount of decomposition in a few minutes.
FIG. 13.
INTRODUCTORY II
21
In this way silver chloride can be separated into silver and chlorine,
just as by light. It is the mechanical energy which is the agent, and
part of it is consumed in producing the change, and only the balance
appears as heat. Conversely, the production of mechanical energy, as
FIG. 14.
the result of chemical change, is seen in the behavior of explosives
and in the working of our muscles. Thus, mechanical energy may be
used up or produced in chemical changes.
Summing our experience up, we may state that no change in com-
position occurs without some concomitant, such as the production or
consumption of heat, light, electricity, or, in some cases, mechanical
energy.
Classification of the Concomitants of Change in Composi-
tion: Energy. — The problem of classifying (i.e., placing in a suitable
category) things like heat, light, and electricity has occupied much
attention. They do not possess mass. In all changes in composition,
one of these natural concomitants is given out or absorbed, sometimes
in great amount, yet, in none is any alteration in weight observed.
Nor may these concomitants be overlooked, simply. A conception is
a real thing ; a religious belief may be most real and potent. There
are many things which are real, although they are not affected by
gravitation. In the present instance we reason as follows :
A brick in motion is different from a brick at rest. The former
can do some things that the latter cannot. Furthermore, we can
22
INORGANIC CHEMISTRY
easily make a distinction in our minds. The brick can be deprived of
the motion and be endowed with it again. Thus, we can get the idea
of motion as a separate conception. Similarly, we observe that a piece
of iron behaves differently when hot, and when cold, when bearing a
current of electricity, and when bearing none. We conceive then
of the brick or the iron as having a certain amount and kind of matter
which is unalterable, and as having motion, heat, or electricity added
to this or removed. Thus, we describe our observations by using two
categories, one of which includes the various kinds of matter, and
the other, various things whose association with it seems to be invari-
able and is often so conspicuous.
At first sight, these concomitants of matter seem to be quite
disparate. But a relation between them can be found. If the heat
of a Bunsen flame or of the sun is brought
under a hot-air motor (Fig. 15) violent
motion results. Again, if the motor is
connected with a dynamo, electricity may
be generated. Still again, if the current
flows through an incandescent lamp, heat
and light are evolved. Conversely, when
motion is impeded by a brake, heat appears.
When a current of electricity is run
through the dynamo, motion results. But
the most significant facts are still to be men-
tioned. The heat absorbed by the motor
is found to be greater when the machine
is not permitted to move and do work,
than when it is. Thus, it is found that
when work is done some heat disappears,
and is, in fact, transformed into work.
Similarly, when the poles of the dynamo are properly connected
and electricity is being produced, and only then, motion is used
up. This is shown by the effort required to turn the armature
under these circumstances, and the ease with which it is turned
when the circuit is open. So, with a conductor like the filament
in the lamp, unless it offers resistance to the current and destroys
a sufficient amount of electricity, it gives out neither light nor
heat. Finally, motion gives no heat xinless the brake is set, and
effort is then demanded to maintain the motion. These experiences
lead us to believe that we have here a set of things which are funda-
FIG. 15.
INTRODUCTORY II 23
mentally of the same kind, for each form can be made from any of
the others. We have, therefore, invented the conception of a single
thing of which heat, light, electricity, and motion are forms, and to it
we give the name energy : energy is -work and every other thing which
can arise from work and be converted into work (Ostwald).
Closer study shows that equal amounts of electrical or mechanical
energy always produce equal amounts of heat. There is never ob-
served any loss in any of the transformations of energy any more
than in the transformations of matter. Hence, J. K. Mayer (1842),
Colding (1843), and Helmholtz (1847) were led independently to the
conclusion that in a limited system no gain or loss of energy is ever
observed. This brief statement of the results of many experiments
is called the law of the conservation of energy.
A current form of this law, namely, that "the total amount of energy in the
universe is a constant quantity," is open to the same objection as the correspond-
ingly flamboyant form of the law of conservation of mass criticized above. It has
a more effective sound than the one we have given. Unfortunately, it is not only
immensely in excess of any statement that present results of scientific work can
justify, but is probably far beyond the limits of possible scientific observation.
Scientific statements of fact can never err by being too conservative.
Matter and Energy as Concepts, and Definitions of the Lat-
ter. — The foregoing paragraphs about energy bring up the question of its relation
to matter. This relation can be made clear only by a somewhat elaborate dis-
cussion of our fundamental conceptions.
The only real, first-hand knowledge which we possess, is that of our states of
consciousness. All else, consisting, for example, of the way in which we inter-
pret and describe our experience, is constructed out of our heads, so to speak.
Now, we become aware of certain things which we call sensations, and seek to con-
struct a mode of correlating and describing them. Our universal habit is to speak
as if they were produced by something outside our minds, and so we begin the
manufacture of an external universe. In course of doing this, we encounter
some things which seem to occupy no particular space, which move from object
to object, and possess no weight. One of these affects our eye, or a piece of
chloride of silver, for example, yet escapes touch, and passes through glass as
easily as through a vacuum. After consideration of our experience with this sort
of thing, some of which has been detailed above, we decide that we shall posit the
existence of energy.
Other things we encounter which appeal to the sense of touch and seem to
possess more definitely located 'qualities, including weight. Another conception
is needed to account for these ; so we establish the category of matter.
Thus, we make shift to describe our sensations by the help of these two con-
structs, much as in analytical geometry we describe the location of a point by
means of two coordinates. Energy and matter are, therefore, products of thought
and not, primarily, objective realities. In chemistry, however, we always speak
24 INORGANIC CHEMISTRY
of them objectively. Historically, the order in which these two concepts were
named and defined was the opposite of that in which they stand above. Yet
attempts to organize a conception, corresponding to energy, in response to a need
of which thinkers were conscious, were not wanting before the nineteenth century
opened. To go no further back than the days of phlogiston, we can easily per-
ceive a certain resemblance between this concept and that of energy. The idea
that heat was an " imponderable " had its origin much earlier, and shows the ex-
istence of the same effort to find a second fundamental conception different from
matter.
There is much confusion of thought in many of the current definitions of
energy. For example, it is often said to be " that which causes change in matter. "
This definition is not easy to bring into harmony with common experience in
chemistry. Thus, when heat is applied to mercuric oxide, the change follows.
But with iron and sulphur, the union of the two substances is a condition antece-
dent to the evolution of heat. It is as often true that change in matter causes the
manifestation of energy as the reverse. Matter and energy are on the same plane.
They are conceptions used jointly in describing what we observe. Neither is
secondary to the other. We do not consider any particular one of the coordinates
in geometry as secondary to the other, or as being affected by the other.
The recent theory of chemical potential and of the factors of energy (q.v.) seeks
once more to ascribe the tendency to change (physical or chemical) in matter to
the state of the energy associated with it. It is, therefore, incidentally, so con-
structed as to favor the definition just given.
The definition that "matter is the vehicle of energy " is obviously just as diffi-
cult to harmonize with the above mode of deriving the two conceptions. One axis
is not spoken of as the vehicle of the other in geometry.
The innate desire to reduce our distinct categories to the smallest possible
number may be seen in the history of this subject. The ancients sought the
amalgamation of the two by regarding heat and light as imponderable forms of
matter. In some quarters it is now believed that the one conception of energy is
sufficient, and that matter may be put into the same category as being, for example,
composed of minute particles of electricity (electrons).
The conception of ether was devised because those of matter and energy did
not suffice for the description of all the phenomena of light. It is on the same
plane with matter and energy. Lord Kelvin's effort to reduce the three catego-
ries to two (energy and ether) by assigning the r61e of matter to vortices in the
ether is familiar to students of physics.
Application of the Conception of Energy in Chemistry. — At
first sight it looks as if the statement that energy is conserved is not
applicable in chemistry. Heat and electricity, for example, seem to
be produced and consumed, in connection with changes in composition,
in a mysterious manner. We trace light in an incandescent lamp
back to the electricity, and this in turn to the mechanical energy, and
this again to the heat in the engine. But what form of energy gave
the heat developed by the combustion of the coal under the boiler, or
INTRODUCTORY II 25
by the union of iron and sulphur in our first experiment ? Since we
do not perceive any electricity, light, heat, or motion, in the original
materials, and yet wish to create an harmonious system, we are bound
to conceive of the iron and the sulphur, and the coal and the air, as
containing another form of energy, which we call chemical or internal
energy. Similarly, when heat is used up in decomposing mercuric
oxide, or light in decomposing silver chloride, we regard the energy as
being stored in the products of decomposition in the form of chemical
energy.
The Actual Quantities of Different Kinds of Energy which
may be Obtained from a Fixed Amount of One Kind. — It will
render all the above clearer if we give some numerical illustrations : A kilogram
of water after falling (in a vacuum) 428 meters (about one-fifth of a mile), under
gravity, possesses 428 kilogram-meters of mechanical (kinetic) energy. When the
motion is arrested, the energy of motion is transformed into heat and raises the
water one degree centigrade in temperature. We describe this amount of heat
as 1000 calories (small) ; that required to warm one gram of water one degree
(between 0° and 100°) being called one calorie. A kilogram of any other falling
material would give the same amount of heat (1000 cal.), although, of course,
if its specific heat were smaller than that of water, the temperature to which it
would be raised would be higher, and vice versa.
Here the acting force is the attraction of gravitation, which is a special case.
In absolute units, 1 g. falling 1 cm. generates energy enough to do 981 ergs of work.
So that the thousand grams falling 428 meters, generates
1000x42,800x981=42,000,000,000 ergs of energy.
The erg being so small, we often use the joule (=10,000,000 ergs). This amount
is the same as 4200 joules.
Now, any body of the same mass, moving with the same final velocity, however
set in motion, will also possess the same energy and give 1000 cal. The final
velocity in the above case is V2 gs. = 9,164 cm. per second. The energy of motion
(£ mv2) of one thousand grams of matter moving with this velocity is = £ x 1000 x
(9164)2 = 4200 joules, as before. If the source of mechanical energy were a hot-
air motor (or an engine) of one horse-power, then, since one horse-power repre-
sents a development of 746 joules per second, the 4200 joules of energy would be
produced in about 5J seconds by this means.
If, instead of being turned into heat, all the energy of motion had been con-
verted into electricity, the quantity of the latter would have illuminated a 16
candle-power incandescent lamp for 84 seconds ( = 1.4 minutes). Such a lamp re-
quires the value of 50 joules per second of electricity and, therefore, in 84 seconds
uses up the 50 x 84 = 4200 joules of energy. As an engine of 1 horse-power pro-
duces this amount of energy every 5$ seconds, such an engine, if none of the energy
were lost, could maintain nearly 15 lamps of this kind.
Finally, if the 4200 joules of electrical energy were applied to decomposing
nitrate of silver in ordinary aqueous solution, it would liberate 6| grams (about
i oz.) of silver from combination.
26 INORGANIC CHEMISTRY
Considerations Connected with Chemical Energy : Free
Energy. — These conclusions compel us, for the sake of consistency, to
think of all our materials as repositories of energy as well as of mat-
ter, each of these constituents being equally real and equally impor-
tant. A piece of the substance known as " iron" must thus be held to
contain so much iron matter and so much chemical energy. So sulphur
contains sulphur matter and chemical energy. Thus, by a substance
we mean a distinct species of matter, simple or compound, with its
appropriate proportion of chemical energy.
Since a chemical phenomenon involves a change in the condition of both the
matter and the energy constituents, both have to be considered in describing the
transformation. Thus, in the first of the three experiments (p. 10) we must say that
certain weights of iron matter and sulphur matter, each with a different but defi-
nite quantity of chemical energy, constituting the substances iron and sulphur re-
spectively, are transformed into iron-sulphur matter with the same total weight,
but less chemical energy, constituting the substance ferrous sulphide, with libera-
tion of an amount of energy, chiefly in the form of heat, equivalent to the diminu-
tion in the quantity of chemical energy. As the chemical energy which may be
liberated from the same substance differs in amount according to the material with
which the substance is interacting, the total content cannot be stated. We can.
measure only the amount transformable in a given chemical change, and we call it
the free energy which is available under the given circumstances. The above
action may be represented more compactly, thus :
Iron, free Ch.E + Sulphur, free Ch.E — » Ferrous sulphide + Heat E.
That the ferrous sulphide still contains chemical energy is shown by the fact
that it interacts with many substances with further liberation of heat. At each
step, however, the stock of energy in the material becomes smaller and the chemi-
cal activity less.
The Fotirth Characteristic of Chemical Phenomena. — In the
course of this discussion it has become clear that it is a characteristic
of a chemical phenomenon that, besides a change in the state of the
matter, there is always an alteration in the amount of chemical energy
in the system. This alteration consists in the production of chemical
energy from, or the transformation of chemical energy into, an equiva-
lent amount of some other form of energy.
The absorption or liberation of energy accompanying a chemical
transformation of matter is often, of the two, the more important fea-
ture. We do not burn coal in order to manufacture carbon dioxide
gas. We are glad to get rid of the material product through the chim-
ney. It is the heat we want. We do not employ zinc in batteries
with the object of making zinc chloride or zinc sulphate. So we use
INTRODUCTORY II 27
the electricity, and throw the products away when we refill the jars.
It is the same with burning illuminating-gas or magnesium powder
when we want light, and with eating food, which we do, chiefly, to get
energy to sustain our activity. We do not run electricity for hours
into a storage battery in order to make a particular compound (lead
dioxide, for example), but in order to save and store the energy for
future use. Fully half the total amount of chemical change used in
industry and life is set in motion by us, exclusively, on account of
the energy changes it involves.
As will be seen in the two following sections, observation of the
amount of the energy absorbed or liberated in chemical changes is
also of the greatest importance in the scientific study of chemical
phenomena.
The Progress of Chemical Actions Spontaneous or Forced:
Fifth Characteristic. — It will be recalled that iron and sulphur,
when the interaction has once been started, continue to unite (p. 11)
spontaneously until the chemical change is complete. Similarly, zinc
interacts with dilute sulphuric acid (p. 20) until one of the substances
is exhausted. On the other hand, mercuric oxide continues to decom-
pose only so long as it is heated. Now, the first two actions liberate
energy in the form of heat or electricity, while the latter consumes it.
A study of all chemical changes shows that a distinction can be drawn
on these lines between those which proceed spontaneously and those
which do not. It is, in fact, a characteristic of chemical phenomena
that, in those actions which proceed spontaneously to completion, the
free internal energy in the system diminishes by transformation into
some other form of energy.
This is a particular case of a principle which applies to all changes,
both physical and chemical. All spontaneously occurring phenomena
are accompanied by loss of energy. A stone falls, losing potential
energy, but never rises spontaneously; a body cools, losing heat, etc.
The form of energy liberated in a chemical change which proceeds
spontaneously is most commonly heat. Hence, exothermal actions,
as those which produce heat are called, are usually of the spontaneous
order, and endothermal actions, which absorb heat, generally are not.
There are, however, numerous exceptions to this rule, both in physical
and in chemical phenomena. Thus, water evaporates spontaneously,
although in doing so it absorbs heat. There is a loss of energy in
other ways more than sufficient to offset the acquisition of heat. So,
28 INORGANIC CHEMISTRY
also, very many substances dissolve spontaneously in water and absorb
heat in doing so. Phosphonium iodide (q.v.) decomposes spontaneously
into phosphine and hydrogen iodide, although heat is thereby rendered
latent. Ammonium carbonate (q.v.) spontaneously gives off ammonia,
although in this chemical change heat is absorbed.
In any chemical change which proceeds spontaneously, the energy
liberated may take the form of electricity when the materials are
arranged so as to constitute a battery.
Chemical Activity. — Actions in which there is a relatively large
loss of chemical energy, and, therefore, a considerable liberation of heat
or electrical energy, proceed rapidly ; that is to say, in them a large
proportion of the material is changed in the unit of time. Those in
which less free energy is transformed proceed more slowly. The speed
of the chemical change, and the quantity of energy available because
of it, are closely related. Now, we are accustomed to speak of materials
which, like iron and sulphur, interact rapidly and with liberation of
much energy as "chemically active." Thus, relative chemical activity
may be estimated (1) by observing the speed of a change (see Speed of
chemical actions), or, in many cases (2) by measuring the heat devel-
oped (see Thermochemistry), or (3) by ascertaining the electromotive
force of the current the change gives, when arranged in the form of
a battery (see Electromotive chemistry). These different methods
will be discussed in later sections.
It is evident that the chemical activity of a given substance will
not be the same towards all others. Thus, iron unites much more vig-
orously with chlorine than with sulphur, and, with identical amounts
of iron, more heat is liberated in the former case than in the latter.
With silver, sodium, and many other substances, it does not unite at
all. One of the tasks of the chemist is to make such comparisons as
this (see Specific chemical properties, p. 67). Evidently, the substances
containing the most chemical energy will be in general the most
active.
The "Cause" of Chemical Activity. — The reader will un-
doubtedly be inclined to inquire whether we can assign any cause for
the tendency which substances have to undergo chemical change.
Why do iron and sulphur unite to form ferrous sulphide, while other
pairs of elements taken at random will frequently be found to have
no effect upon one another under any circumstances ? This question
INTRODUCTORY II 29
is so likely to occur to the reader that it should be dealt with at once.
The answer is that we do not know. Questions like this have to go
without answer in all sciences. What is the cause of gravitation?
We know the facts which are associated with the word — the fact
that bodies fall towards the earth, for example — but why they fall
we are unable to say. So, with chemical change, we can state all the
facts we know about it, but even then we cannot say why it takes place.
The words " affinity " and "attraction" are sometimes advanced as if they
supplied some explanation of chemical activity. Now, we have seen that an ex-
planation in science (p. 10) is a description of the details of some process, either
in terms of known facts, or by the use of some imaginary but plausible and help-
ful machinery. Here no facts are known. Even imaginary machinery has not
yet been conceived by any one. So that these terms are words simply, and do
not meet either of the conditions required of an explanation. They are names
for " the tendency to undergo chemical change," and that is all.
All nouns, such as table or book, are general terms applicable to many more
or less various individuals. Some special nouns are used in chemistry. For ex-
ample, affinity names the tendency to undergo chemical change, and distinguishes
this tendency by name from cohesion, or the tendency to unite physically.
Catalysis names a kind of chemical change in which some specific substance must
be present, yet itself undergoes no change. Dissociation names the kind of chemi-
cal change in which decomposition occurs with rise in temperature, and recombi-
nation when the temperature falls. But none of these terms, as such, is an explana-
tion. It does not explain the concussion of two railway trains to name it a
collision.
Of course, if we have some genuine explanation, applicable to all the other
known cases of a class, any newly discovered example falls heir at once to this
explanation. This would be true of a dissociation, where the kinetic theory and
the law of mass action describe the details of all such phenomena. Here the ex-
planation lies, not in the name, but in the knowledge we have of other instances
of the same behavior. The name, of course, suggests the whole theory, if such a
theory exists. But with affinity, or the tendency to enter into chemical action, we
have no theory for any of the samples of the class. We are entirely ignorant as
yet of the details of its mode of operation, equally so in every case, and, in fact,
know nothing at all about it save that affinity exists and that we can measure its
intensity. So the name cannot remind us of any explanation, for none has been
suggested.
As words, the best one can say of them is that they are rather unfortunately
chosen. Affinity suggests kinship, sympathy, or affection. But the suggestion
that such human emotions control the behavior of iron and sulphur is too wild
and too remote from common sense to furnish any assistance. Attraction hints
at some preexisting bond of a material kind which draws the substances together,
for we cannot conceive of action at a distance without some intervening medium
of communication. But we have no other evidence of the existence of an instan-
taneously adjustable harness capable of drawing materials into chemical action.
It is harder to reduce this idea to comprehensible shape than to do without it.
30 INORGANIC CHEMISTRY
If we are sometimes inclined to think that these are more than class-words,
and do suggest some explanation, we have only to carry the same idea further to
be landed in absurdity. Using similarly crude analogies, we might suppose that
the elements were guided by scent, like dogs, or by sight, like birds, or by feeling,
like fish, and so on ad infinitum, arid forget that the fact itself was after all much
simpler than the explanation. Affinity is simply a fanciful name for a real thing.
"Cause" in Science. — The word "cause" was employed in the
heading of the last section, and it will be observed that no cause was
found. This is the invariable rule in physical or chemical phenom-
ena. We kno\7 of no causes, in the sense in which the word is com-
monly employed.
The word has only one definite use in science. When we find
that thorough incorporation of the three materials is needed to secure
good gunpowder, we say that the intimate mixing is a cause of its
being highly explosive. By this we simply mean that intimate mix-
ture is a necessary antecedent of the result. A cause is a condition or
occurrence -which always precedes another condition or occurrence.
Misuse of the word "cause " is frequent. The law of gravitation is not the cause
of the behavior of falling bodies. It is simply a condensed narration of the facts
about falling bodies, and was made long after the first bodies fell. Affinity, or
the tendency to interact chemically, is the imagined antecedent of chemical change.
Such causes, if we call them causes at all, are invented by way of supplying ante-
cedents to things that appear to lack them, that our sense of symmetry may be
satisfied withal. They are occasionally useful. But the less fiction we employ
in the science, the less will be the danger that the student will mistake fictions for
facts, or even fall under the delusion that it is a habit of science to spend more
thought in making gratuitous assumptions than in ascertaining facts.
Of Simple and Compound Substances. — If we place before
a physicist samples of iron, ferrous sulphide, and sulphur, he will re-
port that there are three absolutely distinct substances represented,
because they show three different sets of physical properties. A
chemist, on the other hand, while admitting the accuracy of the re-
port, in view of the criterion used by the physicist, which indeed he
uses himself (cf. p. 5), will insist that there are only two perfectly
distinct kinds of matter in the set, because he can make the second from
matter furnished by the other two. The same sharp contrast in the
points of view arises when mercury, mercuric oxide, and oxygen, or
any similar set of substances, is submitted to the same two tribunals.
In a sense, chemistry reduces the kinds of different matter to a much
smaller number than does physics or any of the other sciences, and
so it is the final authority in all questions involving matter. By the
INTRODUCTORY II 31
chemist, dozens of physically distinct substances are regarded as
closely related because they all can be made with iron, or when de-
composed give it ; hundreds are alike in that sulphur enters into their
composition ; thousands are compounds of oxygen. In fact, the num-
ber of kinds of matter which are perfectly distinct in the strictly
chemical point of view is quite limited.
The conception contained in the last statement was not reached
until centuries of effort had been spent in trying to make gold out of
pyrite (a shining yellow mineral), silver out of lead, and similar
fruitless tasks. The first to put our modern view into definite lan-
guage was Lavoisier in his Traite de Chimie (1789). His investiga-
tions had already entirely overthrown the hypothesis of the phlogis-
tians (p. 9), and had been based upon simple ideas of combination
or decomposition of different kinds of matter. His work showed that
decomposition had its limits. Mercuric oxide could be decomposed
into mercury and oxygen, but no means was found of breaking these
up in turn and producing any fresh substances from them. The kinds
of matter composing these simple materials he named elements. The
element is to be regarded as an ultimate chemical individual just as
the substance is the physical individual. The definition of an ele-
ment is therefore : a distinct species of matter which has not been
shown to be composite.
The caution which prompted Lavoisier to use, as he did, the words
" has yet been," was justified by the fact that several substances, in
his time regarded as elementary, were afterwards shown to be com-
pound. Thus, quicklime was a simple substance until Davy, in 1808,
prepared the metal calcium and showed that quicklime was a com-
pound of this metal with oxygen. Discoveries similar to this have
been made on more than one occasion since.
Until recently a body made up of one or more specific elements had never
been found to yield any simple substance different from those used in preparing
it. In other words, one element had never been turned into another. And this
is still true so far as the familiar elements are concerned Experience with radium
(q.v.), which seems to decompose and give helium, however, may lead to a revis-
ion of the conception of an element. Indeed, even the more radical hypothesis,
that all the so-called elements are made of one fundamental material, has never
lacked supporters. Perhaps the recent discoveries in regard to electrons will lead
to the conclusion, already formulated by J. J. Thomson, that the kinds of simple
matter differ only in the number and arrangement of the corpuscles constituting
their atoms.
32 INORGANIC CHEMISTRY
Element and Simple Substance Not Equivalent Terms. —
We have seen that all substances, that is, physical individuals, must
be thought of as containing both matter and energy (p. 26). It is
certain that iron, if it could be deprived of the energy it loses in com-
bining with sulphur, would be, before combination, as different from the
free metal iron as is the compound ferrous sulphide itself. Now the
development and use of the idea of elements just given, shows that
by the word " element," applied to iron, for example, the chemist refers
only to the matter of this specific variety. It is this part alone which
passes unchanged from the state of metallic iron into combination and
vice versa, or from one state of combination to another. An element is
therefore a kind of matter which never exists alone. It is always
combined with more or less cliemical energy, and often with some
other element or elements as well. A. simple substance, on the
other hand, like all substances, has independent existence, and con-
tains but one element combined with a certain quantity of energy.
A. compound substance, or compound, contains more than one
element together with a certain amount of energy.
That this interpretation of the habit of thought of chemists is correct is shown
by the fact that they speak of both the yellow and the red form as " the element
phosphorus," and also describe phosphorus pentoxide and phosphine as " contain-
ing the element phosphorus." The only thing that is common to all four is, of
course, matter of a specific variety, to which, therefore, the term "element phos-
phorus " must primarily apply. The individuality, marked by certain specific
properties, of the yellow or the red variety or of one of the compounds as a sub-
stance is as much dependent on the energy it contains as on the matter. Hence the
term substawe covers the energy as well as the one or more kinds of matter con-
tained in the material.
Having regard to this distinction, an element was defined above as a " variety
of matter," and not as a substance, and the phraseology used throughout the para-
graph maintains this distinction. No mention was made of energy in the defini-
tion or elsewhere, for, in comparing simple substances and compounds in the
abstract, the only distinction is in the matter. Following each simple variety of
matter into and out of combination, or from one state of combination to another,
is one of the tasks of the chemist. It is also his business to observe the energy
changes. But this is an exercise quite distinct from the other, and constitutes a
separate task in which different methods of observation have to be employed.
The element is thus an abstraction, reached by leaving the energy out of con-
sideration and thinking of the matter only. It is invented for the purpose of
describing the difference between the concretes, the simple substance and the com-
pound substance. To use element and compound as if they were the concrete
things to be contrasted, involves confusion by mixing two distinct categories.
As we have only one name for the element and the corresponding simple sub-
stance, it is well, when there is danger of ambiguity, to call the simple substance,
not " iron," " sulphur," or " the element, " but/ree iron, free sulphur, and in gen-
INTRODUCTORY II 33
eral terms the free element, or the elementary substance. The same kinds of matter
in combination are then the element iron or combined iron, and the element sulphur
or combined sulphur.
• The chemist's work is directed wholly by the thought that the individual ele-
ment (the matter), after combination, is still present in the compound in some form
which is at least guasi-discrete. The readiness of the element to be released once
more under suitable conditions seems to favor this point of view. But with the
energy in a substance it is different. We cannot easily think of the portions of
this, which came in with the several constituents, as being any longer attached to
particular parts of the compound. Indeed, in some actions (chiefly exothermal)
the free energy diminishes and a part escapes, while in others the reverse occurs.
Whatever energy the compound contains, it possesses as a whole, and a state of dis-
tribution between the constituents is seldom taken into consideration. The specu-
lations concerning the relations of the atoms in complex molecules, which have
played a large and useful part in organic chemistry, seem to constitute the only
notable divergence from this point of view.
These remarks are introduced because they help to demonstrate the logical
necessity of the chemist's habit of maintaining a sharp separation, in his mind at
least, between questions involving changes in the state of combination of matter
and questions of energy (see Formulae and equations, p. 57) .
Among the substances which we have been handling, iron, sulphur,
mercury, oxygen, and hydrogen are the free forms of elements. On
the other hand, the substances which we have shown to be composite
are ferrous sulphide, rust, mercuric oxide, silver nitrate, and common
salt. It will be seen that by combination of a limited number of ele-
ments, two, three, or four together, in varying proportions, all the
known distinct substances might easily be accounted for. The list of
elements whose individuality has been established appears upon another
page (Chap, xii) ; of these the larger number are not frequently
encountered. More than 99 per cent of terrestrial material is made up
of eighteen or twenty elements, of which the quantities of the first
eleven, as estimated by F. W. Clarke, are given in the following table :
Oxygen 49.98 Calcium 3.51 Hydrogen 0.94
Silicon 25.30 Magnesium 2.50 Titanium 0.30
Aluminium 7.26 Sodium 2.28 Carbon 0.21
Iron 5.08 Potassium 2.23 99.61
The evidence of the spectroscope shows that the sun and stars con-
tain many of the very same elements as does the earth.
Some of the Fundamental Ideas used by Chemists and the
Corresponding Terms. — To the chemist it is above all important
that he should be able to describe in unambiguous terms the phenom-
ena he observes and the inferences he draws.
To do this he requires some fundamental ideas connoted by suit-
34 INORGANIC CHEMISTRY
able terms. Many of these ideas and terms we have been employing
in order to accustom the reader to their use. It is now advisable to
take them up more systematically.
Any particular specimen of matter, such as a piece of sulphur, a
portion of water, a piece of ferrous sulphide, a fragment of granite, or
some nitrate of silver solution, we call a body. There are thus as
many bodies as there are discrete portions of matter. A body may be
heterogeneous, or made up of visibly unlike parts, as granite and a
mixture of iron powder and sulphur are ; or it may be homogeneous,
or alike in all parts, as are pieces of sulphur and ferrous sulphide and
portions of water and nitrate of silver solution.
Examination of these homogeneous bodies shows that the sulphur,
ferrous sulphide, and water differ from the nitrate of silver solution in
having but one physical component, while the last contains two com-
ponents, nitrate of silver and water, separable by physical means.
The last is like a mixture, only it is homogeneous. Again, the first
three differ amongst themselves, the first being a simple body, with
one chemical constituent, and the two others being compounds having
each two chemical constituents. We speak of the components of a
mixture or a solution because the parts are laid together and retain in
the former case all, and in the latter much, of their identity. But of a
compound we use the word constituents because the parts are built into
each other and have lost their identity.
When a body, say a specimen of sulphur, contains a little of some
other physical component, we speak of it as impure sulphur. This
does not mean that it contains dirt in the ordinary sense of the term.
A little magnesium chloride is a common impurity in table salt (sodium
chloride), and, by absorbing moisture, renders it more moist in damp
weather than it would otherwise become. " Chemically pure " means
that the quantities of the impurities which the material is most apt to
contain have been reduced below the amount which would interfere
with the most exact chemical work for which the substance is com-
monly employed. Absolutely pure bodies are unknown.
By convention we continually speak of "pure" hydrochloric acid,
or of "pure " sulphuric acid, although there may be more than 60 per
cent of water present in the former, and 7 per cent in the latter. By
this we mean to distinguish the former, for example, from " commer-
cial " hydrochloric acid, which contains impurities like sulphuric acid
and a coloring matter in addition to the water. The water is in fact
disregarded, since it is assumed to be present in all cases.
INTRODUCTORY II 35
We distinguish one body, sa}r one piece of sulphur, from another
by its weight, form, or volume. Each particular specimen differs from
every other in these attributes. These are general attributes possessed
by matter and are used in chemistry for measuring quantity.
When all the bodies to which we should apply the name " sulphur "
are compared, we find that, although some are in fine powder, and
others in lumps of various shapes or in crystals, and thus differ in
weight, form, and volume, they nevertheless have many qualities in
common. These qualities we call specific properties, or properties
common to a species. The material composing all the bodies of one
species we call a substance. Some of the specific properties charac-
terizing a substance and common to all specimens of one species are
color, odor, crystalline structure, hardness, melting-point (temperature
of fusion), solubility in water or other solvents, boiling-point (tempera-
ture above which, at 760 mm. pressure, the substance is gaseous), spe-
cific gravity, specific heat, and conductivity for electricity. Thus,
sulphur is yellow, has little odor, crystallizes in the rhombic system,
has a hardness of 2.5 on a scale of ten, has the m.-p. 115° C., is not
perceptibly soluble in water but dissolves in carbon disulphide (41 : 100
at 18°), has the b.-p. 448° C., the sp. gr. 2, the sp. ht. 0.18, and is a very
poor conductor. In the first two chapters, while presenting the ex-
perimental facts required for our discussion, we have had to speak
of substances very frequently (e.g. pp. 26 and 32), and have done so
always in the above sense.
Slight variations from the standard properties of the substance usually indi-
cate the presence of an impurity homogeneously incorporated. The precise ways
in which the properties are affected in such cases will be considered under solu->
tions.
The " substance " will be seen to be of an abstract nature. It is a conception
built up by selecting (or abstracting) the properties common to all specimens.
Hence we classified chemistry as an abstract-concrete science (p. 3). The bodies
under observation are concrete, the classification of the results is under conceptions
of an abstract nature like this one.
That all bodies of a like kind have many identical properties is the most funda-
mental fact in chemistry. Being a general fact, we call it a law, and word it as fol-
lows : The specific properties of a substance are constant in all speci-
mens. Experience having given us confidence in its universality, we take it for
granted in all o^frwork (cf. First characteristic, p. 5).
There are still other qualities which a body (or specimen of mat-
ter) may possess. It has, for example, a certain temperature, press-
ure, motion, or electric charge. These we speak of as conditions of
36 INORGANIC CHEMISTRY
the body rather than properties. In the use of the body they may be
altered, and some of them may be removed or added, arbitrarily.
Thus there are three kinds of qualities to be considered. The
attributes, like weight, form, and volume, do not belong to substances
but to bodies. The specific properties, like color, solubility, and odor,
belong by right to substances, although we sometimes speak of them in
connection with bodies. The conditions, like temperature and motion,
belong to neither, for they can be altered without changing either the
body or the substance.
When the word "properties " is used in speaking of a substance, we generally
refer to the specific properties only, for a substance is by definition the bearer of
nothing but constant, that is, unchangeable qualities. The weight and volume, the
temperature and pressure, are variable, and they do not enter into the conception
of the substance. This distinction sets the abstract nature of the " substance " in
high relief.
Methods of Work and Observation in Chemistry. — It is not
the end of chemical work to make generalizations or laws, like the char-
acteristics of chemical phenomena (pp. 5, 16, etc.), or conceptions, like
those dealt with in the preceding paragraph. These are simply the
means by the help of which chemical work, whether it be investigation,
commercial analysis, or manufacturing, may be carried on more system-
atically. Together they constitute our system for classifying the facts
with a view to ready reference. The sample experiments (pp. 10-14),
if reexamined, will show that we there employed most of the cate-
gories of our classification which have so far been described.
Thus, in the experiment with iron and sulphur (p. 10), it was first
'our object to find out whether the bodies had interacted chemically on
being mixed. To do this we endeavored to ascertain whether any por-
tion of the mixture had acquired new specific properties (p. 35).
Here we used the law (p. 7) that all the properties of the products
of a chemical interaction are different (pp. 5-7) from those of the initial
substances. We also noted the specific properties of the substances
concerned, purposely omitting all mention of quantity and tempera-
ture, because attributes (p. 35) like the former and conditions (p. 35)
like the latter do not characterize substances (p. 35) in general, as dis-
tinct from particular specimens, and cannot be used for identification.
We found that a part dissolved in carbon disulphide and the remainder
was all magnetic.
While all the specific properties, of which a few are mentioned on
p. 35, find application in identification, the first seven iu that list are
INTRODUCTORY II 37
most frequently used. And of these seven the solubility in various
solvents and the boiling-point are by far the most important. This is
because they are best suited for separating mixtures. Thus, here, we
first removed the sulphur by dissolving it, and then, after evaporating
away the solvent, we completed the identification by recognizing the
color and form of the residue. Similarly we recognized by its appear-
ance and magnetic property the iron that was left undissolved.
Separation by the use of one of these two properties, then application
of the others for identification, is the general order of procedure in
chemistry. Most of the other properties cannot be recognized readily
in mixtures) as a moment's thought will show. The general color and
the specific gravity of a mixture, containing unknown substances in un-
known proportions, for example, tell us little about the corresponding
properties of the components. Magnetic properties may be used for
separation, also, but iron is almost the only substance which shows
them markedly, so that their application for the purpose is very
limited.
The use of the boiling-point (temperature of vaporization) for sepa-
ration was illustrated in the evaporation of the carbon disulphide to get
the sulphur alone. This solvent has a very low boiling-point (46° C.)
and therefore a high vapor pressure at the temperature of the room.
By virtue of this it evaporates rapidly. The sulphur (b.-p. 448° C.) is
not volatile under these conditions and remains behind.
In connection with this investigation we employed several of the
common methods of manipulation used by the chemist. These methods
are derived from the conceptions described in last paragraph. Thus
we treated the mixture with a solvent (Fig. 11), on the assumption
that if it was heterogeneous (p. 34) the components would each behave
as if alone present. We then filtered, a method invented for dealing
with a heterogeneous mixture consisting of a solid and a liquid.
Decantation is often used in such cases when the solid is specifically
much heavier than the solvent and settles readily. We allowed the car-
bon disulphide to evaporate spontaneously, and this is our favorite
method of dealing with a mixture which is homogeneous, and therefore
would run through a filter as a whole without suffering separation.
When the liquid has a higher boiling-point than 50-60° C., as water has,
we use heat from a steam-bath or Bunsen flame to promote the evapora-
tion. In evaporation we allow the vapor of the liquid to escape, be-
cause it is the less volatile, dissolved body that we wish to examine.
When we desire, on the contrary, to examine the liquid, the vapor
38
INORGANIC CHEMISTRY
must be condensed. This method, which we have not yet had occasion
to employ, is called distillation (Fig. 16). The jacket round the long
tube is filled with a stream of cold water, which, on account of its high
specific heat, quickly cools and condenses the vapor. The resulting
liquid is caught in a flask.
These methods may be adapted to the investigation of any similar
problem. Thus, gunpowder is made by the intimate mixing of sulphur,
charcoal, and saltpeter (potassium nitrate). If no chemical interaction
whatever has occurred, a sample will be wholly separable into these
components. If a partial change has taken place, a certain amount of
material with different properties will be discovered in the mixture.
If the change has been complete, no portion of the original substances
will be found. We must first study the specific properties of each of
the ingredients separately, in order that a plan of separation may be
devised, and that we may have a basis for comparison with the products
of the separation.
The experiments with mercuric oxide (p. 12) and with silver
nitrate (p. 13) simply ring the changes on the same conceptions.
When we heat the former it is resolved into its constituents. We se-
lected the experiment because the conditions are such that separation
of the products from each other and from the original material occurs
spontaneously without fiirther manipulation. The separation depends
INTRODUCTORY II 39
here on the boiling-points of the materials. That of oxygen is very
low ( — 183° C.) ; so that, being a gas even at ordinary temperatures, it
conies off. That of mercury is much higher (357° C.), but the heat pro-
duced by the flame (1000°-1200° C. in the bottom of the test-tube) is
more than sufficient to vaporize this product also. The speedy con-
densation on the cooler part of the tube, causing separation of the two
products, is due to the wide difference in boiling-points. The part of
the mercuric oxide which is still unchanged, on the other hand, is
entirely involatile even at 120(P C. ; so that it does not mingle with
the vaporized products at all.
In the action of silver nitrate on sodium chloride (p. 13), it was
solubility that furnished the means of separating the products. The
silver chloride is practically insoluble in water, while sodium nitrate is
very soluble. To recover the latter, the boiling-points were then con-
sidered, and the more easily vaporized water was driven off.
When, following the latter experiment, the silver chloride was exposed to light,
resolution into silver and chlorine took place. The products became separated
spatially at once, because chlorine is a gas (substance of low b.-p.) and silver is a
solid (usually a substance of high b.-p.).
In every case the specific properties — color, crystalline form, and so
forth — by which recognition is effected were mentioned.
In connection with the foregoing, the following points should be
noted :
1. That coincidence in two or three specific properties is generally
sufficient to establish identity.
2. That we usually have to separate the components of a mixture
or solution before identification can be effected.
3. That if our problem were the malting of a particular compound
whose constituents were known, we should study the products of each
attempt by the above method. Our object being to get the compound as
a pure substance, we should try to arrange the composition of each of
the initial substances (apart from the necessary constituent) so that any
by-products formed should have .such properties that we could readily
effect a separation of them from the desired compound.
4. That a wide knowledge of specific physical properties is required
for intelligent chemical work (cf. pp. 36-39 and 65).
5. That our methods are purposely limited so as not to separate
chemically combined, but only physically mixed, forms of matter : after
a physical individual has been isolated, and even then only if it has new
40 INORGANIC CHEMISTRY
properties, and is not recognized as a known substance, we next proceed
to separate it into its chemical constituents so as to learn which con-
stituents it contains and in what relative proportions by weight.
Summary. — In this chapter we have added considerably to our
conception of the scope of chemistry (cf. p. 16). Although our survey
is by no means yet complete, we may condense our results as follows :
Chemistry deals with the changes in composition and constitution
which substances undergo and with the transformations of energy
which accompany them. To convert the isolated facts into a science we
classify related parts under laws, such as those of conservation of mass
(p. 17) and of energy (p. 23), and under conceptions, such as those of
chemical energy (p. 25), element (p. 31), body (p. 34), substance (pp.
26, 32, and 35). We also distinguish between attributes, specific
properties, and conditions (pp. 35-36). In the last paragraphs we
have indicated briefly the use to which these conceptions and this
classification are put.
Chemical laboratory work consists largely in the separation, recog-
nition, and description of substances. The importance, especially, of
thorough familiarity with specific properties and the influence of con-
ditions (for example, temperature) to these ends is shown by the ex-
amples (pp. 36-39). The system of classification as a whole is part of
the everyday mode of thought of the chemist, for thought consists
largely in comparing and contrasting, and our system of classification
furnishes the norm of this so far as chemistry is concerned. Learning
chemistry consists, therefore, in learning this classification and becom-
ing habituated to its use.
The influence of conditions has as yet been barely touched. It will
be dealt with more explicitly as occasion offers. The attribute of
quantity, which has already received some attention (pp. 8, 37), will
form the basis of discussion in the next chapter.
Exercises. — 1. Define the following terms, and find illustrations
of each, other than those given on pp. 34-39 : mixture, physical com-
ponent, chemical constituent.
2. Describe (a) a red-hot rod of iron, (&) an aqueous solution of
sugar, employing all the terms given on pp. 34-36 so far as they are
applicable.
CHAPTER III
INTRODUCTORY III
The Law of Definite Proportions, Sixth Characteristic of
Chemical Phenomena. — In making a chemical compound, may we
use varying proportions of the constituents ? The controversy between
Berthollet and Proust, in which the former supported the affirmative
and the latter the negative side of this question, was one of the chief
features of the chemical history of the early part of the nineteenth
century. The ways of forming or decomposing a compound, or of
carrying out a more complex chemical change, may be varied indefi-
nitely. The apparatus, the mode of experiment, and the proportions
of the materials, may be altered at our will. But, in spite of an
enormous amount of careful work, no case of variation in the propor-
tion of the constituents actually used or produced in a given chemical
action has come to light. If too much of one constituent, for example,
is taken, a part simply remains unchanged. A higher temperature
may hasten the chemical action, but it does not affect the quantitative
composition of the products, provided the resulting substances are of
the same nature. It was the work of Stas (1860-65) which settled the
question by proving that even slight variations cannot be detected. It
is, therefore, a characteristic of chemical phenomena that : In every
sample of each compound substance formed or decomposed, the pro-
portion by weight of the constituents is always the same. This state-
ment of fact is known as the law of definite proportions. When the
composition of a siibstance seems to be variable, it is always found on
closer examination that mechanical mixtures of some kind were being
mistaken for pure substances.
Another form of statement, which is a corollary of this one, and is applicable
more directly to complex chemical actions, is: The ratio by weight of any one of
the factors or products of a chemical change to" any other is constant.
The Law of Multiple Proportions, Seventh Characteristic. —
In the course of even a very limited experience in the examination of
chemical ccmpounds, we should be bound to find that among the sub-
stances which we examined were cases of two or move perfectly dis-
41
42 INORGANIC CHEMISTRY
tinct bodies, made of the same elements, but on analysis found to con-
tain different proportions of the constituents. We should discover
that each of those distinct substances had a constant composition pecul-
iar to itself. Thus, we find in nature a beautiful yellow mineral
known as pyrite. A pure sample of this contains nothing but iron and
sulphur, and yet the composition of the substance is entirely different
from that of ferrous sulphide (p. 12). The percentage of the. constit-
uents in each is as follows :
Ferrous sulphide. Pyrite.
Iron 63.59 46.62
Sulphur 36.41 53.38
If we calculate from these data how much of one of the elements is
combined with identical amounts of the other in each of the two com-
pounds, a simple relation emerges from the seemingly unrelated pro-
portions. Thus, taking one part of iron in each, we find 63.59 : 36.41
::!:»(= 0.5725) and 46.62 : 53.38 : : 1 : x (= 1.145). Thus the
quantities of sulphur (0.5725 and 1.145) combined with one part of
iron, in ferrous sulphide and pyrite respectively, are in the ratio 1 : 2.
The reader will find by calculation that if the quantity of sulphur,
instead of that of iron, be fixed at any value, the proportions of iron
in the compounds will then stand in the ratio 2 : 1.
Again, sulphur burns in the air, forming with oxygen a gaseous sub-
stance whose odor is familiar. But a different compound of the two
elements, generally seen in the form of a white fibrous solid, is much
used in dye-factories. In the former, the proportion of sulphur to
oxygen is almost exactly 1 : 1 (50 per cent of each), while in the lat-
ter it is 1 : l£ (40 per cent sulphur and 60 per cent oxygen). Thus
the two different proportions of oxygen combining with one part of
sulphur are in the ratio 1 : 1^ or 2 : 3.
The existence of this simple relation is not an accident, and con-
fined to these cases. It is a rule to which no exception has yet been
found. More complex examples yield the same result.
Thus there are known over two hundred compounds of carbon and hydrogen,
all different in composition. But if we fix the quantity of one of the elements and
find the amount of the other which, in each compound, is combined with that
quantity, the ratio of these amounts.to each other is expressible by integral num-
bers. Take, as a sample, the compositions of four of these substances : methane,
acetylene, ethylene, and naphthalene (moth-balls) contain respectively 3, 12, 6,
and 15 parts of carbon combined with one part of hydrogen.
These ratios are not those of approximately whole numbers. The
more carefully our determinations of the proportions of the constitu-
INTRODUCTORY III 43
•
ents in the several compounds have been made, the more exactly do
integral numbers represent the numerical relations between the results.
This principle was discovered by Dalton (1804), and was embodied
by him in a statement known as the law of multiple proportions,
which ran somewhat as follows : If two elements unite in more than
one proportion forming two or more compounds, the quantities of one
of the elements, which in the different compounds are united with
identical amounts of the other, stand to one another in the ratio of in-
tegral numbers, which are usually small.
The Measurement of Combining Proportions. — The most
exact measurement of the proportions in which the elements combine
to form compounds involves manipulations too elaborate to be gone
into here. Operations of the same nature are described ia works on
quantitative analysis. One or two brief statements, diagrammatic
rather than accurate, will show the principles, however.
If we take a weighed quantity of iron in a test-tube and heat it with
more than enough sulphur (an excess of sulphur), we get free sulphur
along with the ferrous sulphide (p. 12), and no free iron survives.
We may remove the free sulphur by washing the solid with carbon
disulphide. The difference between the weights of ferrous sulphide
and iron gives the amount of sulphur combined with the known
quantity of the latter.
As an example of the study of rusting, we may weigh a small
amount of copper in the form of powder in a porcelain boat and pass
oxygen over the heated metal (Fig. 17). The formation of cupric oxide
takes place rapidly. If we limit the oxygen, part of the copper may
remain unaltered ; if we use it freely, the excess will pass on unchanged.
A given weight of copper cannot be induced to take up more than a
certain amount of oxygen, and use of a less amount simply limits the
amount of copper transformed into oxide. The original weight of
the copper, and the increase in weight, representing oxygen, give us the
data for determining the composition of cupric oxide. The record of
the result is usually made in the form of the quantity of each con-
stituent in a hundred parts of the compound, and is called the per-
centage composition.
The data furnished by one rough lecture-experiment, for example,
were as follows :
Weight of boat empty 3.428 g.
Weight of boat + copper 4.278 g.
Difference = •weight of copper 0.850 g.
44
INORGANIC CHEMISTRY
Weight after addition of oxygen . . . ; *. 4.488 g.
Weight without oxygen 4.278 g.
Difference = weight of oxygen 0.210 g.
The proportion of copper to oxygen, so far as this one measurement
goes, is therefore 85 : 21.
=€E
Fio. 17.
To find the percentage of each constituent, we observe that the pro-
portion of copper is 85 : 85 + 21, or fifa of the whole. That of the
oxygen is ^V^ of the whole. Thus the percentages are :
Copper, 106: 85 :: 100 : x x = 80.2
Oxygen, 106 : 21 :: 100 : x' x' = 19.8
• Naturally, the mean of the results of a number of more carefully
managed experiments will be nearer the true proportion. The percen-
tages at present accepted as most accurate are 79.9 and 20.1.
In the case of mercuric oxide, we may decompose a known weight
of the oxide (p. 12) and afterwards weigh the mercury and ascertain
the oxygen by difference.
Finally, a strip of the metal magnesium may be set on fire in the air.
It gives out a dazzling white light in burning, and on this account the
powdered metal is used in making flash-light powder for photography.
The product is magnesium oxide, a white substance, which partly rises
as a dense smoke and partly falls on the ground. In a loosely closed
porcelain vessel (Fig. 18) the metal may be burned s.lowly, with the
help of the heat from a small flame, and the oxide may be retained.
INTRODUCTORY III
45
The weight of magnesium ribbon taken and the increase in weight due
to oxygen give the data for calculating the proportions of the constit-
uents.
The following figures show the results of experiments in these and
other simple cases, and represent the percentage composition of the
products', only two places of decimals being given in each case. They
will be required for the discussion of the next topic.
(1)
(2)
(3)
Cupric oxide
Copper, 79.9
Oxygen, 20.1
Mercuric oxide
Mercury, 92.69
Oxygen, 7.41
Water
Hydrogen, 11.18
Oxygen, 88.81
(4) Cupric sulphide (7)
Copper, 66.48
Sulphur, 33.51
(5) Mercuric sulphide (8)
Mercury, 86.18
Sulphur, 13.81
(6) Hydrogen sulphide (9)
Hydrogen, 5.92
Sulphur, 94.07
Cupric chloride
Copper, 47.3
Chlorine, 62.7
Mercuric chloride
Mercury, 73.8
Chlorine, 26.2
Hydrogen chloride
Hydrogen, 2.76
Chlorine, 97.23
The Law of Combining Weights, Eighth Characteristic. — The
ratios in the above list represent the true proportions by weight in the
various compounds, but natu-
rally the individual numbers
constituting those propor-
tions have no chemical sig-
nificance whatever. They are
arbitrary values selected so
that the constituents of the
proportion may together make
100. Each pair represents
the constant ratio which is the
mean result of numerous ex-
periments.
We begin the effort to re-
duce these numbers to order
by selecting one element as
our starting-point, and by tak-
ing some convenient weight
of it as the basis. As we
shall see, it makes no difference what choice we make in either re-
spect. To avoid waste of time, we shall, therefore, use oxygen, as it is
the element generally preferred by chemists for the purpose. The
reason for this preference will be apparent later (see pp. 48, 51).
FIG. 18.
46 INORGANIC CHEMISTRY
We should naturally be inclined to use 1 part of the element as our
basis. But our later steps involve finding out what amounts of the
other elements combine with this quantity, and we perceive that the
amount in the case of hydrogen will be only 0.126 parts. We calculate
this from (3) : 88.81 : 11.18 : 1 : x (= 0.126). If, however, we take
8 parts of oxygen, this amount of hydrogen is also increased eight
times and becomes 1.008. As no element is found to combine in
smaller proportions than hydrogen, we are satisfied that a scale for our
numbers based on 8 parts of oxygen will not involve any values less
than 1. The choice of scale is purely one of convenience.
Our further procedure is determined by a desire to demonstrate a
particular relationship, and so we must keep a fixed principle in mind.
This principle is that of calculating a series of combining proportions
of which each succeeding member is the amount of some element which
combines with the amount of the one preceding it in the list. Oxygen
is the initial member of the series.
From (1) we calculate the amount of copper combined with 8 parts
of oxygen, thus: 20.1 : 79.9 : : 8 : x (= 31.8). We next look for an ele-
ment which combines with copper — any such, whether it appeared in
the above list or not, would do — and from (4), for example, find, by
the same method of calculation, that 31.8 parts of copper unite with
16.03 parts of sulphur. Then we select any element that combines in
turn with sulphur, and from (5) learn that 16.03 parts of sulphur unite
with 100 parts of mercury. Next we observe from (8) that 100 parts
of mercury combine with 35.45 parts of chlorine. Finally, from (9) we
calculate that 35.45 parts of chlorine unite with 1.008 parts of hydro-
gen. Setting these results down in order, we obtain the series:
OXYGEN COPPER SULPHUR MERCURY CHLORINE HYDROGEN
8 31.8 16.03 100 35.45 1.008
Since each of these numbers represents the amount of the corre-
sponding element which is actually found to combine completely with the
quantities of two other elements, we may call them combining weights.
It must be observed, however, that those of our measurements of which
we have made use in deriving them relate each number in the series to
its two immediate neighbors only.
Now (2) gives the proportion in which oxygen combines with mer-
cury. By calculation, 7.41 : 92.59 : : 8 : x (= 100), we find that 8
parts of oxygen combine with 100 parts of mercury. It must be
pointed out explicitly that this result for mercury and oxygen is en-
INTRODUCTORY in
47
tirely independent of the previous one for mercury and sulphur, by
which the value 100 was originally derived. It is based upon the in-
dependent measurement of the proportion of mercury in mercuric oxide,
and we could not possibly have foretold what this proportion would be.
It appears then that the above combining weights, at first measured for
neighboring pairs only, are also the weights according to which remoter
combinations take place. Further study confirms this speculation.
Thus (3) shows that oxygen and hydrogen combine in the proportion
8: 1.008, and (7) gives the proportion in cupric chloride, and (6) that in
hydrogen sulphide exactly as they stand in the series.
A few other pairs in this set are capable of uniting. For example,
oxygen and sulphur, as we have seen, form two compounds in which
the proportions are found to be 2 x 8 : 16.03 and 3x8: 16.03. Sulphur
and chlorine unite in the proportions 16.03 : 2 x 35.45 and 2 x 16.03 :
35.45. Sulphur and copper form another compound, cuprous sulphide,
and in it the proportion is 16.03 : 2 x 31.8, instead of 16.03 : 31.8.
These examples are also, incidentally, illustrations of the law of mul-
tiple proportions (p. 43).
These relations become clearer when represented diagrammatically.
Thus the five elements, omitting oxygen, give the following compounds
and no others :
Copper -
(31.8)
-1:1-
-1:2-
•1:1-
-Sulphur-
(16.03)
-2:1-
-1:1— Mercury— 1:1— Chlorine — 1:1— Hydrogen
(100) (35.45) (1.008)
-2:1-
-1:2-
-2:1-
-1:1-
-2:1-
It will be observed that hydrogen forms stable compounds with only two of the
other four elements in the series. If, however, oxygen had been included, compounds
of oxygen with all the other five would have demanded recognition. This illustrates
the reason given below for the preference of oxygen as the fundamental element.
It ought to be added that combinations of three or more elements are not un-
common, but neither does examination of these cases reveal anything in conflict
with the principle which our study of the combining weights is bringing to light.
For example, oxygen, chlorine, and hydrogen combine in the proportions :
2x8: 35.45 : 1.008, 6x8: 35.45 : 1.008, and 8 x 8 : 35.45 : 1.008.
48 INORGANIC CHEMISTRY
A complete study of all known combinations of all the elements
shows that no compound containing any two or more of the elements is
known, whose composition cannot be expressed in terms of the combining
weights found in a series like the above, by using proper integral fac-
tors when necessary.
This statement would remain true whatever element we chose as
the initial one of the series (see Exercise 1, p. 52), whatever combining
weight we assigned to that element, — for this affects only the scale of
all the members, not their relative values, — and whatever the order in
which the succeeding elements were taken. It applies to all the ele-
ments without exception, for the above series might have been extended
to include all the known elements. It represents, in fact, not a prop-
erty of our method of manipulation, but of chemical combination itself.
It will now be seen that, since the order is a matter of indifference,
a long series like the above is not needed, and would lead only to cumu-
lative errors. The exact determinations of the combining weights of
most of the elements have actually been made by direct union with oxy-
gen or with the help of but one intermediate step. Again, if the
question had been one of mathematics, hydrogen, the element with the
lowest combining proportions, would have furnished the basis and unit
of the scale. But the question was the practical one of getting the most
accurate measurements for the relative magnitudes of the numbers, so
oxygen was chosen instead. Nevertheless, the value 8 was selected in
order that the advantage of having a mathematical unit, or something
close to it, in the combining weight of hydrogen, might be retained also.
The law of combining weights may be put briefly thus : The pro-
portions by weight in which all chemical combinations take place can
be expressed in terms of small integral multiples of fixed numbers
called combining weights, one for each element. It describes what is
perhaps the most striking of all the characteristics of chemical action.
It is, perhaps, hardly necessary to point out that the laws of definite and mul-
tiple proportions are simply partial statements whose whole content is included in
this far greater generalization. No one chemist succeeded in discovering this
property of combining weights. The work of J. B. Richter, Dalton, and many
others contributed to it.
Witnout this fact, the remembering of the compositions of chemical
substances, necessary as it is to the chemist, would have been com-
pletely beyond the power of any ordinary memory. With it, the task
becomes comparatively simple. It is only necessary to decide on the
best system of values for the combining weights, and then, regarding
INTRODUCTORY III 49
the value of this for each element as the unit of weight for that element,
to express the proportions of the element in every compound by the
proper multiples. Thus, given a list of the combining weights, one
for each element, only the small integral multiples have to be kept in
mind in connection with each compound.
The reader will require a little time, however, before he becomes
accustomed to the use, not of a single unit of weight, but of a different
one for each element. Chemistry is the only science in which the
physical unit of weight, which is the same for all materials, is not
employed for every purpose. The physical manipulations of the
chemist are carried out with the use of physical units, but the chemi-
cal results are expressed in terms of individual unit quantities of the
several elements, the combining weights.
The individual units actually iised for each element are not in all
cases identical with those we have given. The final values will be
discussed in the next section but one.
Most of the first exact determinations of combining weights were made by
Berzelius before 1830. It should be added that, while the combining weights, with
the exception of that of oxygen which is the standard, are never actually whole
numbers, although they often approach such integral values, the integers used to
multiply them, when they are employed to express combining proportions, are so
most exactly. Even in determinations by methods of the highest refinement, the
factors to be used in multiplying the combining weights are always found to di-
verge from whole numbers by amounts within the known errors of the method of
measurement.
Equivalents. — The combining weights may be viewed from an-
other standpoint. The relation between the first three members of
the series discussed in the last section, for example, may be stated
thus : 16.03 parts of sulphur and 8 parts of oxygen combine with iden-
tical amounts of copper, namely, 31.8 parts. Either of them will satisfy
the same amount of this third element. In fact, they are equivalent
for the purpose of combination. Similarly, 31.8 parts of copper and
100 parts of mercury are equivalent, because either will hold 16.03
parts of sulphur in combination. •
Now, copper will actually displace mercury from combination with
a third element. The copper will combine with the third element, and
the mercury will be set free. So, if we know the combining weight of
copper, but not that of mercury, we can make a measurement of the
amount of mercury displaced by 31.8 parts of copper. This amount
(100 parts), which we then call the " equivalent," will be also the re-
50
INORGANIC CHEMISTRY
quired combining weight. It will be so because this amount of mercury
was formerly in combination with that quantity of the third element
with which 31.8 parts of copper are now united. The terms combin-
ing weight and equivalent both refer to the same members, and the
latter simply emphasizes a particular experimental method. Follow-
ing the custom of chemists, we shall hereafter refer to the combining
weight on the scale oxygen = 8 as the equivalent of the element. The
equivalent weight of an element is the quantity of the element -which
combines with or displaces 8 parts of oxygen or 1.008 parts of hydrogen.
Atomic Weights. — The chemist frequently uses the idea of
equivalents and the values we have given them. But far more often
he employs a slightly differing set of numbers, which, for reasons that
will appear in a subsequent chapter, he calls atomic weights. The fol-
lowing list shows the elements whose equivalents we have been dis-
cussing, along with one or two others, added by way of furnishing a
fair sample, and gives both sets of weights for the purpose of com-
parison :
EQUIVA-
EQUIVA-
ELEMENT.
LENTS (COM-
BINING
ATOMIC
WEIGHTS.
ELEMENT.
LENTS (COM-
BINING
ATOMIC
WEIGHTS.
WEIGHTS).
WEIGHTS).
Oxygen .
8
16
Iron .
27.95
55.9
Copper .
31.8
63.6
Magnesium
12.18
24.36
Sulphur
16.03
32.06
Carhon . .
3.00
12.00
Mercury
100.0
200.0
Aluminium
9.03
27.1
Chlorine
35.45
35.45
Sodium .
23.05
23.05
Hydrogen
1.008
1.008
Bromine
79.96
79.96
It will be seen that some equivalents have been multiplied by two,
the first four and those of iron and magnesium, for example ; some
have been multiplied by three, like that of aluminium ; some by four,
like that of carbon ; and some remain unchanged, like those of chlo-
rine, hydrogen, sodium, and bromine.
The reasons for this manipulation of the simple equivalents found
by experiment is based upon theoretical considerations. As it is im-
possible, until we reach certain important facts which cannot be intro-
duced here, to explain these considerations, the discussion of the rea-
sons for the changing of the numbers will be postponed until after
these facts are before us. Suffice it to say that great advantages are
found to attach to these modifications in the values. The step from
equivalents to atomic weights (q-v.) is taken before the justification of
INTRODUCTORY III 51
it can be given, because otherwise formulae (see next chapter) could
not be used in the earlier chapters, and so the advantages their em-
ployment offers would be sacrificed.
A little thought will show that the atomic weights have all the
properties which we have shown to belong to the combining weights
(equivalents). The atomic weight is the unit of weight (p. 49) actu-
ally used in expressing the proportions of each element in all its
compounds. The integral factors are, of course, different from those
which would be employed in expressing the composition of the same
substance in terms of equivalents, because many of the latter have
been multiplied by small integers already in course of being made into
atomic weights. But the multiplication has in every case been by an
integer, so that no change in the properties of the numbers has oc-
curred.
To the reasons given above for the choice of oxygen as the funda-
mental element, and the value 8 for its combining or equivalent weight,
one other may now be added. The majority of the atomic weights,
calculated on this basis from the experimental results, fall so close to
being integers that the nearest round numbers are exact enough for
ordinary use. Thus in the above list nine of the twelve atomic weights
are within 0.1 of the nearest whole number. This convenience disap-
pears when, for example, hydrogen with the value 1 is made the basis.
The reader will inevitably find difficulty at first in thoroughly
grasping the significance of these numbers. It may, therefore, be of
some assistance if a hint is thrown out which will suggest a concrete
basis for this curious property. These numbers appear to mean that,
when we wish to make a chemical compound, we may choose any two
elements from the list, and, if it is found that they can combine at all,
we have only to take the atomic weights, worked out from other com-
binations of each element, and we shall find that they will exactly
suffice for this case of chemical union. If complete combination of
both materials does not take place, then trial will quickly show what
multiples of the atomic weights will result in this. The situation
seems to suggest that the constructing of chemical compounds depends
upon the putting together of ready-made " parts," like those of a
watch or a bicycle. The parts seem to be " interchangeable," and
each element seems to be furnished to us by nature in ready-made
packets suitable for application in building up any chemical structure.
A complete list of atomic weights is printed on the inside of the
cover at the back of this book.
52 INOKGANIC CHEMISTRY
Summary. — This chapter adds an important item to our state-
ment of the scope of the science (cf. p. 40), which, therefore, now
reads as follows : Chemistry deals with the quantitative study of the
changes in composition and constitution which substances undergo and
with the transformations of energy which accompany them. To ex-
press the quantitative relations which are observed, a different unit of
weight is employed for each element, and is known as the atomic
weight of the element.
There are other important characteristics of chemical phenomena
mostly concerned with the conditions (p. 35), but those which have
been given are sufficient, for the present, to guide us in the systematic
study of the behavior of the elements and their chief compounds.
It may not be out of place to indicate which are the most important conditions.
The first condition whose influence we are likely to notice in chemical work is
that of temperature. The accelerating effect of rise in temperature on the speed of
all chemical changes (see Chap, v), and van 't Hoffs law (q.v.) in regard to the
effect of temperature on the direction of chemical change, describe the most impor-
tant characteristics of this influence.
The second condition whose effects we continually observe is that of concentra-
tion. This, and not chemical affinity, as many suppose, determines chemical be-
havior in the majority of familiar actions. It is described by the law of concentra-
tion (q.v.), or " mass action," as it is often inappropriately called to the great
detriment of clearness. Erin's method of obtaining oxygen furnishes the first con-
spicuous case of the influence of this condition which we shall encounter. If this
and many other examples are passed over without discussion, it is only because
we must wait until much chemical experience has been gained before this principle
can be understood. Pressure is the special name for concentration in gases.
A third condition of great importance in many — perhaps most — chemical
actions is the presence of a catalytic agent (q.v.).
Exercises. — 1. Starting with mercury as the basal element and
100 as its combining weight, calculate from the data on p. 45 a series
of combining weights for the other five elements. Then show that this
series has the same properties as that discussed in the text.
2. There is another oxide of copper, namely, cuprous oxide, which
differs in properties and composition from cupric oxide. In it the
ratio of copper to oxygen is 2 x 31.8 : 8. If this compound had been
used in deriving the combining weight of copper on p. 46, what effect
would this procedure have had on the rest of the series ? Would the
properties of the series have been affected ?
3. Calculate the equivalent (combining) weight of iron from the
composition of ferrous sulphide (p. 42).
CHAPTER IV
INTRODUCTORY IV
A CONSIDERATION of the contents of the foregoing chapter will show
that the complete description of a chemical change must be exceed-
ingly involved. In a moderately complex action, such as that of
sodium chloride upon silver nitrate, we should say that sodium chlo-
ride, composed of one atomic weight each of sodium, and chlorine, when
brought in contact with silver nitrate, composed of one atomic weight
each of silver and nitrogen and three atomic weights of oxygen, gave
silver chloride, composed of one atomic weight each of silver and
chlorine, and sodium nitrate, composed of one atomic weight each of
sodium and nitrogen and three atomic weights of oxygen. Such a
statement, while it would give all the facts from the quantitative
point of view, would be difficult to grasp and lacking in perspicuity.
Symbols, Formulw, and Equations. — In order to represent the
nature of a chemical change in a form which may be taken in at a
glance, the chemist is in the habit of using certain symbols, first intro-
duced by Berzelius. Thus, the letters Ag represent one atomic weight
(i.e., 107.93 parts) of silver (argentum), and O represents one atomic
weight (i.e., 16 parts) of oxygen. In other words, the symbol of an
element means one chemical unit weight of the element. Since many
elements begin with the same initial, two letters have frequently to be
used to distinguish them. When the names of the elements are not
the same in all languages, resort is frequently had to Latin. Thus, Cu
stands for one combining weight of copper (cuprum), Fe is used for
iron (ferrum), Hg for mercury (hydrargyrum). Prom German we have
Na for sodium (natrium^) and K for potassium (kalium}. To repre-
sent a compound, the symbols of the elements which it contains are
placed side by side, small numbers indicating multiples of the atomic
weights where they occur. Thus, sodium chloride is represented by
the symbols NaCl, silver nitrate by the symbols AgK03. A combina-
tion of symbols is called a formula. The symbols of a formula, taken
by themselves, do not stand for any definite quantity ; each is one
53
54 INORGANIC CHEMISTRY
factor of a proportion. Ag means the proportion of 107.93 parts of
silver to the proportions of the other elements represented by the other
symbols which may be connected with it.
The value of these symbols lies not only in the compact way in
which the resulting formulae present the composition of compounds,
but also in the use which may be made of them in showing at a glance
the details of a chemical change. The chemical action just mentioned
appears as follows :
NaCl + AgN03 -> AgCl
This expression contains all that was conveyed by the words which
were written out in full above. The arrow indicates that the materials
on the left-hand side pass, in the chemical transformation, into those
on the right-hand side. Such symbolic expressions are called
equations.
One other variation is in frequent use. If we could make pyrite
(p. 42) by the union of iron and sulphur, we should require two
atomic weights of the latter to one of the former. The equation
would run thus :
Fe + 2S -> FeS2.
It will be observed that we employ the form 2S before combination
and S2 (in FeS2) after it. The reasons for this usage will become clear
as we proceed. We note simply that 28 means 2 separate atomic
weights of sulphur, as 3FeS2 would mean three separate formula-
weights of pyrite. The same substance (cf. p. 32), sulphur, might
appear as 5S or 8S in other equations, according to the proportion
needed. But FeS2 is a group of three atomic weights united chemi-
cally. The substance pyrite never contains anything but two atomic
weights of the element sulphur, and its formula is invariable. Thus
the regular integers multiplying the atomic weights in the composi-
tion of a particular compound are written after the symbols of the
elements, while more arbitrary factors which change from one use of
the substance to another are written in front.
We shall find later that there are two substances containing nothing but
oxygen, and that each is a compound of the element with itself. The molecular
formulae of these two are O2 (oxygen) and O3 (ozone). Thus O or 2O or 3O
would all be used for different proportions of the substance 0, if such a substance
were known, and 0 would be used for a substance made of oxygen, but different
from oxygen or ozone. Molecular formulae will not be employed here until after
Avogadro's hypothesis has been discussed. They will then be used exclusively.
The object in writing a series of formulae in the above manner is to show
INTRODUCTORY IV 55
that the system upon the left-hand side, consisting of certain substances whose
composition and properties we know, is, under the conditions of the experiment,
unstable, and changes into the system upon the right-hand side, whose nature
we also know. The materials on the two sides are essentially different, for the
transformation represented is a chemical change. It is somewhat anomalous,
therefore, that, to connect two sets of things which are essentially different, the
sign = is usually employed. To call this a chemical equation is still more anoma-
lous, since it is precisely in the chemical point of view that the difference between
the two sides is most strongly to be emphasized. It represents two sets of things
which are different and not alike chemically. The physical properties of the two
sets of substances are likewise totally unlike. There is only one respect in which
the materials on the two sides agree, and that is that their mass is not different.
This is, however, merely an example of the law of conservation of matter, and
need not, therefore, be specially commemorated in the form in which we write
every equation. It may be assumed that the equality in mass holds for all chemi-
cal changes until some case where it does not hold shall have been discovered.
Above all it must be remembered that the chemical equation is not an algebraic
expression ; it is subject to none of the rules of algebra. It is a brief expression,
in terms of the combining weights, of the distribution in kind and quantity of the
constituents of a system before and after chemical change.
Making Formulae and Equations. — The making of a chemical
equation, it is needless to say, involves (1) a knowledge of the pro?
portion of the constituents in each compound used and produced.
These are ascertained* by experiment, and are at first stated in ordinary
physical units of weight. We require (2) a knowledge of the chemical
unit weights — the atomic weights — which have been accepted by
chemists for each element concerned, since the proportion *is to be
expressed in terms of the chemical units of weight proper to each ele-
ment. By factoring the terms of the first proportion so that one factor
in each case is the atomic weight, we discover whether multiples of
the atomic weights will be required to represent the composition of the
substance, and if so what these must be. We may then introduce the
symbols in place of the atomic weights, and construct the formula for
each compound. It is necessary (3) to know the proportions in which
the various compounds concerned are used and produced. These give
us the multiples of each formula as a whole, and we are then in pos-
session of everything required to make the equation.
It will be seen that not all of these facts are independent of one
another, so that a complete measurement of everything is not in prac-
tice required. Thus, knowing the total number of atomic weights of
oxygen used on the left-hand side of an equation, we see from the
laws of the conservation of mass and matter that this number must
appear on the other. So that if, for example, but two products are
56 INORGANIC CHEMISTRY
formed, and only one of them contains oxygen, we obtain by inference
some definite information about the composition of this substance. In
other words, in obtaining the complete information needed in making
equations, we may frequently reduce the actual work of measurement,
and reason out the minimum number of proportions that must be
measured, by a careful study of the nature of each change. Then, use
of the law just mentioned and of our general knowledge of the nature
of the action will enable us to complete the equation. It will be seen
at once that since we have no means of knowing a priori what the
proportions in the resulting compound will be, if we make an experi-
ment with unmeasured quantities of the substances, we must lack the
data for making the formula. It is, therefore, absolutely impossible to
construct a chemical equation unless we have obtained by experiment
and by inference all the information required under the heads enum-
erated above. Placing the symbols of the elements side by side, for
example, does not suffice, since we do not know what multiples of the
atomic weights may be involved. «
Let us take an actual illustration which will make all this clear.
Suppose the problem is to make the formula of dried rust and the
equation representing its production from iron, and that very refined
experimental methods are not used. By weighing before and after the
change we get the weight of the iron and of the corresponding amount
of oxygen in the rust it produces. If we took 2 g. of iron we should
get about 2.86 g. of rust. So that the proportion of iron to oxygen is
2
jr-^-r- The atomic weights are 55.9 and 16 respectively. Since these
O.oo
have been selected so as to conform to the law of combining weights
(p. 48), this proportion must be expressible in the form - — — '— J
(/ /{ -L O . w
sc
in which - is a ratio of two small integers. To find the value of this
y
ratio we have,
x x 55.9 _ 2
y x 16.0 ~ O86'
, . , x 2 x 16.0 2
from winch, - = g-gg-^g = 0.65, or about g -
Substituting these values for x and ?/, and the symbols for the atomic
2 x Fe
weights, the proportion appears as = — '—^- > and the formula of the
o X U
INTRODUCTORY IV 57
compound is Fe203. We perceive at once that the equation must be,
2Fe + 30 — > Fe203,
since we must take enough material to make the product. In this case
one measurement (the ratio of iron to rust), the knowledge of the
atomic weights (which, of course, depend upon the experimental work
of chemists), and the law of conservation of mass, constitute the
necessary and sufficient equipment for learning all about the change
quantitatively, and, therefore, for making the equation.
When more than two constituents make up a compound, the same
principle is used. Thus, if we find a substance containing 27.09 per
cent of sodium, 16.50 per cent of nitrogen, and 56.41 per cent of
oxygen, we resolve each percentage into the desired factors by divid-
ing it by the corresponding atomic weight. This gives us :
PERCENTAGE Ax. WT. FACTOR USING SYMBOLS -j- 1.175
Sodium, 27.09 = 23.05 x 1.175 or Na X 1.175 Na x 1
Nitrogen, 16.50 = 14.04 x 1.175 or N x 1.175 N x 1
Oxygen, 66.41 = 16.00 x 3.526 or O x 3.526 O x 3
Now the composition of all known compounds can be expressed by in-
tegral multiples of the atomic weights. We must, therefore, find the
whole numbers which stand in the same ratio as the three fractional
factors. We may divide all the factors by 1.175, since the composi-
tion of the compound is a question of ratio simply. This gives the
integral numbers 1, 1, and 3, and the formula NaN03.
It is hardly necessary to add that a chemical equation gives the
proportions of the materials and nothing more. The physical condi-
tions, for example, whether the substances are dissolved in a liquid, or
are in the state of gas, or are at a high temperature, have no place in
it. Yet it is absolutely necessary that these facts should be known
also. The physical properties of the substances concerned, and also
the energy in the form of heat or electricity which may appear or dis-
appear in the process, are likewise left entirely out. A question in
regard to the nature of a particular chemical change demands in
answer a full statement of all these things. The equation is therefore
an essential part, but only a part, of such a statement.
That the formulae and equations can deal only with the material part of the
substances undergoing change, and not with their energy (p. 26), is shown by a
moment's consideration. Consistency is to be secured only by holding that the
symbol S, whether alone or combined with others, stands for the matter-part of the
sulphur (for the element, in fact, see p. 32). It is 32 parts by weight of sulphur-
58 INORGANIC CHEMISTRY
matter. Only in this way does it preserve the same significance on both sides of
the equation S + O2 — -> S02. If S on the left side stood for the free substance sul-
phur, then it would stand for 32 parts of sulphur-matter plus the appropriate
amount of energy. In this case the S on the right side would have a different sig-
nification, and represent a less amount of energy. This is only the beginning of
the difficulty, for we then find that S in H2SO4 represents the same weight of
sulphur-matter with still another proportion of energy, and S has as many inter-
pretations as there are formulae in which it occurs. Clearly, all references to en-
ergy should be rigidly excluded from equations, and thermochemical data can never
be given in connection with them without complete sacrifice of consistency. In
this book, however, the habit of writing thermochemical equations, being universal,
is frequently followed when thermochemical data are given.
Units of Measurement in Chemical Work. — In chemical work
temperatures are invariably measured on the Centigrade scale. The
temperature of a mixture of ice and water is the zero point. The tem-
perature of the steam which rises from water boiling under a pressure
of one atmosphere is represented by 100°. The interval between those
two points is divided into one hundred equal parts.
For the expression of length, weight, and volume, the metric system
is employed. The unit of this system is the meter, which is subdivided
into decimeters, centimeters (cm.), and millimeters (mm.). For small
measurements the last subdivision is taken as the unit. A cubic centi-
meter (c.c.) is the unit of volume for small measurements. For larger
ones the liter, which contains 1000 cubic centimeters, is used. The
unit of weight is that of one cubic centimeter of water at 4°, the tem-
perature of maximum density. This is called the gram.* For larger
amounts of material the kilogram, which contains 1000 grains (1000 g.),
is frequently employed. The meter is equal to about 39J inches in
ordinary measures, and the centimeter is very nearly § of an inch.
One liter is about ^ of a cubic foot and contains 61 cubic inches or
35 fluid ounces. One hundred grams is about 3^ ounces avoirdupois,
and one ounce equals 28.35 grams.
Calculations in Chemistry. — In the laboratory it is frequently
desirable that we should know what amount of some substance may be
obtained by a given chemical action from another, or what amount of
material must be used to obtain the desired amount of some product.
This information is readily accessible, since measurements of quantity
* In point of fact, the grain is the one-thousandth part of the weight of the
standard kilogram kept in Paris. This differs from the weight of 1 c.c. of water at
4° by less than 0.01 per cent.
INTRODUCTORY IV 59
in connection with most chemical changes are 011 record. The simplest
and most easily handled form of this record is found in the formulae
of compounds, and in the equations representing the changes which
they undergo. It is most convenient, therefore, when a question of
this kind occurs, to ascertain and write down, first, the equation.
Having then before us the information in regard to the quantities in
the most condensed form, we may use such parts of this information
as are required for the problem in hand.
Suppose, for example, that the question is in regard to the weight
of oxygen which may be obtained from 120 g. of mercuric oxide. We
write down the equation, and, if the numbers are not familiar to us,
we ascertain the atomic weights. These we place then below the
symbols by which they are represented, thus:
HgO -> Hg + 0
200 + 16 200 16
Reading this equation, it appears that one formula-weight, or 216 parts
by weight, of mercuric oxide give one atomic weight, 16 parts, of oxy-
gen, and the question is, What weight of oxygen will be obtained from
120 grams of the oxide ? The answer may be stated by simple pro-
120
portion : 216 : 16 : : 120 : x, or 16 x 7^ = Answer. The reader
must conquer a tendency to speak of the symbol 0 as representing "1
part" of oxygen : it stands for 16 parts. The word "part" refers to
physical units exclusively.
In the very simplest cases it might not seem necessary to write out
the equation as a preliminary step, but, wherever the action is more
complicated, there will be great danger of errors arising if this is not
done. It will be noticed that not all the data which we have written
down are necessarily used. In general, only two of the three or more
weights which the equation represents will be required. The first
thing after writing down the equation is to be sure that the question is
brought into relation to the information which the equation gives, and
is perfectly understood. Any obscurity on this point is sure to lead to
miscalculation.
Exercises. — 1. What weight of sulphur is contained in 100 g.
of pyrite ?
2. What weight of rust may be obtained with a supply of 10 g. of
oxygen ?
60 INORGANIC CHEMISTRY
3. How much silver is contained in 100 g. of an impure specimen
of silver chloride which is 33 per cent sand ?
4. What are the formulae of the substances possessing the follow-
ing percentage compositions ?
i ii in
Magnesium, 25.57 Sodium, 32.43 Potassium, 26.585
Chlorine, 74.43 Sulphur, 22.55 Chromium, 35.390
Oxygen, 45.02 Oxygen, 38.025
5. What are the percentage compositions of substances possess-
ing the following formulas ? Mn304, KBr, FeS04.
CHAPTER V
OXYGEN
Historical and Introductory. — Almost one-quarter of the at-
mosphere, by weight, is free oxygen. Water contains nearly 89 per cent
of oxygen in combination, and this element constitutes about 50 per
cent of common materials like sandstone, limestone, brick, and mortar.
On account of its predominance over other elements in quantity, and
the exceptional capacity which it exhibits for forming compounds with
a great variety of other elements, the systematic study of chemistry,
may conveniently be begun with oxygen.
While many elements which are less easily obtainable than oxygen
have been recognized as distinct substances for many centuries, oxygen
did not attain this position until it was
first prepared by Priestley in 1774.
The reason of this was that gases are
not so easy to handle and distinguish as
solids or liquids, and consequently very
slow progress was made in the study of
them. Priestley was particularly in-
terested in examining the nature of the
gases which were evolved by some ma-
terials when heated. His plan was to
fill an elongated glass vessel with
mercury (Pig. 19) ; to invert this in a
trough filled with the same metal, and,
after allowing the substance under ex-
amination to float up into the top of the
tube above the mercury, to expose it to
the rays of the sun concentrated by a large burning lens. Priestley
found that one material, then known as "mercurius calcinatus per se"
(mercuric oxide), gave off an unusual amount of a gas, or "air" as he
called it. He found that this gas supported combustion extremely well
and, later, that it was respirable and favorable to the life of small ani-
mals, such as mice. He did not, however, recognize that atmospheric
61
FIG. 19.
62
INORGANIC CHEMISTRY
air was a mixture, and that the substance he had obtained was in reality
identical with that component of the air which has the same properties.
Simultaneously with Priestley's work, Scheele in Sweden had been
working in much the same way, and had obtained the same gas from
niter, mercuric oxide, and other substances, publication of his results
being delayed, however, until 1777.
Meanwhile Lavoisier, in Paris, who had been studying the rusting
of metals in the air, heard of Priestley's experiments, and demonstrated
that the latter' s " good air " was really a constituent of common air,
and combined with metals when they formed rusts, or " calces," as
they were then called. He proved this conclusively by heating mer-
cury in the retort shown in Fig. 20. The apparatus was arranged in
such a way that a definite vol-
ume of air was inclosed, within
the bell-jar and the retort, by
the larger quantity of mercury
which filled the trough. During
the heating of the mercury in
the retort, the familiar red
powder (mercuric oxide) was
formed on its surface, and sim-
ultaneously the volume of air
diminished, until after twelve
days both of these changes
ceased. The air had suffered
a shrinkage equal to about one-fifth of its volume, and an easily weigh-
able quantity of the oxide had been obtained. The gas which remained
had lost the power of supporting combustion and life, and hence was
named by Lavoisier " azote " (Gk. d priv. and fay, life). In English it
is called nitrogen. The oxide on being heated more strongly by itself
gave off a gas whose volume exactly corresponded with the shrinkage
undergone by the inclosed air, and this gas possessed in an exaggerated
degree the properties which the air had lost. The proof that oxygen
was a constituent of the atmosphere was therefore complete. Lavoisier
named the new element oxygen, or acid-producer (Gk. o£vs an acid, yevvav
to produce), from the fact that the compounds formed by its union with
many elements gave acid (sour-tasting) solutions when they were mixed
with water. Cavendish pointed out almost immediately that there
were sour.-tasting substances which contained no oxygen, so that the
name has no longer any significance.
FIG. 20.
OXYGEN 63
Preparation of Simple Substances. — There are two general
ways of obtaining simple substances. If the element occurs uncom-
bined in nature, as sulphur and gold do, it is only necessary to free it
from foreign materials (impurities) with which .it is mixed. If no such
supply exists, or if the purification is difficult, then some compound,
natural or artificial, is decomposed.
The decomposition, in turn, may be effected in two ways The
compound may be forced apart by the application of energy, usually
in the form of heat or electricity, as in Priestley's experiments (cf. pp.
19-21). Or the desired constituent may be liberated by offering to the
other constituents some substance with which they will unite (see
Preparation of hydrogen). When oxygen is to be liberated, the former
is the more easily applicable plan.
In selecting our source, we are naturally influenced by the cost of
the material, as well as by the ease of the process. Thus, gold oxide
yields oxygen by the application of very little heat, but it is extremely
expensive. Quicklime is very cheap, but does not give up its oxygen
even at the temperature of the electric arc.
Preparation of Oxygen. — 1. Oxygen may be separated from
the other substances mixed with it in the atmosphere by liquefying
the air (see Liquid air), allowing the nitrogen, which is more volatile,
to escape, and finally compressing into tanks the oxygen which evap-
orates last. This is a purely mechanical process.
2. There are many compounds which, when heated to temperatures
under 2000° such as we can obtain with the aid of a Bunsen burner, a
coal fire, or a blast-lamp, give up their oxygen. Some of them are
minerals, but most of them are manufactured articles. Of the min-
erals, pyrolusite (manganese dioxide, Mn02) is an example. It usually
contains the elements of water also, and hence moisture is evolved at
the same time. A substance identical with the mineral hausmannite
(Mn304) remains. Amongst the artificial sources are mercuric oxide,
expensive, but historically interesting (p. 12) ; barium peroxide, used
in manufacturing oxygen on a large scale (Erin's process) ; and potas-
sium chlorate, the most convenient for laboratory use. Many other
substances of this class will be encountered in the sequel.
Erin's Oxygen Process starts from barium oxide (q.v.~). Barium
oxide BaO closely resembles quicklime CaO, but differs from this sub-
stance in the fact that when heated in air to about 500°, it rapidly
acquires additional oxygen and gives barium peroxide. When barium
64 INORGANIC CHEMISTRY
peroxide is raised to a yet higher temperature, this extra oxygen is
given up again. Barium oxide contains one chemical unit weight
each of the two constituents and takes up another of oxygen, so that
the equation for the primary action is :
BaO + 0 -» Ba02.
The subsequent decomposition of the peroxide, during the stage in
which the oxygen is made, is the exact opposite : Ba02 — > BaO + 0.*
The commercial advantage of the method lies in the fact that the
barium oxide remaining after the second stage can be used over and
over again. This, as will be seen, is in reality a chemical method of
obtaining oxygen from the air.
In practice an improvement on the above principle makes working
more economical. It is found that if the barium oxide is maintained
at a temperatiire of 700°, intermediate between the two just mentioned,
oxygen is absorbed when air is forced under pressure into the tubes
containing the oxide. A valve at the extremity of the tubes permits
the escape of the nitrogen. When the combination with oxygen is
completed, the pumping apparatus is reversed, and, a partial vacuum
being created, the oxygen in combination is given off without any
alteration in temperature being necessary. Thus a great waste of fuel
is avoided, and the process is rendered more nearly continuous. This
method furnishes oxygen about 96 per cent pure, and suitable for sale
in compressed form in cylinders.
Potassium Chlorate (y.v.) is a white crystalline substance used in
large quantities in the manufacture of matches and fireworks. When
heated to a moderately high temperature in a tube similar to that in
Fig. 5, it gives off a very large volume of oxygen. Examination shows
that the whole of the oxygen it contains can be driven out. The white
material which remains after the heating is identical with the mineral
sylvite. To the chemist it is known as potassium chloride, and, when
decomposed, it yields one atomic weight each of potassium and chlo-
rine. Its formula is thus KC1. We may infer, therefore, that the
composition of the original substance will be representable by the for-
mula KC10,,,, where x is the number of atomic weights of oxygen.
Measurement and calculation show x = 3. The formula is therefore
* In cases where an action is reversible, and the direction depends on condi-
tions which may be altered, we write both equations in one :
BaO + O <=» BaO,.
OXYGEN
65
KC108, and the equation for the decomposition (see also under Chlo-
rates) :
KC10S -» KC1 -f 30.
To learn the value of x, we ascertain the loss in weight (= oxygen) which a
known quantity of potassium chlorate sustains when heated in a hard glass" tube
closed at one end. By subtraction we get the weight of potassium chloride for-
merly combined with the oxygen. In an actual experiment, 2.998 g. of potassium
chlorate gave 1.169g. of oxygen and left 1.829 g. of the chloride. The atomic
weights of potassium and chlorine are 39.15 and 35.45 respectively, and the
formula-weight of the chloride is therefore 74.6. Dividing the measured weights
of oxygen and potassium chloride by the corresponding atomic and formula weights
(cf. p. 57) respectively: 1.169 -H 16 = 0.07306 and 1.829 +- 74.6 = 0.02452. We
observe that the ratio of the quotients is 2.98 : 1, or almost exactly 3 : 1. The
formula is therefore O x3, (KC1) x 1, or KClOs-
A peculiarity of this action is that admixture of manganese dioxide
increases very markedly the speed with which the decomposition of
the potassium chlorate takes place. Hence, in its presence, and it is
FIG. 21.
generally mixed with the chlorate in laboratory experiments (Fig. 21),
a sufficient stream of the gas is obtained at a relatively low tempera-
ture (below 200°). Without it no oxygen is evolved at all until the
chlorate melts (351°). The dioxide does not begin to lose oxygen
below 400°, and therefore is not itself permanently changed in any
way when used for this purpose.
Familiarity with Physics Required in the Study of Chem-
istry.— In mentioning chemical phenomena, it is inevitable that considerations
of space should limit our statements to the merest indication of the process and
the briefest record of the chemical result. The prodigious disproportion between
66 INORGANIC CHEMISTRY
the meagerness of this fragment and the mass of detail which lies behind it in each
case should be constantly before the mind of the reader. The book gives empiri-
cal knowledge, the laboratory work and the discussion of it furnish the only real
knowledge. The extent and nature of this real knowledge may be shown in con-
nection with any action. As an illustration we may point out some of the prob-
lems which the heating of potassium chlorate presents to one who is trying to
acquire an intelligent acquaintanceship with its chemistry, and has not previously
done the experiment.
First, the substance melts. It must be realized that this is a common occur-
rence which does not necessarily imply any profound change, and may be reversed
by cooling. Later, the liquid appears to boil, and the properties of a boiling sub-
stance must be known. If the observer has been informed in advance that the
body is -homogeneous, he must know that, if it is simply boiling, it will evaporate
completely and leave nothing behind, and that the temperature required to achieve
this will remain constant from the beginning to the end. In order, therefore, to
become aware of the fact that here decomposition is taking place, he must note the •
ways in which the decomposition of potassium chlorate differs from ordinary boil-
ing. For example, if it were a case of boiling, he should expect to find the solid
body condensing on the sides of the tubes, and note the fact that no such conden-
sation is observed, with the appropriate inferences. He should observe that, in
the later stages at least, the agitation of the liquid does not cease when the flame
is removed, although this would undoubtedly occur in a case of simple boiling.
He must further observe the changes in the consistency of the material and the
way in which it finally becomes thick and may even solidify. Even the most ex-
perienced investigator would have to make many careful experiments before he
could definitely classify the nature of the phenomenon being observed. The first
inference would probably be that the phenomenon was certainly not one of mere
ebullition. In some ways it is like the evaporation of a solution, obtained, say,
by the melting of a substance in its own water of crystallization. Yet, this hypoth-
esis would not explain to the thoughtful observer even the more obvious features
of the phenomenon, for the liquid which was acting as a solvent would have to be
amazingly volatile if the absence of any condensation on the walls of the tube was
to be accounted for.
The illustration need not be elaborated further. These remarks are sufficient to
show that even the simplest experiment presents an almost limitless field for the
discussion of important questions which are more or less common to all chemical
phenomena. It must be noted, also, that chemical change is in itself not percept-
ible by the senses, and that only physical properties and physical phenomena are
observed (cf. pp. 36-40). The chemical facts, such as the general nature of the
change, the conditions under which, and the facility with which it occurs, are
reached solely by inference. The above example shows the ready and thorough
knowledge of physics which must be at the command of every individual effort to
study even the simplest chemical phenomenon. It is only when the physics as
well as the chemistry of the change have been mastered that the " real knowledge "
to which reference was made above has been gained.
Physical Properties of Oxygen. — Oxygen, as a gas, resembles
air in being colorless, tasteless, and odorless. It is slightly heavier
OXYGEN 67
than air ; its density, using the physical standard of air = 1, is 1.105.
The chemist often uses hydrogen as his standard, and oxygen is
15.900 times (Morley) as heavy. One liter of oxygen, at 0° and 760
mm. barometric pressure, weighs 1.42900 grams (Morley). The gas
dissolves to some extent in water, the solubility at 0° being four vol-
umes of gas in one hundred volumes of water (at 20°, 3 : 100). The
critical temperature (q.v.~) of oxygen is — 118° C. At that tempera-
ture, fifty atmospheres pressure is required to liquefy it. Liquid
oxygen, which was first made by Wroblewski, has a pale-blue color,
and boils under one atmosphere at — 182.5°. Its specific gravity at
— 182.5° is 1.13 (water = 1) : that is to say, 1 c.c. weighs 1.13 g. By
cooling with a jet of liquid hydrogen, Dewar froze the liquid to a
snow-like, pale-bluish solid. The magnetic properties of the element
are clearly shown by the force with which a tube of liquid oxygen is
attracted by a magnet.
Specific Cfiemical Properties. — Under this head we describe
the chemical behavior of a substance, enumerating the other sub-
stances, simple or compound, with which it unites or interacts (p. 14),
stating the conditions peculiar to each action, and estimating the
intensity of the tendency to chemical change in each case. In the
case of a simple substance like oxygen we note particularly with how
many of the other elements it can form compounds, how far it unites
with them directly, and in how many cases the compounds have to be
made by indirect means. In general, we call those simple substances
active which unite with many other simple substances and do so by
direct union. Oxygen, for example, is active, and nitrogen (q.v.^) is
relatively inert. The intensity of a chemical action is judged by its
speed, by the electricity it can generate, or, roughly, by the heat evolved
during its progress (p. 28).
Chemical Properties of Oxygen. — Sulphur, when raised in
advance to the temperature necessary to start the action, unites vigor-
ously with oxygen (Fig. 22), giving out much heat and producing a
familiar gas having a pungent odor (sulphur dioxide). This odor is
frequently spoken of as the " smell of sulphur," but in reality sulphur
itself has no odor, and neither has oxygen. The odor is peculiar to
the compound of the two. It may be added that the mode of experi-
mentation can be changed and the oxygen led into sulphur vapor
through a tube. The former then appears to burn with a bright
flame, giving the same product as before,
68
INORGANIC CHEMISTRY
Warm phosphorus combines with oxygen with even greater vigor,
and forms a white, powdery, solid compound (phosphoric anhydride),
which absorbs moisture from the aqueous vapor in the air and quickly
forms a solution in this water. In both these cases the products differ
from oxygen, not only in odor and
in being a gas and a solid, respec-
tively, but notably in that, when
shaken with water, they dissolve
and interact to * form acids (see
below).
Burning carbon, in the form of
charcoal, when plunged into the gas,
glows much more brightly than in
ordinary air. In this case the pro-
duct is a gas (carbon dioxide).
When this gas is shaken with
" lime-water," a solution of calcium
hydroxide Ca(OH)2 (q.v.~), a white
precipitate of calcium carbonate
CaC03 is formed.
Finally, metallic iron, which is
simply rusted by air (diluted oxy-
gen), burns in pure oxygen with
surprising brilliancy. Globules of
a molten product fall from the iron,
and when they have cooled are
found to consist of a dark-gray
brittle material, which we recognize as identical with blacksmith's
hammer-scale, and with a well-known ore of iron (magnetic oxide of
iron).
Experiments like the above with the whole series of simple sub-
stances would show that oxygen unites in a similar way with nearly
every member of the list, and often, though not always, with the same
vigor as in the case of these examples. In the case of one or two
elements, such as gold and platinum, the compounds are obtainable by
double decomposition, and not directly. With the five members of the
helium group, of which no chemical compounds are known, and with
fluorine, oxygen does not combine. These three sentences summarize
the chemical properties of oxygen.
Oxygen can unite with many of the same elements when they are
FIG. 22.
OXYGEN 69
already in combination. Wood, for example, is composed of carbon
and hydrogen, with a small amount of oxygen. When previously
heated, it is decomposed, and the constituents unite with oxygen form-
ing carbon dioxide and water.
The Making of Equations Again. — To learn the exact nature
of interactions like those used as illustrations above, quantitative ex-
periments must of course be made and the usual method employed to
obtain the formulas of the products. Thus, for example, a known
weight of sul-
phur is placed in -m-J-1 ^
a porcelain boat =U-E -> 2 ^-l
(Fig. 23), which
has already been
weighed. The
U-shaped tube to
the right con-
tains a solution FlG> 23>
of potassium hy-
droxide which is capable of absorbing the resulting gas. The oxygen
enters from the left. When the sulphur is heated, it burns in the
oxygen, and the loss in weight which the boat undergoes shows the
amount of sulphur consumed. The gain in weight of the U'tube
shows the weight of the compound produced. By subtracting from
this weight that of the sulphur, we get the quantity of oxygen. The
proportion of the constituents and the steps in the calculation are as
follows :
PERCENTAGE* AT. WT. FACTOR -=-1.561
Sulphur, 50.05 = 32.06 x 1.561 = S X 1.561 S x 1
Oxygen, 49.95 = 16.00 x 3.122 = Ox 3.122 O X 2
The formula of the product is therefore. S02 and the equation S + 20
Similarly, phosphoric anhydride may be shown to have the for-
mula P205, carbon dioxide C02, and magnetic oxide of iron Fe304.
The results given by the above experiment are usually inexact. The tendency
to the formation of sulphur trioxide, often heightened by catalytic action (see
* Here the percentages are employed. The actual weights found in one experi-
ment, however, may be divided by the atomic weights and the same result obtained
(p. 65) . In fact, any two numbers bearing the proper ratio to one another may be
used.
70 INOBGANIC CHEMISTRY
below) of the porcelain of the boat, raises abnormally the proportion of oxygen.
The principle of the experiment is easy to imderstand, however.
In the case of phosphorus a similar plan of experiment may be used. Instead of
attempting to receive the solid product in a U-tube, however, it must be caught by
a plug of glass wool in the main tube of hard glass, and a drying tube will be
needed at the end to prevent admission of moisture from the air. The increase in
weight of the hard glass tube represents the oxygen taken up. Care and leisurely
performance are needed to make the experiment successful.
The measurement of the composition of carbon dioxide * gives the most exact
results when carefully conducted. The data and working in these cases and in
that of iron are as follows :
PERCENTAGE AT. WT. FACTOR -=- .704
Phosphorus, 43.66 31.0 x 1.408 Px2
Oxygen, 56.34 16.0 x 3.521 O x 5
-H 2.272
Carbon, 27.27 12.0 x 2.272 C x 1
Oxygen, 72.72 16.0 x 4.545 O x 2
-^-0.431
Iron, 72.38 = 55.9 x 1.295 Fe x 3
Oxygen, 27.62 = 16.0 x 1.726 O x 4
The density of the vapor of phosphoric anhydride leads to the formula P40,0,
but, as the substance is never used as a gas, the simpler formula is generally
preferred.
Oxides. — Substances containing one element in combination with
oxygen are called oxides, and processes like those described above are
called oxidizing processes, or oxidations. When the same element
forms more than one oxide, the names of the oxides indicate the differ-
ing proportions. Thus we have barium oxide (or monoxide) BaO, and
barium peroxide (or dioxide) BaO,, magnetic oxide of iron Fe304, fer-
rous oxide FeO, and ferric oxide Fe203. In cases like the two last
the terminations -ous and -ic applied to the metal correspond to the
smaller and larger proportions of oxygen, respectively, which the metal
is able to hold in combination. Oxides of the form Fe203 are often
called sesquioxides because they contain one and a half (Lat. sesqui-,
one-half more) unit weights of oxygen to each one of iron.
Many oxides, like those of iron, are quite indifferent to water, but
others, like those of sulphur .and phosphorus, interact with it (see
under Water). Some give sour solutions, containing acids dissolved in
the excess of water. Such solutions turn blue litmus, a vegetable dye,
red. Others give solutions with a taste like soap or borax, and here
* Described in the author's Laboratory Outline of General Chemistry.
OXYGEN 71
the dissolved substance is called a base (g.v.'), and turns litmus blue.
Thus sulphur dioxide and phosphoric anhydride give sulphurous acid
and phosphoric acid respectively :
S02 + H20 -+ H2S08
P205 + 3H20 -» 2H3P04.*
If the product, to whichever class it belongs, is not volatile, it may be
obtained by evaporating the excess of water. In the case "of sulphur-
ous acid the above action is reversed by evaporation and the sulphur
dioxide and water both pass off ; in that of phosphoric acid, the white
crystalline acid is obtained. In consequence of their relation to the
acid, differing from it in not containing the elements of water, these
oxides are often called anhydrides.
The discussion of the formation and properties of ozone (q.v.'),
which is an oxide of oxygen, cannot be taken up until we are in pos-
session of the means of understanding the difference between the two
substances.
Combustion. — Sfrnce oxygen is a component of the atmosphere,
chemical actions in which it plays a part are familiar in daily life.
Violent union with oxygen is called in popular language combustion
or burning. Yet since oxygen is only one of many gaseous substances
known to the chemist, and similar vigorous interactions with these
gases are common, the term has no scientific significance. The union
of iron and sulphur, even, gives out light and heat, and is quite simi-
lar in the chemical point of view to combustion.
In connection with this, however, it may be worth while to notice
the distinction between combustible and incombustible substances.
Things which are incombustible, using the term in its popular sense,
may be divided into two classes. There are those substances which
already contain all the oxygen which they can hold in combination.
Such are the oxides whose formation we observed in the experiments
described above. In everyday life, limestone, sand, bricks, and most
rocks are illustrations. The other substances ordinarily classed as
incombustible are those which do not unite with oxygen, as it is found
in the air, with sufficient vigor. The iron used in the construction of
fireproof buildings is the commonest example of this class.
* Here summation of the formulae on the left would give H6P2O8, but in such
cases, unless there are reasons to the contrary, the common factor is put in front
and the formula reduced to its lowest terms.
72 INOKGANIC CHEMISTRY
Oxidation. — The rusting of metals differs from combustion only
in speed. Thus, magnesium ribbon, when left lying in the air, gradu-
ally becomes covered with a white crust. By scraping this off, and
continually exposing the surface of the metal, the whole of the latter
may eventually be transformed into the white powder. The product
is identical in composition and properties with the oxide formed by
combustion (p. 44). In the case of iron, burning gives us the
magnetic oxide (Fe304), while rusting in moist air yields a hydrated
ferric oxide (Fe2O3 -f Aq*). The products differ in composition, but
are closely related.
This process of slow oxidation, although less conspicuous than
combustion, is really of greater interest. Thus the decay of wood is
simply a process of oxidation whereby the same products are formed
as by the more rapid ordinary combustion. Large volumes of pure
water are mixed with sewage, the object being not simply to dilute the
latter but to introduce water containing oxygen in solution. This has
an oxidizing power like that of oxygen gas, and, through the agency of
bacteria, quickly renders dissolved organic matters innocuous by con-
verting them for the most part into carbon dioxide and water. In our
own bodies we have likewise a familiar illustration of slow oxidation.
Avoiding details, it is sufficient to say that the oxygen from the air
taken into the lungs is carried by the blood throughout our tissues and
is there used for oxidizing waste materials. The carbon dioxide is
carried back to the lungs by the blood, and finally reaches the air dur-
ing exhalation. To supply the place of the material thus removed, we
are' under the continual necessity of building new tissue from the food
which we eat. If we cease to eat, we become lighter and weaker, show-
ing that a real portion of our structure is gradually being consumed by
oxidation.
The opposite of oxidation, the removal of oxygen, is spoken of in
chemistry as reduction.
Means of Altering the Sjteed of a Given Chemical Action : By
Change of Temperature. — That the same change may proceed with
very different speeds according to conditions is a familiar fact. For
example, raising the temperature increases the rapidity of all chemical
interactions. Thus, cold iron combines with oxygen very slowly, giv-
* The formula H0O may not be used excepting to indicate a definite proportion
of the elements of water (18 parts). Where the proportion varies according to
circumstances, as here and in the case of solutions, the contraction Aq is employed.
OXYGEN 73
ing rust, while white-hot iron sheds quantities of scales of an oxide,
formed in the few moments that it is under the blacksmith's hammer.
White-hot coal unites with oxygen in the air to form carbon dioxide
and seems to disappear before our eyes, while in the cellar, even in
warm weather, we observe no appreciable diminution in its amount.
The chemist, however, has reason for considering that even here the
difference is one of degree only. No temperature can be found at which
the interaction definitely begins. We believe that every change, pro-
vided, like these, it involves a liberation of energy (cf. p. 27), proceeds
with some speed at every temperature. A rough estimation, based on
experiment, shows that on an average, other things being equal, every
rise in temperature of ten degrees doubles the amount of material
changed per second, and conversely.
If, on bringing two materials together, the chemist observes no
marks of chemical action, he immediately begins cautiously to heat the
mixture. This appeal to the accelerating effect of a rise in tempera-
ture is always made as a matter of course.
The common expressions used in chemistry in describing temper-
atures, along with the corresponding readings of the thermometer, are
as follows :
Incipient red heat, about 525°. Yellow heat, about 1100°.
Dark red heat, " 700°. Beginning white heat, " 1300°.
Bright red heat, « 950°. White heat, " 1500°.
Rapid Self-sustaining Chemical Action and Means of Initiat-
ing it. — When a piece of wood is set on fire at one end, the heat pro-
duced by the action itself raises the temperature of neighboring portions
until their speed of union becomes equal to that of the part originally
lighted. In this way the whole becomes finally inflamed. When we
blow the blaze out, the great excess of cold air suddenly lowers the
temperature of the wood, and of the gas rising from it, and rapid union
ceases. Whether a given set of materials can maintain itself at a tem-
perature proper to violent interaction will depend on the amount of heat
developed by the action itself on the one hand, and the losses of heat
by conduction and radiation on the other. If the latter are great, the
former must be greater. Thus the union of iron and oxygen per se
gives heat enough to warm the materials to the burning temperature and
leaves much over for radiation. But iron in air, which is four-fifths
nitrogen, can receive the oxygen only one-fifth as fast at the start, and
even more slowly as, later, the nitrogen accumulates round it. And be-
74 INORGANIC CHEMISTRY
sides, all the nitrogen has to be heated to, perhaps, 2000°. The task is
too great. The union is impeded and the iron is not oxidized fast
enough to generate the heat required to maintain everything at this
high temperature. Poor conductors of heat, like wood and candles, fare
better. Powdered iron, with its particles presenting large surface to
the air relatively to the weight of material in each particle to be heated,
burns well.
The initial supply of heat required to start violent exothermal chem-
ical actions, of which alone we are here speaking, must not be confused
with the heat subsequently developed as the action proceeds. The lat-
ter is usually much greater. Indeed, the preliminary supply varies
with circumstances, and may be made as small as we choose by limiting
the area first heated and using ordinary precautions against radiation
and convection. Indeed, in practice, a single spark from an induction
coil often takes the place of more clumsy methods of raising the tem-
perature. The heat produced by the interaction itself, however, is fixed
in amount, and depends only on the materials and their quantity. En- '
dothermal actions differ completely from those under discussion. In
them a definite, large quantity of heat has to be furnished, and the ac-
tion instantly ceases if the supply fails.
Heating is not the only means used to give the initial acceleration
to a self-sustaining chemical change. The materials in a match-head
are capable of undergoing a great transformation. Yet, so slowly does
this proceed at ordinary temperatures, that matches may be kept in
efficient condition for years. Here a rather violent vibration is em-
ployed to hasten the torpid action in a small part of the material, and
the heat produced by the resulting action quickly ignites the whole.
The same explanation accounts for the explosion of gun-cotton by a
percussion fuse.
Other Means of Altering the Speed of a Given Chemical
Change: By Change in Concentration; by Catalysis; by Solu-
tion. — Even when the temperature remains constant, there are other
changes in the conditions (p. 52) which may be used for accelerating or
for moderating the speed of chemical interactions. The most impor-
tant of these is, a change in the concentration of the interacting sub-
stances. Another is the presence of a catalytic agent. The condition
of solution might be accounted still another.
The abatement in the activity of the oxygen found in the air by the
nitrogen which is mixed with it, is a question of concentration If the
OXYGEN 75
concentration of pure oxygen under atmospheric pressure is taken as
unity, that of oxygen in air is only about 0.2. And the speed of inter-
action of a body, other things being equal, is directly proportional to
its concentration. This is not an obscure law, but merely common
sense put into definite language. The opportunity which one substance
has for getting at every part of another will be one factor in determin-
ing the speed with which the resulting transformation will take place.
And this opportunity, other things being equal, depends on the thick-
ness or density with which the substance is scattered in the region of
action. In the case of a gas, this factor is measured by its partial press-
ure. Hence, lights burn badly at great elevations, where the oxygen
is very tenuous. On the other hand, powdered charcoal interacts so
rapidly when ignited in liquid air, where the oxygen is highly con-
densed, that an explosion takes place, whereas in common air it burns
feebly. Again, when oxygen is compressed in contact with barium
oxide at 700° it combines to form the dioxide (p. 64) ; when the press-
ure of the oxygen in contact with the latter is reduced, oxygen is lib-
erated (see Chemical equilibrium).
When an extra substance increases the speed of a chemical change,
seemingly by its mere presence, without itself suffering any permanent
change, we call this catalytic (Gk. Kara, down ; Averts, the act of loosing)
or contact action . The word was originally used for cases of decomposi-
tion. The foreign body is called the catalytic or contact agent, and the
process catalysis. The effect of manganese dioxide on the decompo-
sition of potassium chlorate (p. 65) is of this nature. When some of
the chlorate is melted carefully, so as to avoid superheating, scarcely
any evolution of oxygen can be perceived at this temperature (351°).
If now a thin glass bulb filled with powdered manganese dioxide be
broken in the molten mass, the oxygen is given off in torrents in con-
sequence of the enormous acceleration of the decomposition. Yet the
manganese dioxide itself may be recovered unchanged from the residue.
It is found that many actions owe what appears to be their normal
speed to the presence of a trace of water vapor. Thus many of the
elements show no visible tendency to unite with carefully dried oxy-
gen, even when they are strongly heated in it. Addition of a trace of
moisture, however, brings about instant combustion. So water is to be
regarded as one of the commonest contact agents (see Chemical prop-
erties of chlorine).
A few cases of retardation of an action by a catalytic agent are known. Thus,
a little benzyl alcohol or mannite added to the solution will retard the oxidation of
76 INORGANIC CHEMISTRY
sulphites by the air. Hence positive and negative catalysis both occur (see ab-
stract of an address by Ostwald in Nature, LXV, 522, given also in full in the
Zeitschrift fur Elektrochemie VII, 995).
The effect of solution in hastening a chemical change was seen
when we examined the interaction of sodium chloride and silver nitrate
(p. 13). With the solutions the action was seemingly instantaneous.
If we had attempted to bring it about by rubbing the dry substances in
a mortar, hours of work would have left much of the original bodies
still unchanged. Even heating would not have produced so prompt an
effect. It is obvious that the intimate access which every part of each
solution gains to every part of the other accounts to some extent for
the difference (see lonization). Chemical actions, as will be seen in the
sequel, are very frequently carried out in aqueous solution in order to
take advantage of the favorable influence of this condition.
Thermochemistry. — As we have seen (p. 27), the free or available
chemical energy of a system undergoing chemical change usually ap-
pears in the form of heat. Since it is often instructive to consider the
amount of heat produced, some of the elementary facts of thermo-
chemistry must be explained.
The chemical interactions to be studied thermally are arranged so
that they may be carried out in some small vessel which can be placed
inside another containing water. The heat developed raises the tem-
perature of this water. Where gases like oxygen are concerned, a
closed bulb of platinum forms the inner vessel. The quantity of heat
capable of raising one gram of water one degree in temperature, be-
tween 0° and 100° Centigrade, is called a calorie. So that 250 grams of
water raised 1° would represent 250 calories, and 20 grams of water
raised 5° would represent 100 calories.
While in physics the unit of quantity is the gram, in chemistry the
unit which we select is naturally that represented by the formula of
the substance. Thus, the heat of combustion of carbon means the heat
produced by combining twelve grams of carbon with thirty-two grams
of oxygen, and is sufficient to raise 100,000 grams of water one degree.
This is expressed as follows :
C + 20 -> C02 + 100,000 cal.
In other words, the combustion of less than half an ounce of carbon
will raise one kilogram (over two pounds) of water from 0° to the
boiling-point
OXYGEN 77
It is always found that the same quantities of any given chemical
substances sustaining the same chemical change under the same condi-
tions produce or absorb, according as the action is exothermal or en-
dothermal (p. 27)-, amounts of heat which are equal.
The rate at which a given chemical action is allowed to take place
has no influence on the total amount of heat consumed or produced.
It may not at first sight appear obvious that rusting evolves heat, but
a delicate thermometer would show that a heap of rusting nails was
somewhat higher in temperature than surrounding bodies. Poor
conductors, like oily rags and ill-dried hay, show a tendency to spon-
taneous combustion owing to accumulation of the slowly developing
heat of oxidation. The warmth of our own bodies is in part due to the
same cause.
In accordance with invariable experience expressed in the law of
the conservation of energy, when an action is chemically capable of
reversal, the contribution of the same amount of heat which it de-
velops will exactly suffice to drive the chemical change in the opposite
direction. The heat contributed is simply used to restore the amount
of chemical energy proper to the original system. Thus, the union of
one chemical unit weight each of mercury and oxygen (p. 62) produces
30,600 cal. :
Hg + 0 -+ HgO + 30,600 cal.,
and the decomposition of one formula-weight of mercuric oxide (p. 12)
demands the same amount of heat in order that free mercury and oxy-
gen, with their appropriate proportions of chemical energy, may be
recovered.
In practice it is found that all chemical changes are not capable of
reversal by the use of the sources of heat available in the laboratory.
A quantity of heat, equivalent to that produced by any chemical action
on a small scale, is very easily provided, but something more appears
to be necessary. The heat provided must be of a certain temperature,
otherwise it is quite ineffective. For example, the heat produced by
the union of calcium and oxygen is within the limits of ready meas-
urement,
Ca + O -> CaO + 131,000 cal.,
and the supply of this amount (or even of unlimited amounts) of heat
to calcium oxide (qiiicklime) is easily achieved. Yet this method is
quite ineffective to produce decomposition (p. 63) of the product.
78 INORGANIC CHEMISTRY
Apparently we have not sufficiently high temperatures for the purpose
at our command (see Factors of energy).
It may be noted in this connection that the temperatures required
to produce reasonably rapid decomposition vary within a wide range.
Some substances can be kept only below 0°, and decompose when
allowed to become warm. Others, like the oxides of gold and platinum,
require a little heating (p. 63). Many, like quicklime, are not broken
up even at the temperature of the electric arc. When the energy is
applied in the form of electricity (p. 19), instead of heat, the range is
incomparably more easily within the reach of the means ordinarily at
our disposal. There is no substance, provided it is of such a nature as
to be affected by the electric current at all, which cannot be decomposed
by a current with an E.M.F. of 10 volts or less, while currents of 110
volts and over are commonly accessible. It is partly on this accoimt
that electrical processes have become so common in industrial chem-
istry.
One of the most important principles of thermochemistry is the
law of constant heat summation. If a system of substances can be
transformed into another system of substances by different stages or
by more than one route, then the algebraic sum of the heats absorbed
or produced in the various stages is the same. Thus, barium oxide
might be formed directly from the proper proportions of the constitu-
ents, or it might be made by preparing the dioxide (p. 63), and then
driving out half of the oxygen contained in the latter. The quantities
of heat involved in these changes are as follows :
Ba + 0 -» BaO + 124,400 cal. (1)
( Ba + 20 -» Ba02 + 141,600 cal. (2)
( Ba02 -» BaO + 0 - 17,200 cal.* (3)
When the last two equations are added algebraically, canceling terms
such as Ba02 and O, which are common to both sides of the final
equation, the chemical action is seen to be the same as in (1), and the
balance, in favor of heat produced, is 124,400 calories as before. If
in such cases the sum of the heats was not the same, it would follow
that by using different plans of procedure we could prepare different
specimens of the same substances containing different proportions of
chemical energy. This, however, we have never been able to do (p. 77).
* Since (3) is to be added to (2) so as to give (1), the formulae of (3) must be
so placed that if the initial substances of (1) are on the left side of (2), the products
of (1) may be on the right side of (3) and vice versa.
OXYGEN 79
The quantities of heat liberated in two chemical changes are often
measures of the relative amounts of available chemical energy in the
systems before the change, and, therefore, often furnish a measure of
the relative chemical activities of the two sets of substances. The
comparison may safely be made in certain cases when the conditions
under which the two actions take place are precisely alike. Formerly
it was supposed that the heat liberated was always proportional to the
chemical activity of the substances, but we have already shown cause
(pp. 27-28) why this general statement cannot be true.
Exercises. — 1. Construct the equations for the combustion of
phosphorus, carbon, and iron in oxygen (p. 70).
2. When 1 g. of sodium burns in oxygen, it produces 1.7 g. of the
oxide. What is the formula of the latter, and the equation (p. 69) ?
3. Which are the components (p. 34) of the liquid made by treating
phosphoric anhydride with water ? Which are the constituents (p. 34)
of phosphoric acid?
4. How should you show that, in the making of oxygen from a
mixture of potassium chlorate and manganese dioxide, the latter
remains unchanged ? Which properties (p. 37) are you employing for
this purpose ?
5. Discuss the union of iron and sulphur (p. 11) and the decompo-
sition of mercuric oxide (p. 12) in their relation to the explanations on
pp. 73-74.
6. How many calories are required to raise 500 g. of a substance
of specific heat 0.5 from 15° to 37° (p. 76)?
7. The combustion of 1 g. of sulphur to sulphur dioxide develops
2220 calories. What is the heat of combustion of sulphur (p. 76) ?
8. Outline briefly the proof that thermochemical data are not accu-
rate measures of chemical activity (p. 79).
CHAPTEE VI
THE MEASUREMENT OF QUANTITY IN GASES
WE have spoken of measuring the proportion by weight of the
oxygen used in several chemical changes, but in our illustrations we
have never weighed the gas itself. We have always (e.g. p. 69)
obtained its quantity by subtracting the weights of solid or liquid
bodies. In practice this method often serves the purpose.
Our preference for weighing as a means of ascertaining quantity
of matter is largely due to the fact that the weight is independent of
the physical or even chemical condition of the substance. Yet, with
proper precautions, we may learn the quantity of matter by means of
any other attributes which are proportional to it. Now the volume is
such an attribute. In determining the quantity of a liquid, where
rapidity with no great accuracy is desired, the volume is frequently
measured. In the case of gases the error made in measuring the vol-
ume is less, as a rule, than in measuring the weight.
The Variable Concentration of Gases. — A little experience
with gases soon shows us that measurement of volume alone does
not necessarily give any definite idea of the quantity of matter
present. The denseness with which the gaseous matter is packed (the
concentration of the gas) in the vessel must somehow be defined, as
well as its volume, in order that there may be specification of the
quantity of matter.
Gases vary markedly in chemical activity with changes in their concentration
(cf. p. 75), and thus the consideration of this condition (p. 62) is continually
forced upon the chemist. Solids and liquids do not alter their denseness of pack-
ing (concentration) very noticeably even when compressed severely or changed in
temperature. So that concentration need not be considered in the case of pure
bodies in the solid or liquid forms. Such substances can be scattered through a
variable space by solution in some suitable solvent, however, and then their degree
of packing or concentration becomes an important factor in their chemical behavior
also.
The principle used for estimating the concentration of a gas may
be illustrated by means of the arrangement in Fig. 24. Except that
80
THE MEASUREMENT OF QUANTITY IN GASES
81
a little gas (any gas will do) remains shut off by the mercury in the
left limb of the tube, the whole apparatus has been evacuated. The
reservoir can be turned upward, and thus larger amounts of mercury
may be introduced into the tube.
Now the portion of the mercury below the dotted line is essentially
a balance, that is to say, it will move in one direction or the other if
the stresses on either side change. At
present these stresses must be equal. On
the right pan of the balance, so to speak,
the stress is represented by the weight of
the column of mercury above the dotted
line. As there is nothing in the tube
above this mercury, the weight of the
latter is all that this side of the balance
sustains. On the left pan of the balance
there must be an equal stress, and this
stress can be caused only by the gas con-
fined in the shorter limb. The nature of
a gas suggests that this stress must be
exercised on the walls of the tube also,
although they naturally do not exhibit
its effects. This stress we call the press-
ure or tension of the gas.
The height of the surface of the mer-
cury on the right above that on the left
having been measured, more mercury may
now be added from the reservoir, and the
difference in the two levels again noted.
The gas cannot have diminished in amount,
yet it now occupies a smaller space, and is, therefore, packed more
closely — its concentration is greater than before. If, for example, the
difference in level is now twice as great, it will be found that the con-
centration of the gas is also twice as great (its volume having become
half of the original volume). Whatever amount of mercury is added,
we shall always find that the concentration of the gas is proportional
to the height of the mercury. But this in turn is proportional to the
weight of metal. The weight of mercury on one side must, therefore,
be equal to the stress or pressure or tension of the gas on the other
side which balances it. Hence, the concentrations of a sample of any
gas are proportional to the corresponding pressures it exercises. We
FIG. 24.
82
INORGANIC CHEMISTRY
determine, therefore, the denseness with which any sample of gas is
packed by measuring its pressure.
Method of Allowing for Varying Concentration in Measur-
ing Quantity in Gases. — The principle just stated is applied to the
measurement of the quantity of matter in a sample of gas by per-
mitting the concentration of the sample to alter until it is equal to
that of the atmosphere at the moment. Then we
read off the volume now occupied, and simulta-
neously we ascertain the pressure by observing
that of the atmosphere. Each of these two oper-
ations is facilitated by a special arrangement of
apparatus.
A gas to be measured is always confined so
that some liquid constitutes one part of the barrier
to its escape. The very simplest form of the
apparatus is shown in Fig. 25. To render the
concentration (and pressure) of the gas equal to
that of the atmosphere, the cylinder 'containing
the gas is lowered (or raised) until the levels of the
liquid inside and outside are the same. When the
system is in this condition the stress of the gas on
the inner surface must be equal to that of the
atmosphere on the outer one, otherwise movement
of the liquid would occur. The volume of the gas
is then read directly from the graduation on the
cylinder. Often the cylinder, or other vessel, is
FIG. 25. closed with a ground-glass plate, placed quickly in
erect position, and weighed. The weight of water
which is then required to fill it to the brim gives more exactly the
volume occupied by the gas (1 g. water = 1 c.c.). When special
modes of admitting or handling the gas have to be provided for,
the apparatus may be more complex. But the principle of adjust-
ment is always the same. In exact work, mercury is employed to
confine the gas. Water serves the purpose of rough work with
gases which, like oxygen, are but slightly soluble in it. When
water is used, the volume is too great by the space occupied by
the vapor of the water which is mixed with the gas (see Mixed gases),
and a correction must always be made on this account.
The pressure or tension of the atmosphere at any moment is meas-
THE MEASUREMENT OF QUANTITY IN GASES
83
ured by means of a simplified form of the apparatus in Fig. 24. The
reservoir is omitted. The atmospheric air being the gas whose concen-
tration is to be measured through its pressure, the short limb is left
open. The resulting apparatus (Fig. 26) performs its functions in the
same way as does the more complex one. The only difference is that
mercury is automatically added to or withdrawn from the right side
by the motion of the metal resulting from
changes in the pressure of the air. The
reading of the vertical height between the
lower and upper surfaces of the mercury
gives a number which is proportional to the
weight of mercury on the right side of the
balance and, therefore, to the (equal) stress
of the atmosphere on the left. This is called
the barometric reading (unconnected), after
the name of the instrument.
To make different readings, taken when
the mercury is at different temperatures,
strictly proportional to the weight of the
metal, the observed height is always reduced
to that which would have been shown by
the same weight of mercury at 0° in the
same apparatus. A thermometer, and a table
of temperatures with the corresponding cor-
rections to be subtracted from the uncorrected
reading ((7, Fig. 26), must be used.
Knowing now the volume occupied by
the sample of gas when its concentration is
equal to that of the atmosphere and the,
barometric reading, which is proportional to
this concentration, the measurement of the
amount of matter in the sample has become
definite so far as concerns the variability of concentration with change
in pressure.
The recorded results of measurements made as above at different
times are still unsatisfactory because the data for samples of the same
kind of gas differ in the value of the pressure as well as in that of the
volume. To make the results easily comparable in respect to the
amount of matter they represent, one further step is needed. All
the data are recalculated so as to show the volume each sample would
FIG. 26.
84 INORGANIC CHEMISTRY
have occupied if the pressure had been equal to the weight of 760 mm.
of mercury, which is the average height of the barometer at the sea
level in 45° of latitude.
We have seen that the concentration of a given quantity of a gas
is proportional to its pressure (p. 81). But, volume occupied is the
inverse of concentration. Thus the same rule may be stated in the
form, the volumes occupied by a sample of any gas are inversely pro-
portional to the pressure at each volume. The fact was discovered by
Boyle (1660) who stated it in this way. In still other words, the
product of the several pressures and corresponding volumes of a
sample of gas is a constant.
A numerical illustration will show the mode of applying this rule.
We measure 200 c.c. of a gas at atmospheric pressure, and the barom-
eter reads 742 mm. The question is : What would be the volume
of this amount of gas at 760 mm. barometric pressure ? It will be
200 x ^f § c.c. = volume at 760 mm. It is unnecessary to use any
formula, but absolutely essential to ask : Is the new pressure greater
or less than the old ? Here it is greater. Hence, according to the
law, the new volume will be less, so that the fraction must be arranged
with the smaller number in the numerator.
Boyle's law may be illustrated by using a long tube bent like a barometer (Fig.
26) but having the short limb closed and the long one open. Strips of paper mark
the levels of the mercury, which are at first alike on both sides, and register the
volume of the air in the short limb at a pressure of one atmosphere. The reading
of the barometer at the time, say 750 mm. , is noted. Then mercury is added through
a funnel inserted in the long limb, until the level in this limb is 750 mm. above
that in the other. A stick cut to 750 mm. length, and held beside the tube, will
conveniently show when this has been accomplished. The pressure in the open
limb being now two atmospheres, the volume of the air will be found to have
diminished to one-half its former value.
For pressures lower thai one atmosphere, a different arrangement must be
used. A graduated tube, closed at one end, is partly filled with mercury and in-
verted in a tall, narrow cylinder full of the same metal. The tube is then clamped
in any position, such that the mercury level in the tube is above that in the cylinder.
The reading of the barometer is next noted. The volume occupied by the air in
the tube is then read, and the difference in height of the two mercury surfaces is
measured by means of a graduated rule. Subtracting this height from that of the
barometer gives the pressure of the air in the tube. The position of the tube is
then altered, and the same measurements repeated, as often as is wished. The
product of each volume and the corresponding pressure will be a constant number.
The law is expressed mathematically by letting pl and w, represent the initial press-
ure and volume, p2 and v2 the new pressure and volume, and so forth. Then
P\°\ = P?vz = constant for that particular specimen of gas. For a given sample
of gas, any one of the four values may be calculated if the other three are known.
THE MEASUREMENT OF QUANTITY IN GASES
85
Boyle's law states the facts with sufficient accuracy for all ordinary
purposes. But in reality no two gases behave precisely alike in re-
spect to change in concentration when the pressure is altered (see Molar
weights). The same gas does not even behave in precisely the same
way at high, intermediate, and low pressures. The
ideal gas, which should behave uniformly, we
call the perfect gas. With most gases, at low
pressures concentration increases more, and at
very high pressures much less than the rule in-
dicates (see Kinetic hypothesis).
The Relation of the Volume of a Gas to
Temperature. — In the foregoing we have as-
sumed that there were no temperature effects.
But, as a matter of fact, a rise in temperature
will at once produce an increase, and a fall in
temperature a decrease in the pressure of an
inclosed sample of a gas. Hence a record of the
pressure alone will fail to indicate the concentra-
tion definitely, and volume and pressure together
will still leave the amount of material unspecified.
The temperature must, therefore, be given as well.
Our descriptions of different samples of gas
having thus become once more incomparable,
we apply the same kind of remedy as before.
We calculate the vohime which each specimen
of gas would occupy at 0°.
The rule for this calculation may be demonstrated in a rough way
as follows : The large, graduated bulb (Fig. 27) is surrounded by a
vessel which can subsequently be filled with ice water or with water
of any temperature up to 100°. About one-half of its volume is occu-
pied by the gas. The mercury which fills the rest is connected with
a reservoir, so that the levels of the metal can be made alike, and the
pressure of the gas be maintained constantly the same as that of the
atmosphere. When, now, the volume occupied by the gas at 0° is
read, and warmer water is introduced, we find that the volume gains
u|^ of its value at 0° for every degree through which its temperature
rises. If it is cooled below 0°, it loses ^ jg- of its volume at 0° for
every degree through which the temperature is lowered. Observation
gives practically the same value for all gases.
Fro. 27.
INORGANIC CHEMISTRY
The following graphic method will put these facts in a clearer
light. In Fig. 28 we have on the left a thermometer scale divided
r 100°
Vol. of
273 c.c.
- 373 c.c.
- 372 c.c.
r 371 c.c.
- 275 c.c.
- 274 c.c.
- 273 c.c.
- 272 c.c.
- 271 c.c.
Temp.
Abs.
t
373°
372°
371°
L 275°
- 274°
- 273°
- 272°
- 271°
FIG. 28.
3°
2°
1°
into degrees Centigrade. The middle line represents the volumes of
a given sample of gas which correspond with the successive tem-
peratures. If we, for convenience, take a volume of 273 c.c. of a gas
at 0° and warm this through 1°, then at 1° its volume, having gained
g-|7 of its original value, becomes 274 c.c. At 2° it has gained another
2^3 of the volume it had at 0° and becomes therefore 275 c.c., etc.
If cooled below O^it loses ^}7 of the volume, becoming 272 c.c., etc.
If the gas be heated to 100° it gains £f§ of its volume at 0° and
becomes 373 c.c. We might infer that if it were cooled through 273°
from 0° it would lose f ^§ of its volume, in other words, it would dis-
appear. This temperature has not yet been reached, and in any case
all gases would presumably liquefy before reaching it. Our state-
ments apply to ordinary temperatures only, and within them the law
holds with considerable strictness. Now the series of numbers on the
middle line (Fig. 28) are all 273 units larger than the temperatures
Centigrade in the line to the left of the figure. If we alter the gradua-
tion of the Centigrade thermometer by adding 273 to each temperature,
THE MEASUREMENT OF QUANTITY IN GASES 87
we secure the scale on the right, which exhibits degrees of the very
same length as before, but has the additional advantage that the num-
bers expressing temperature are the same as those expressing volume.
This is what we call the absolute scale of temperature. In this arti-
ficial case we started with 273 c.c. at 0° C. (273° Abs.), and the abso-
lute temperature is here always numerically equal to the volume. If
a different volume had been taken at 0°, then the volume assumed by
the gas at each temperature would have borne a constant ratio to the
volumes recorded. Hence, the volumes assumed by a sample of gas
at different temperatures, the pressure remaining constant, are in the
same proportion as the corresponding absolute temperatures. This is
the modern way of stating a fact which was first discovered by Charles
of Paris (1787). Obviously, if the volume remains constant, then the
pressure will be proportional to the absolute temperature. The bare
fact underlying our statement of the law is that all gases suffer equal
increments (or decrements) in volume (or pressure) for equal changes
in temperature.
The discovery of this fact is generally attributed to Dalton, who first published
an investigation of the subject, embodying this result, in 1801, or to Gay-Lussac,
who made a more complete investigation in 1802. Inasmuch, however, as in
Chemistry we have another important principle known as Gay-Lussac's Law, and
several which are connected with the name of Dalton, it is on the whole fortunate
that we are justified in attributing this discovery to Charles.
The application of this law may best be illustrated by an example.
We obtain 200 c.c. of a gas at 20° and wish to know what volume it
would occupy at 0°. To answer the question, we convert the Centi-
grade temperatures to the absolute scale by adding, algebraically, 273
to each. Thus 200 x |J| = the volume at 0° required. No formula
is needed. We simply ask whether the new temperature is higher or
lower than the old one. Here it is loiver. The new volume will
therefore be smaller than the old one. So we take care to place the
smaller number in the numerator.
The law may be put in mathematical form thus : If we make vu the volume at
0° and w, the volume at the temperature £15 then
1 - (\ l t\ - f27'3 + *'
The number 273 + £, is the value of the temperature increased by the number
273, and the sum is spoken of as the absolute temperature. The intervals between
the degrees on this scale are the same as those on the Centigrade thermometer,
but the numbers corresponding to them are all increased by 273. Using 2', for the
absolute temperature corresponding to tf, the last expression, above /becomes equal
88 INORGANIC CHEMISTRY
to v0 ^\-g 2\. Now, similarly, for any other temperature, T2, there will correspond
another volume, B2, and the same relation will hold, namely, »2 = v0 2}j T2. Equat-
ing both values of VQ and calling ^?, the coefficient of expansion, a, we get,
In words, the volumes occupied by a sample of a gas are proportional to tbe abso-
lute temperatures. The last formula enables us to calculate any one of the four
values when we know the other three. It is preferable, however, that beginners
should use the method employed in the illustration given in the preceding para-
graph. Simple mathematical expressions, like the one representing this law, are
not made to save us the trouble of remembering the law itself, and it would be
unfortunate if their use led us to forget it.
An instructive graphic demonstration of the law is given by Ost-
wald in his Principles of Inorganic Chemistry, p. 75.
The behavior of gases in respect to changes of temperature and
pressure are perfectly independent of one another, so that the above
laws may be applied to any example, either in succession, using the
answer for the first calculation in making the second, or simultane-
ously. Thus 200 c.c. of gas at 742 mm. pressure and 20° become
200 x Ifffi X f|f = 183.8 c.c. at 0° and 760 mm.
Mixed Gases. — Two gases at the same temperature, provided
they do not interact chemically, do not interfere with each other's
pressures when mixed. Thus, if they are forced into the same volume,
the pressure of the mixture is equal to the sum of those of the com-
ponents (Dalton's law, 1807). The gases are therefore still thought
of individually, and the share which each gas has in the total pressure
is called its partial pressure. This, like any other gaseous pressure, is
proportional to the concentration of the particular gas in the mixture.
For example, a gas measured over water contains water vapor.
The partial pressure of this, called the aqueous tension (q.v.), which
is definite for each temperature, must be subtracted from the total
pressure. The remainder is the partial pressure of the gas being
measured, and this remainder is used as the pressure of this gas in
any calculation.
Densities of Gases. — The application of the laws of Boyle and
Charles enables us to express the quantities of matter in samples of
gases in a definite and readily comparable manner. In describing
chemical changes, however, it is continually necessary to express
quantities of gases by weight. The relation between volume and
weight for each kind of gas must, therefore, be ascertained. If we
THE MEASUREMENT OF QUANTIFY IN GASES
know, for example, the weight of one liter of each gas at 0° and 760
mm. pressure, conversion to other weiguts, volumes, temperatures,
and pressures can be made. The one-vhousaudth part of this value,
the weight of 1 c.c., is called the density of the gas. Often, however,
the relative weights of equal volumes, with that of air or hydrogen as
unity, receive this name.
For most chemical purposes a high degree of accuracy is not re-
quired. Different arrangements of apparatus are used according to
circumstances. The most direct method is to employ a light flask
provided with a rubber stopper and stopcock (Fig. 29). By means
of an air-pump the contents of the flask are removed, and it is weighed.
This gives the weight of the empty vessel. The gas, whose density
is to be ascertained, is then admitted, and care is taken that it finally
fills the flask at the pressure of the atmos-
phere. The flask is closed and weighed
again. The increase represents the weight
of the gas. At the same time the tempera-
ture and barometric pressure are read.
The volume is determined by displacing
the gas once more from the flask, filling
with water, and weighing again. The
difference in weight between the empty
flask and the flask full of water, in grams,
represents the volume of the content of the
flask in cubic centimeters. This volume
is reduced to 0° and 760 mm. by the rules
discussed above, and we have then a volume
of the gas and the corresponding weight.
To illustrate, let us suppose that the
volume of the flask is 200 c.c. and that it
is filled with oxygen at 20° and 742 mm.
The weight, we will suppose, is found to be
0.26 g. We ascertained (p. 88) by calcula-
tion that at 0° and 760 mm. this volume
would be 183.8 c.c. The weight of a liter
is given by the proportion 183.8:0.26::
1000 :x. Here x = 1.415 g. When the
operation is performed carefully, and the weighing carried to the
nearest milligram instead of the nearest centigram, a result more
nearly approaching the exact one (1.429) may easily be reached.
Fm. 29.
90 INORGANIC CHEMISTRY
To get the density of oxygen referred to hydrogen as unity, we
must divide the answer by the weight of a liter of hydrogen (0.08987 g.).
In the above case the quotient is 15.74. The accepted value is 15.90.
The density referred to air as unity is similarly obtained by dividing
by 1.293, the weight of a liter of air at 0° and 760 mm. pressure.
If a suitable pump is not available, the flask, in this case provided with two
openings, is weighed without preliminary exhaustion. This gives the weight of
the vessel plus that of the air it contains. A continuous stream of the gas is then
passed into the flask until the air has been completely displaced. The vessel is
then closed and another weighing made. Finally the gas is displaced by water,
and a third weighing taken. The temperature and barometric pressure are noted
as usual. The last weighing gives the volume as before. Knowing that one liter
of air weights 1.293 g. at 0° and 760 mm., we may calculate readily the weight of
the air which the flask contained at the observed temperature and pressure. When
this is subtracted from the number obtained in the first weighing, we have the
weight of the empty flask. Subtracting this in turn from the second weighing, we
have the weight of the gas. We obtain thus the weight of a known volume of the
gas at a known temperature and pressure, and finish the calculation as before.
The values of the densities of gases are of great significance in the
chemical point of view. A number of them are given in connection
with the discussion of molar weights (q.v.^).
Vapor Densities of Liquids and Solids. — The densities of
vapors are as important to the chemist as those of gases and, solids
and liquids being more numerous, are even more frequently measured.
The apparatus must be specially adapted to the purpose. In its
simplest form (Dumas' method) it consists of a bulb (160-200 c.c.)
provided with a long, narrow tube (Fig. 30). This corresponds to the
flask used for gases. The bulb is first weighed full of air, and is then
charged with a considerable amount of the substance, usually a liquid.
It is then suspended in a bath whose temperature can be maintained
at some point above the boiling-point of the substance. Boiling water
(100°) serves for liquids boiling below 100°. When the vapor has ex-
pelled all the air, and is no longer seen to issue from the tube, the
bulb contains nothing but the vapor. The tip of the tube is then
sealed with a blow-pipe. The temperature of the bath is taken and
the barometer is read. When the bulb has cooled, it is weighed for
the second time. The tip of the tube is next filed off under water,
which rushes in and fills the bulb almost entirely. A third weighing,
which should include the portion of the neck filed off, gives, after
subtraction of the weight of the vessel, the content of the bulb. From
THE MEASUREMENT OF QUANTITY IN GASES 91
the content and the temperature and pressure at the time of the first
weighing, we calculate the weight of air which the bulb originally
contained. When this is subtracted, we obtain the weight of the empty
bulb. Subtracting that in turn from the second
weighing, we get the weight of the substance
which, in the state of vapor, just filled the bulb
at the temperature of the bath (say 100°) and
at the pressure indicated by the barometer.
From these data we calculate the density re-
ferred to hydrogen (or air) as unity.
The reduction to 0° and 760 mm. pressure
by rule gives, of course, a fictitious result.
The vapor would condense to the liquid form
before 0° was reached, if the cooling were actu-
ally carried out. But the value for the density
as it would be at 0° and 760 mm. has to be cal-
culated to facilitate comparison with the cor- FIO. 30.
responding values for other substances. The
results have no physical significance, but are highly important to the
chemist.
Exercises. — The foregoing cannot be understood unless some
problems involving the laws of gases are actually worked.
1. Eeduce 189 c.c. of gas at 15° and 750 mm. to 0° and 760 mm.
2. Eeduce 110 c.c. of gas at — 5° and 741 mm. to 0° and 760 mm.
3. Convert 500 c.c. of gas at 25° and 700 mm. to 18° and 745 mm.
4. A Dumas bulb full of air weighs 13.3125 g. After being filled
with the vapor of carbon tetrachloride at 100°, it weighs 13.7969 g.
Filled with water it weighs 141.3 g. The barometric reading (corr.)
is 755 mm. What is the vapor density referred to air at 0° and
760 mm. ?
5. The density of a substance referred to air is 3.2. What is
the density referred to hydrogen ? What will be the volume occupied
by 10 g. of the substance at 20° and 752 mm. ?
CHAPTER VII
HYDROGEN
HYDROGEN, although discovered by Paracelsus in the sixteenth
century, was confused with other combustible gases, and its independ-
ent nature was first established by Cavendish in 1766. Somewhat
later (1781), the latter showed that hydrogen when it burned gave water
vapor, of which he condensed a large quantity to the liquid form.
Taken in conjunction with Lavoisier's proof that oxygen was the
active substance in the air (1777), this fact showed that water was a
compound and not a simple substance. The new element was named
hydrogen (Gk. vSwp, water ; yewav, to produce).
Occurrence. — The free element is found, mixed with varying pro-
portions of other gases, in exhalations from volcanoes, in pockets found
in certain layers of the rock-salt deposits, and in some meteorites. The
air contains a mere trace of it, not more than one part in 30,000. Its
lines are very prominent in the spectrum of the sun and of most stars.
In combination, it constitutes about 11 per cent of water. It is an
essential constituent of all acids. It is contained also, in combination
with carbon, in the components of natural gas, petroleum, and all
animal and vegetable bodies.
We have seen (p. 33) that, using the physical unit of weight, which is the
same for all substances, the element hydrogen stands ninth in order of plentiful-
ness. But the chemical work that elements can do should rather be reckoned by
the relative numbers of atomic weights which are available. When Clarke's num-
bers are recalculated to this basis, and the number of chemical unit weights of
oxygen is called 100, hydrogen assumes a position more in harmony with its im-
portance :
Oxygen . . 100.00 Aluminium. . 8.57 Iron. . . 2.97
Hydrogen . 30.10 Magnesium. . 3.29 Calcium . 2.81
Silicon . . 28.52 Sodium . . . 3.17 Potassium . 2.64
Acids. — In making hydrogen, the acids are used almost exclu-
sively. Hence, some statements in regard to their nature must be
made before we can use them intelligently.
92
HYDEOGEN 93
The common acids are hydrochloric acid (HC1, Aq), sulphuric acid
(H2S04, Aq), and nitric acid (HN08, Aq). The usual forms are mix-
tures containing water, the variable amount of the latter being indi-
cated by the symbol Aq. The first is a solution of a gas, hydrogen
chloride. The " pure concentrated " hydrochloric acid used in lab-
oratories contains nearly as much of the gas (39 per cent by weight)
as the water can dissolve. When heated, it readily gives up part of
the gas, and the effervescence attending this must not be mistaken for
evidence of chemical action. The " commercial " acid contains im-
purities and is also less concentrated. The "concentrated" sulphuric
acid is an oily liquid containing practically no water. The " com-
mercial" sulphuric acid contains 6 to 7 per cent of water, besides
impurities. The " pure concentrated " nitric acid contains 70 per cent
of liquid nitric acid. The " commercial " acid, 53 to 62 per cent and
impurities in relatively small amounts. Acetic acid (HC2H302, Aq)
is a solution of a liquid in water. All the "dilute" acids contain 90
to 95 per cent of water. The water, as a rule, takes no part in the
chemical changes in which the acids are concerned, and is therefore
omitted from the equations.
The name " acid " is restricted to one class of substances having cer-
tain definite characteristics. Hydrogen is the only essential constitu-
ent of all acids. When free from water they do not conduct electricity.
Their aqueous solutions have a sour taste and change the color of
litmus from blue to red. We shall note presently two other prop-
erties which acids show when dissolved in water : — They conduct and
are decomposed by the electric current, and their hydrogen (or one
unit weight of it in the case of acetic acid) is displaced by certain
metals.
In describing the chemical behavior of acids, we speak of the
material combined with the hydrogen as the negative radical (see next
section). Thus the negative radicals in the above acids are Cl, S04,
N08, and C2H802, respectively. The first (Cl) is a simple radical, the
others complex. In many interactions the complex radicals move as
units from one state of combination to another.
Preparation of Hydrogen by Electrolysis. — A supply of
free, unmixed hydrogen not existing, we are compelled to prepare it
from compounds by the two general plans (p. 63) used for liberating
simple substances. When the first plan, the direct application of
energy, is employed, we find that electricity serves the purpose best.
94
INORGANIC CHEMISTRY
The common compounds of hydrogen, like hydrogen chloride and water, are
not easily decomposed by heat, and in most cases, at best, a mixture of gases would
be obtained. The difficulty in separating the resulting gases makes the use of this
form of energy unsuitable. On account of its ability, not only to liberate the con-
stituents from combination, but also to deliver the positive and the negative parts
of the compound in separate places, electricity alone is available.
If we dissolve any acid in water, and immerse the wires from a
battery in the solution, bubbles of hydrogen begin to appear on the
negative wire (the cathode) and rise to the surface. All the other con-
stituents, whatever they may be, are
attracted to the positive wire (the anode)
and are set free in some form at its
surface. It is on account of this be-
havior of the radicals of acids that they
are known as "negative" radicals. An
apparatus devised by Hofmann (Fig. 31)
enables us to secure the hydrogen, which
ascends on the left and accumulates at
the top of the tube, displacing the solu-
tion. The other products, if gaseous,
occupy a separate tube on the right
side. The solution displaced by the
gases is forced down and mounts into
the bulb behind. The current of elec-
tricity flows from one wire to the other
through the cross-tube. In the typical
case, with a properly selected acid, the
production of hydrogen ceases when
the acid is all decomposed. The water
alone is an almost complete noncon-
ductor, so that the flow of the electricity
practically ceases at the same time. If
the operation does not come to rest in
this way, its continuance is due to the
regeneration of conducting substances by the interaction with the
water of the materials of the radical liberated at the positive electrode.
When hydrochloric acid is used, we have a close approximation to
the typical case. The equation is :
HC1 — » H (neg. wire) + Cl (pos. wire),
and the chlorine, a soluble gas, remains dissolved in the water near one
FIG. 31.
HYDROGEN 95
pole. When sulphuric acid is employed, the equation is :
H2S04 --» 2H (neg. wire) + S04 (pos. wire),
and the S04 interacts with the water (see Discharging potentials), thus :
S04 + H20 -» H2S04 + 0.
Hence oxygen conies off, and the substance regenerated is here sul-
phuric acid itself. The final results are, therefore, the liberation of
hydrogen and of oxygen and the localization of the regenerated acid
round the positive electrode.
It is worth noting that the acids and water, taken separately, are
all nonconductors. The fact that the mixture does conduct, concomi-
tantly with the decomposition of the acid, is therefore highly sug-
gestive. Solution in such cases must be something more than a mere
physical change of state of aggregation (see lonization).
It is commonly asserted that water is decomposed by a current of electricity
This is true in the sense in which we might say that a man can carry off a hill.
He may eventually remove it, if you give him time. The action of electricity upon
the purest water is exceedingly slow, on account of the very minute conductivity
for electricity which it possesses. Common distilled water owes its appreciable
capacity for conducting chiefly to traces of an acid, namely, carbonic acid, which
it contains. Even when the water is saturated with carbonic acid, however, dilute
sulphuric acid has a conductivity of the order of a thousand times greater. For
our present purpose, therefore, water is declared to be a nonconductor. Yet, as
we shall see, the conductivity of pure water, small as it is, has to be taken into
consideration in certain cases (see Hydrolysis and Electromotive chemistry).
Preparation of Hydrogen by Displacement from Diluted
Acids. — By use of the second plan for liberating elements (p. 63), hy-
drogen may be obtained from acids, through substitution for it of some
element with which the negative radical will unite.
The acids must be diluted with water before rapid action occurs.
The substances capable of displacing hydrogen from them are certain
of the metals, like zinc, iron, and aluminium. The hydrogen escapes
in bubbles, and evaporation of the remaining liquid gives in dry form
the compound of the metal with the other constituents of the acid.
Thus, with zinc and sulphuric acid, zinc sulphate is produced :
Zn + H2S04 -» 2H + ZnSO4 ;
and with tin or aluminium and hydrochloric acid we get stannous
chloride or aluminium chloride :
96
INORGANIC CHEMISTRY
Sn + 2HC1
(Tin)
Al + 3HC1
. 2H + SnCl2
(Stannous chloride)
. 3H + A1C1,
The water undergoes no change during the action, although its presence
is essential. It is simply a part of the apparatus. Any acid may be
used, although with many the action goes on very slowly. In all
cases the plan of the action is the
same : the metal is said to displace
the hydrogen (see below).
The apparatus for generating
small amounts of hydrogen (Pig. 32)
is arranged so that additional acid
may be added through the thistle
or safety tube. This avoids ad-
mission of air. With a Kipp's
apparatus (Fig. 33) the gas may be
made on a larger scale and its de-
livery can be regulated. When the
stream of gas is shut off by the
stopcock, the pressure of the gas,
as it continues to be generated,
drives the acid away from the metal
and up into the globe above, so
that the action ceases. Yet the
action is ready to begin again the
moment any portion of the stored
gas is drawn off for use.
A rather sharp line can be drawn between those metals which dis-
place hydrogen from dilute acids and those which, like mercury,
silver, and gold, do not (see Electromotive series of the metals).
Contact of the zinc or iron with an inactive metal, like platinum,
always hastens the interaction, and therefore renders the evolution of
the hydrogen, more conspicuous. Such an arrangement is called a
couple, and its efficiency depends on the electric states of the two
metals (see Solution tension).
When water is not used along with the acid, the latter is either
inactive or undergoes a different sort of chemical change. Thus, dry,
gaseous or liquefied hydrogen chloride hardly interacts at all with
zinc. Pure, concentrated sulphuric acid, on the other hand, although
almost unaffected by zinc in the cold, is violently decomposed when
FIG. 32.
HYDROGEN
97
heated. The action, however, is not a simple displacement (see below)
of the hydrogen. The oxygen is removed from a part of the acid, and
water and hydrogen sulphide are formed :
4Zn
4H20
H2S.
Preparation of Hydrogen from Water. — Every one of the
metals which act on dilute acids will also displace hydrogen from
water, and no others will do so. Only
the more active metals, like potassium
and sodium, which would act with
uncontrollable vigor on dilute acids,
can displace the hydrogen rapidly
from cold water. Magnesium and zinc
show obvious action on water at 100°
only, and are much assisted by contact
with another metal. If any action is
to be perceived in the cold, the iron,
nickel, zinc, and magnesium have to be
used in a state of fine powder, with
great surface.
In all cases in which cold or boil-
ing water is employed, the hydrogen
of the water is not completely dis-
placed. The metal forms an hydrox-
ide, such as sodium hydroxide or
magnesium hydroxide :
Na + H20 -> H + NaOH,
Mg + 2H.20 -> 2H + Mg(OH)2.
Sodium, which is one of the con-
stituents of common salt, may be used
to illustrate this sort of action. As
it is lighter than water, it must be
held under the surface by means of a piece of wire-gauze in order
that the gas may be collected (Fig. 34). Most of the water serves
the mechanical purpose of permitting the collection of the gas, and
only a small fraction of it takes part in the change. The solution
has a soapy feeling and turns litmus from red to blue. This color
reaction is the precise opposite of that of acids (p. 93). Substances
FIG. 33.
98
INOEGANIC CHEMISTRY
causing these two effects are called alkalies. Evaporation of the
resulting very dilute solution reveals the sodium hydroxide, the
alkali, as a white solid.
With steam at a red heat, metals like iron, zinc, and magnesium
FIG. 34.
interact vigorously. The metal is placed in a tube in which it can be
strongly heated (Fig. 35X The steam, generated in a flask, enters at
TTP— .
one end of the tube, and the hydrogen passes off at the other. Since,
at a red heat, all hydroxides, except those of potassium and sodium,
HYDROGEN 99
are decomposed into an oxide of the metal and water, Mg(OH)2 — >
MgO + H20, the oxides are formed in this case :
Mg + H20 -» MgO + 2H.
Iron gives the magnetic oxide, Fe304. Hence, to make the equation,
four unit- weights of oxygen, and therefore four formula- weights of
water, are required :
4H20 + 3Fe -» Fe304 + 8H.
Tests. — The effect of a solution of an acid (p. 93) or an alkali (see
above) upon litmus solution is an illustration of a chemical test. In
this case a mere trace of a highly colored body dissolved in water
interacts with a correspondingly minute trace of the acid or alkali
under examination. Yet the product has such marked color that its
presence is immediately recognizable. Tests for various kinds of
materials are eagerly sought by chemists for use in identifying un-
known substances. Any property which is so conspicuous as to be
apparent when very little material is concerned will serve as a test.
Thus the precipitation of calcium carbonate (p. 68) enables us to
recognize the presence of carbonic acid. The test consisted in the
addition of lime-water.
The Other Ways of Preparing Hydrogen. — For special pur-
poses, hydrogen may be made by boiling an aqueous solution of sodi-
um hydroxide with aluminium turnings, when sodium aluminate is
formed : Al + NaOH + H20 — > NaA102 + 3H ; also by heating pow-
dered zinc and dry sodium hydroxide, the product being sodium zin-
cate : Zn -f 2IS"aOH — > Na2ZnO2 + 2H. It may likewise be made by
electrolyzing solutions of compounds of metals which, when free, dis-
place hydrogen from cold water. Thus electrolysis of sodium chloride
or sodium hydroxide in aqueous solution liberates chlorine at the posi-
tive wire and hydrogen and sodium hydroxide (p. 97) at the negative
wire (see Discharging potentials, Chap, xxxviii).
Displacement. — We now have before us illustrations of two sub-
varieties of the third kind (p. 15) of chemical change. In this kind,
compounds are decomposed and the parts combine in a new way.
The first sub-variety was double decomposition, as in the action of
sodium chloride upon silver nitrate :
NaCl + AgN03 -» AgCl
100
INORGANIC CHEMISTRY
In this class of cases two compounds interact, each splits into the
radicals of which it is composed, and two neiv compounds are formed
by union of the radicals crosswise. The actions used in the prepara-
tion of hydrogen differ from these inasmuch as one compound and one
element interact, the compound splits into its radicals, and one com-
pound and one free element are produced:
Zn + H2S04 -> ZnS04 + 2H,
Zn + 2KaOH ->Na2Zn02 + 2H.
The former element, here the zinc, is said to displace the latter, here
the hydrogen, from combination. In double decomposition there
is an even exchange, the sodium,
for example, giving up one
radical (01), and getting another
(N03), whereas in displacement
one element gains a radical
while another loses it, the zinc,
for example, giving up nothing
but getting S04, while the
hydrogen loses S04, and gains
nothing in return.
It should be noted, that
although in the former of the
above illustrations the Ag may
be said to displace the Na from
combination with the Cl, the
FlG- 36- term displace is \ised techni-
cally only when the element
displaced escapes in the free condition.
Purification of Gases. — Hydrogen made in any of the above
ways is impure (p. 34). As made by the first three methods, a good
deal of water vapor is mixed with it. Other impurities, like hydrogen
sulphide and arsine, come from the action of the acid on foreign mate-
rials in the zinc (p. 95). Some of the acid, if it is volatile, will also be
taken over with the gas. When the object for which the gas is being
made demands it, we must know what the impurities to be expected
are, and take proper means of removing them.
Gases are freed from aqueous vapor by means of calcium chloride
or concentrated sulphuric acid, which greedily absorb moisture. The
HYDROGEN
101
former is used in granulated form in straight or bent tubes (Fig. 36).
The latter is applied by saturating pieces of pumice-stone with the
acid and filling similar tubes with the fragments. Or the acid may
be placed in a gas washing bottle (Fig. 37). For extremely complete
drying, a tube may be filled with phos-
phoric anhydride sifted upon glass
beads or glass wool. Forethought must
be used to avoid a drying agent which
will interact with gas. The longer the
gas remains in contact with the drying
agent, the more perfect, up to a certain
limit, is the purification effected. In
all cases, the stream of gas must pass
slowly.
Particles of liquid or solid matter
are always carried along by freshly
made gases. These will pass with the
gas through sulphuric acid without
being affected. A plug of cotton or
of glass wool in some part of the tubing
is required to arrest them.
Valence. — We shall gain nmch
help in the making of equations if we FIG. 37.
now introduce and bring into relation to
the symbols a conception for which the remarks about atomic weights
(p. 50) have paved the way. It will have been observed that the com-
position of the chlorides of aluminium,, tin, and sodium are represented
by the formulae A1C13, SnCl9, and NaCl respectively. Again, the hy-
droxide of sodium is NaOH, while those of magnesium and calcium are
Mg(OH)2 and Ca(OH)2. In making equations we constantly need to
know whether the chloride of an element, say magnesium, is MgCl, or
MgCl,, or MgCl3, or MgCl4, etc., and whether its sulphate is MgS04,
or Mg2S04, or some other combination of the symbols. To answer
questions like this it is not necessary to know the formula of every
compound of each element : the apparent disorder of these numbers
can be reduced to rule, and the reader should endeavor thoroughly to
master the rule before going farther.
If the method by which the atomic weights were derived from
equivalents (p. 50) is now reexamined, the nature of this rule will be
102 INORGANIC CHEMISTRY
seen. It was found, for example, that 9.03 parts of aluminium (p. 50)
combined with the equivalent weights of the other elements, and
therefore with 35.45 parts of chlorine. If this weight of aluminium
had been accepted as the final unit (the atomic weight), then it would
have been represented by the symbol Al, and, since Cl stands for
35.45 parts of chlorine, the formula of the chloride would have been
A1C1. In point of fact, however, a number three times as large as the
equivalent, namely, 27.1, was chosen as the atomic weight of alumin-
ium, and symbol Al stands for this triple quantity. If the equivalent
of chlorine had also been tripled in making its atomic weight, the
amounts represented by the symbols would still have been chemically
equivalent, and the formula would still have been A1C1. But the
equivalent of chlorine was left unaltered. Hence, to get the equiva-
lent amounts (i.e., the actual combining quantities) of the two elements,
we must have 301 with 1A1. The formula is thus A1C18. Now, it
is evident that this tripling of the equivalent of aluminium will affect
the formulae of all its compounds. Whenever it is combined with an
element which, like chlorine, has identical equivalent and atomic
weights, the formula of the compound will be of the form A1X3. In
accordance with this we have the bromide AlBr3. In making the
formulae of compounds of aluminium, the chief thing to be kept in
mind, therefore, is the fact that its atomic weight contains three
equivalents and always combines with three equivalents of another
element. This fact we state by saying that the valence of the atomic
weight of aluminium is three, or simply that the element aluminium is
trivalent.
Similarly, the equivalent of tin is 59.5 and its atomic weight is
119. This atomic weight therefore contains two equivalents* of tin
and combines with two equivalents of any other element. Hence,
the formula of a compound of tin with an element of the chlorine
class will be SnX,. Thus tin is bivalent. In like mariner the equiva-
lent of sodium is 23, and this number was not altered in making the
atomic weight. Hence, the symbol Na stands for one equivalent, and
the formula of the compound with chlorine is NaCl. Elements whose
atomic weights are identical with their equivalents are described as
univalent.
Thus the valence of an element may be defined as the number of
equivalent -weights contained in its atomic weight. Arithmetically
it is the integer by which the equivalent weight was multiplied in
forming the atomic weight. The above explanation shows that
HYDROGEN 103
we may define the valence of an element also as the number of
atomic weights of a univalent element, with which its atomic
weight will combine. A more complete definition will be given pres-
ently.
Sometimes the valence is indicated in the symbol thus : A1IIT>
Sn11, Na1, Cl1, Br1. The table of atomic weights (p. 50) shows the
following additional cases: On, Cu11, S", Hg", H1, Fe", Mg", CIV.
With the help of this list the formulae of compounds may easily be
made. Thus, oxygen is bivalent, and an atomic weight of oxygen,
represented by O, will combine with two atomic weights of a univalent
element as in OnH2r (water), or with one atomic weight of a bivalent
element as in OirSn" (stannous oxide), 0"Hg" (mercuric oxide), 0"Cun
(cupric oxide), OnMgn (magnesium oxide). Again, carbon being quad-
rivalent, the atomic weight combines with four units of chlorine and
of hydrogen in C^ClJ (carbon tetrachloride) and CIVH4Z (methane),
or with two units of oxygen and of sulphur in CIV02" (carbon dioxide)
and CIVS2JI (carbon disulphide). When it combines with a trivalent
element, equal numbers of equivalents of each element must be used,
as in C8IVAl4m (aluminium carbide), where C3 and A14 contain twelve
equivalents each. This method, with exceptions to be noted below,
will give the formulae of all compounds containing only two elements
— so-called binary compounds.
The above mode of handling valence is based upon the notion of
combination in equivalent proportions. Another variety of chemical
change, namely displacement (p. 99), is often of assistance in enab-
ling us to determine the valence of an element. It will be noted
that when Al acted upon hydrochloric acid (p. 96) and combined
with 301, it necessarily displaced the 3H with which the 301 was
formerly united. It was equivalent to 3H for the purpose of holding
301 in combination. It is from this aspect of the relation that the
word " valence " comes. Al is equi-valent to 3H, and, H having the
unit valence, Al is trivalent. Similarly, since one atomic weight of
zinc, represented by the symbol Zn, displaces 2H (p. 95), zinc must
be bivalent. Combining this with the former conception, we reach the
common definition of the valence of an element : The valence of the
atomic weight of an element is the number of atomic weights of hydro-
gen, or of some other univalent element, which it combines with or
displaces.
Formulae are often written so as to show the valence plainly.
Thus, K — 01 indicates, by the single line, that each element is uui-
104 INORGANIC CHEMISTRY
valent. Two or more lines meeting at a symbol, indicates that that
element is bivalent or trivalent :
/ Cl pi Fe=O
A1-C1, Mg'pt, Mg = 0, 0 = Fe-0-Fe = 0, ' >Q.
\C1 Fe=O
Here Al and Fe are trivalent, Mg and 0 bivalent. The lines may be
drawn in any direction, as the last two formulae show. As many
lines proceed from each symbol as will represent its valence. The
resulting structures are called graphic formulae.
The Valence of Radicals. — In the preceding section it has been
seen that the valence of elements can easily be determined when they
are present in binary combination. This is no longer the case when
more than two elements are united together. A study of chemical
changes shows, however, that even here the conception of valence can
still be employed. In the interaction of zinc with dilute sulphuric
acid:
Zn + H2S04 -^ ZnS04 + 2H
the group S04 passes as a whole from combination with 2H to com-
bination with Zn. Hence, although we cannot by inspection deter-
mine the valence of sulphur, we do perceive that the radical SO^ as a
whole, must be bivalent. It occurs, in fact, in all sulphates, as Ag2S04,
MgS04, and A12(S04)3, and in the interactions of these substances it
usually passes intact from one state of combination to another, and
behaves as if it were a unit of a single element of valence two. Again,
in the interaction of salt with silver nitrate (p. 99), we observe that
the radical N08 is univalent. Still again, the compositions of the com-
pounds CaCl2 and Ca(OH)2 show that the radical OH (hydroxyl) is
univalent. The formula NaOH leads to the same conclusion.
This addition to our ideas enables us greatly to extend the list of
substances of which we can write the formulas. Thus, the hydroxides
all contain (OH)1, e.g. A1III(OH)3I (aluminium hydroxide), Sn^OH),1
(stannous hydroxide), Cu^OH)/ (cupric hydroxide). The nitrates
all contain (NO,)1, as : ^(NOg)1 (nitric acid), Mg'^NO,,),1 (magnesium
nitrate). It is to preserve the identity of the radicals that we write
them in brackets and place the factor outside, instead of using the
forms A1O3H3, MgN206, and so forth. In fact, we regard as binary
compounds, substances which commonly interact as if the radicals
were single elements. In all actions in which the radicals preserve
HYDROGEN 105
their integrity, the conception of valence proper to binary compounds
may be used for the more complex compounds also.
How to Ascertain the Valence of an Element or Radical. —
The above shows that the valence of one element or radical may al-
ways be ascertained by examination of the formula of a compound
containing another element or radical of known valence. Thus, when
we know the formula of sodium iodide to be Na1!, or that of hydro-
gen iodide to be HI, we infer that iodine is univalent. The formula
of silica (sand) Si02IJ shows silicon to be quadrivalent, and indicates
that the chloride must be SiCl4. Similarly the formula of calcium
carbonate Ca"C03 shows that the radical C08, which is common to all
carbonates, must be bivalent. Hence, the chemist does not memorize
the valences themselves ; he recovers them when needed by recalling
the formula of a compound containing a more familiar element or
radical.
It is absolutely essential that correct valences should be used
in constructing equations, and, at first, the student will find the task
by no means easy. He should give special attention to this mat-
ter until, by solving the exercises at the end of this chapter, and by
careful examination of all the equations encountered in the text, he
has mastered the subject.
Multiple Valence and Exceptional Cases. — Some elements
show more than one valence. This is as much as to say that an atomic
weight of such an element may form stable compounds with two, or
even more different numbers of equivalents of another element. This
fact has already been mentioned, for it is implied in the law of mul-
tiple proportions (p. 41). Thus an atomic we'ight of tin may form
two different compounds with chlorine, namely, SnnCl, (stannows
chloride) and SnIVCl4 (stanmc chloride). Tin behaves in the same way
towards other elements, however, and we have a series of stannaus
compounds, SuO, SnBr2, and so forth, and a corresponding series of
stanm'c compounds, Sn02, SnBr4, etc. Two different valences of the
same element or radical gives rise therefore to two complete sets of
compounds. The nomenclature used to distinguish the two series
has been discussed before (p. 70). As a rule, an element passes from
one form of combination to another without change of valence. But
compounds of elements like tin can also undergo changes in course
106 INORGANIC CHEMISTRY
of which the valence alters. Cases of this kind will be considered
when they arise (see Preparation of chlorine).
The regular valence of an element cannot be learned by examining the com-
position of a compound chosen at random. Thus FeS, H2S, HgS, and other com-
pounds show sulphur to be bivalent. There is also a series in which sulphur is
sexivalent, as in SO3. But the compound S203, in which sulphur appears to be
trivalent, is an isolated case. Again, FeO, FeS, FeCl2 show iron to be bivalent,
and FeCl3, Fe2(S04)3, etc., show it to be also trivalent. But Fe3O4, the magnetic
oxide, is an exception. Valence has to do mainly with chemical interactions, in
which the element either passes from one state of combination to another without
change of valence, or goes over into a compound of another regular series with
another regular valence- It is not a matter of statics. Hence, questions as to the
magnitude of the valence in isolated compounds like Fe3O4, N20, and so forth, are
at present of minor importance.
A definition of valence differing from those given above is preferred by many
chemists. The atomic weight of a univalent element can hold but one unit of an-
other element in combination. Thus, the weight of chlorine represented by Cl can
hold but one H or one Na in combination. Anatomic weight of a bivalent ele-
ment, although it combines with but one unit of another bivalent element, may
hold as many as two units of a univalent element in combination. But it cannot
hold more. A unit of a trivalent element, however, may hold as many as three
units, provided the other element is univalent. In this point of view the valence
of an element is the maximum capacity of its atomic -weight to hold
atomic weights of other elements in combination.
Valence is often defined as the power of the atomic weight of one element, to
hold units of other elements in combination. But the word power suggests that
valence is a measure of the force with which the elements are held together,
whereas it has to do with the quantity of matter only. Gold is trivalent, but holds
chlorine with incomparably less force than does sodium which is only univalent.
Physical Properties of Hydrogen. — Some of these may be
given in tabular form :
Colorless Crit. temp., about — 234°
Tasteless Sp. Ht. (gas), 3.4
Odorless Boiling-point, -i 252.5°
Density (air = 1), 0.0695 Melting-point (58 mm.),- 260°
Density (H = 1), 1 Sol'ty in Aq, 1.9 vols. in 100 (14°)
Wt. of 1 1., 0.08987 g.
Air is 14.5 times as heavy, hence the gas may be poured upwards
and is used for filling balloons. A liter flask filled with air requires
about 1.2 g. to be added to the tare to restore the balance when the
air is dispaced by hydrogen. Its specific heat is about seventeen
times that of oxygen (0.2). Its thermal conductivity is greater than
that of any other gas. Hence a wire, raised to incandescence in
HYDROGEN 107
air by means of an electric current, cannot be kept at a red heat, even,
by the same current in hydrogen.
Hydrogen was first liquefied in visible amounts by Dewar (1898).
The liquid is colorless, and, when allowed to evaporate rapidly under
reduced pressure, freezes to a colorless solid. All other gases except
helium solidify easily when led into a vessel surrounded by liquid
hydrogen.
Hydrogen is absorbed, for the most part in a purely mechani-
cal way, by many metals. Heated iron will take up 19 times its
volume of hydrogen. Under similar conditions gold takes up 46 vol-
umes, platinum in fine powder 50 volumes, palladium 502 volumes,
and silver none. The maximum absorbed by palladium under favor-
able conditions is 873 volumes. It is still a question whether, in the
case of palladium, a part of the gas is not in combination.
Diffusion. — If a volume of gas is inclosed at one end of a cylin-
der, the rest of which is entirely empty, and is suddenly released from
this confinement, it spreads with extreme speed so as to occupy the
whole of the cylinder to an equal degree. This spreading is not an
effect of gravitation, since it takes place upwards or downwards
with equal celerity. The same phenomenon is observed when, in every-
day life, a bottle of scent is opened. The vapor, on escaping, begins
to penetrate in all directions through the room, showing its presence
by its odor. The motion, as this instance shows, takes place through
a space occupied by another gas more slowly than, but just as surely
as, when the space is empty. The material of gases has in fact an in-
dependent power of locomotion. The resulting phenomenon we call
diffusion. It is constant in rate for each gas under like conditions,
and hydrogen has the greatest speed of diffusion of all the gases.
The interdiffusion of gases and the absence of gravity effect may be shown
.simultaneously. A jar is filled with carbon dioxide and a jar of air is inverted
and placed mouth to mouth with the other. After a few minutes, and in spite of
the fact that carbon dioxide, measured in bulk, is one-half heavier than air, as
much of each gas will be found in the cylinder of the other as in its own. Lime-
water (p. 68) will show the presence of carbon dioxide in the upper jar. The phe-
nomena of diffusion must not be confused with cases like the pouring of hydrogen
upward to displace air in an inverted jar. In this case the gas flows en masse, and
the gravity effect is the very one on which we depend for the success of the experi-
ment. It is when hydrogen scatters itself in a somewhat slower way, and down-
ward and sideways as well as upward, that we have diffusion. The word indicates
the scattering rather than the flowing nature of the phenomenon.
108
INORGANIC CHEMISTRY
The different rates of diffusion of different gases are easily shown
by comparing their several speeds with that of air, when both pass
through a wall of unglazed, porous porcelain.
The porous cylinder A on the left (Fig. 38) contains air and is con-
nected with a wide tube which dips beneath the surface of the water.
When a cylinder H containing hydrogen is brought over it, rapid
escape of gas takes place through the water, showing that a rise in
pressure has taken place inside the porous vessel. Before the cylinder
of hydrogen approached it, the air was moving both outwards and in-
wards through the porcelain, but,
being the same air, the speed of
motion was equal in both direc-
tions, and therefore the pressure
inside was not affected. It is im-
portant to note that there was at
no time rest, there was simply
equal motion in both directions.
When the hydrogen atmosphere
surrounded the cylinder, the hy-
drogen gas moved more rapidly
into the cylinder than the air in-
side could move out, and hence an
excess of pressure quickly arose
in the interior.
The cylinder on the right is
similar, but a vessel C filled with
carbon dioxide, a gas heavier than
air (density 1.53, air = 1), sur-
rounds it. Here the air moves
out faster than the gas can move
in, a reduction in pressure takes
place, and the water rises.
Exact measurement shows that the lighter a gas is in bulk, the
faster its parts move by diffusion in any direction. The rate is
inversely proportional to the square root of the density of the gas.
Thus, for hydrogen and air it is in the ratio VI : V.0695, or 3.8 : 1.
For air and carbon dioxide it is Vl-53 : VI, or 1.24 : 1.
FIG. 38.
Chemical Properties of Hydrogen. — Hydrogen, delivered from
a jet, burns in air or pure oxygen. A cold vessel held over the almost
HYDROGEN 109
invisible blue flame condenses to droplets of water the steam that is
produced. Although the flame gives little light, it is exceedingly hot.
Platinum melts in it easily. In a closed space it produces a tem-
perature of over 2500°. When hydrogen and oxygen are mingled in a
suitable burner, and the flame is allowed to play on a piece of quick-
lime, the latter becomes white-hot at the spot where the flame meets
it. This result is called a calcium light or lime light.
When the gases are mixed in a glass vessel, the chemical action is
very slow at ordinary temperatures, no perceptible amount of union
occurring in a period of five years. If the mixture is sealed up and
kept at 300°, after several days a small part is found to have com-
bined to form water. At 518°, hours are required before the union is
complete. At 600° the interaction is rapid, but not explosive. At
700° the combination is almost instantaneous. Hence contact with
a body at a bright-red heat (p. 73) is required actually to explode the
mixture.
These facts illustrate the effect of temperature on the speed of
chemical changes (p. 72). A rough calculation shows that, since
interactions lower their speed to half its value for every depression of
10° in temperature, at ordinary temperatures this union can hardly
make easily perceptible progress in less than a thousand million years.
This effect of temperature, therefore, accounts for the apparent absence
of action in the cold gases.
Finely divided platinum,* when held in the mixture, hastens the
action in the part of the gases in contact with it. The heat of their union
raises the temperature of neighboring portions and causes explosion of
the mass. The platinum is simply a catalytic agent (p. 75) and
remains itself unaffected. Its role is to increase prodigiously the
vanishingly small speed of the union between the cold gases.
Hydrogen unites directly with a minority only of the simple sub-
stances. It combines rapidly with oxygen, chlorine, fluorine, and
lithium, and more slowly with a few others.
When these elements, especially the first two, are already in com-
bination, hydrogen may still sometimes displace the material with
which they are united. Thus, when one of the oxides of copper or of
iron is heated in a tube through which hydrogen flows, the latter com-
* The most convenient form is obtained by dipping asbestos in a solution of
chloroplatinic acid, and heating it in the blast-lamp. The fibers are covered with
a thin film of the metal :
H2PtCL— >Pt + 4Cl + 2HCl.
110 INORGANIC CHEMISTRY
bines with the oxygen to form water, and the metal is liberated. To
make the equations, we set down first the formulae of the substances
used and produced in each case :
CuO + H -> H20 + Cu,
Fes()4 + H->H20 + Fe.
We then observe that, for each atomic weight of oxygen, 2H will be
required, and amend the equations thus :
CuO H- 2H -> H20 + Cu, (1)
Fe304 + 8H -> 4H20 + Fe.
Then we make the amount of iron produced equal to that taken :
Fe304 + 8H -» 4H2O + 3Fe. (2)
Thus (1) and (2) are the final equations.
These interactions are classed as displacements. In describing
them the chemist would also say that the hydrogen has been oxidized
and that the oxide of the metal has been reduced (p. 72).
An Inapt Use of the Word "Affinity "in Explanation of
Chemical Actions. — It has passed into the common language of
chemistry that actions like the reduction of magnetic oxide of iron,
just mentioned, are "explained" by saying that the hydrogen has a
greater tendency to unite with oxygen, or has a greater affinity for it,
than has iron, and therefore removes the oxygen from combination
with the latter. Plausible as this statement seems, it would be in
most cases, as here, quite incorrect. Under the modes of preparing
hydrogen, we spoke of the action of steam upon iron (p. 99), and gave
the equation : 3Fe + 4H20 — > Fe304 + 8H. To use consistently this
handy method of explaining chemical change by the help of the word
" affinity," we should have to say that the hydrogen has a less affinity
for the oxygen than has iron, and therefore hydrogen is set free and
oxide of iron is formed. It will be seen that this statement is in direct
contradiction to the one made above. Both cannot be true. The fact
is that both are based upon an assumption which is incorrect — the
assumption, namely, that the displacement of one element by another
is always an evidence of the greater affinity of the latter. Until the
means of truly measuring affinity, or, as we prefer to call it, activity,
have been explained, we shall do well, as far as possible, to avoid using
the word.
HYDROGEN 111
The action of catalytic agents is itself a refutation of this blundering assump-
tion. Putting a little platinum in a mixture of oxygen and hydrogen cannot add
to the energy contained in these substances, and cannot therefore increase their
intrinsic tendencies to unite. Yet in its presence an almost nonexistent action
becomes suddenly explosively violent. There are other far more potent factors
than affinity which determine the direction and speed of many chemical changes
(see Chemical equilibrium).
In this connection, it is worth noting that, while increasing the speed of a train
or a ship requires a great addition to the energy expended, and is very costly, in-
creasing the speed of a chemical change requires the expenditure of no energy
whatever. The employment of a couple (p. 96) or a catalytic agent adds nothing
to the energy the separate bodies possessed before they were mixed. And the cata-
lytic agent is recovered unchanged and as efficient as ever at the end. Theoretically,
therefore, these agencies cost nothing. The increased speed in the formation of
the products is obtained gratis. The contact method of making sulphuric acid
(q.v.) illustrates the way in which commerce has taken advantage of this fact.
The Speed of Chemical Actions: a Means of Measuring
Activity. — The speed of a chemical action is measured by the num-
ber of atomic or formula weights of the substance undergoing change
in a given time. Now, one means of measuring the relative chemical
activities of several substances, and, therefore, of the relative amounts
of available chemical energy they contain (p. 26), is to observe the
speed with which they undergo the same chemical change (p. 28).
Thus we may compare the activities of the various metals by allowing
them separately to interact with hydrochloric acid and collecting and
measuring the hydrogen liberated per minute by each. It will be seen,
even in the roughest experiment, that magnesium is thus much more
active than zinc. The comparison must be made with such precau-
tions, however, as will make it certain that the conditions under which
the several metals act are all alike. Thus, in spite of the heat evolved
by the action, means must be used, by suitable cooling, to keep the
temperature at some fixed point during the experiment, for all actions
become more rapid when the temperature rises (p. 72). Again, the
pieces of the various metals must be arranged so that equal surfaces
are exposed to the acid in each case ; for pieces of the same metal,
having, of course, the same intrinsic activity, will nevertheless give
hydrogen more rapidly the larger the surface they expose. Equal
weights of zinc will finally give equal weights of hydrogen ; but if one
of them is in the form of foil while the other is a cylinder, the former,
although it will not last so long, will give much more hydrogen per
minute. Still again, the portions of hydrochloric acid must contain
the same percentage of hydrogen chloride in each case, for the metal
112 INORGANIC CHEMISTRY
will secure the acid it needs with less delay in a more concentrated
solution than in a less concentrated solution, and in the former case
will therefore displace hydrogen more rapidly. When these and other
precautions have been taken, a true comparison of the relative activi-
ties of the metals with respect to this particular action may be made.
It is found that the order in which this comparison places the metals
is much the same as that in which they are placed by a study of other
similar actions. This is natural, since we are really comparing, in each
case, the amount of chemical energy in each metal. A single table
suffices, therefore, for all purposes. This table is given in connection
with a more exact method of comparing the activities of the metals
(see Electromotive series of the metals). The speed of chemical
changes is discussed in greater detail under Chemical Equilibrium (see
also Sulphurous acid).
Exercises. — 1. What are the valences of the negative radicals
of phosphoric acid (p. 71), and of acetic acid (p. 93) ? What must be
the formulae of calcium phosphate, cuprio acetate, aluminium phos-
phate, ferrous carbonate, ferrous sulphate, cupric chloride ?
2. What is the valence of phosphorus in phosphoric anhydride
(p. 71) ? What must be the formulae of the chloride and the sul-
phide of phosphorus, and of aluminium oxide ?
3. What are the valences of the elements in the following : LiH,
NHS, SeH2, BN?
4. WThat are the valences of the metals and radicals in the following :
Pb(N"08)2, Ce(S04)2, KC1, KMn04 (potassium permanganate)? Name
all the substances in 3 and 4.
5. Write the formulae of ferrous and ferric oxides, of ferrous and
ferric nitrates, of stannous and stannic sulphides.
6. What must be the relative rates of diffusion of hydrogen and
of carbon dioxide ?
7. Make equations to represent (a) the reduction of lead dioxide
(Pb02) by hydrogen ; (b~) the actions of aluminium upon cold water
and upon steam at a red heat.
CHAPTER VIII
WATER
THE great quantity of water which, occurs in nature makes it one of
the most familiar chemical substances. The ocean covers about three-
fourths of the surface of the earth, and in most habitable regions lakes
and streams abound. Water is found also in the bodies of both ani-
mals and plants in large quantities, and is indeed essential to the
working of living organisms.
Natural Waters. — The water found in nature varies greatly in
the amount of foreign material which it contains. Sea- water holds
about 3.6 per cent of solid matter in solution, while rain-water is the
purest natural water. Even rain-water contains foreign matter, how-
ever. When we heat it, bubbles of gas form on the sides of the vessel,
showing that oxygen and nitrogen from the air have been dissolved by
the water as it fell. On evaporating a considerable mass of such
water, we find that, aside from dust, crystals of chemical substances,
such as ammonium nitrate, may be recognized in the residue. Of well
and surface waters, some which contain calcium sulphate, calcium-
bicarbonate, and compounds of magnesium in solution are described as
hard. Others contain compounds of iron, and still others are effer-
vescent and give off carbon dioxide. These are called mineral waters.
All of the dissolved substances are obtained by the water in its prog-
ress over or under the surface of the ground.
Water which is to be used for domestic purposes is examined, not
only to ascertain the amount of the ingredients which produce hard-
ness, but also with reference to the proportion of organic matter which
it may hold in solution. This usually gains access to the water by
admixture of sewage (p. 72). It is not the organic matter itself which
is deleterious, but the bacteria of putrefaction and disease which are
likely to accompany it. Inoculation of culture media with the water
can alone show whether or not the latter are present.
Purification of Water. — The foreign materials which water may
contain are divisible into two kinds, — dissolved matter and suspended
113
114 INORGANIC CHEMISTRY
matter. No water is free from either of these varieties of impurity.
In chemical laboratories distilled (p. 38) water of a more or less pure
kind is always employed, but this represents only a crude purification.
By using a platinum still and condenser, water of much greater purity
can be obtained. Yet, on account of the solvent power of water, it is
impossible to keep such a liquid even for a short time. Ordinary
glass dissolves in water to a very noticeable extent.
The purity of water is most easily investigated by measurement of
its resistance to the passage of electricity. A column only one milli-
meter long, of the purest distilled water that can be made, has a greater
resistance than a copper wire of the same cross-section and long enough
to reach a thousand thnes round the earth at the equator. During a
few minutes' exposure to the air, or contact with a glass vessel, how-
ever, a sufficient amount of foreign material (p. 95) of high conductiv-
ity is taken up to diminish its resistance very greatly.
For ordinary purposes the suspended matter which water contains
is removed by filtration (p. 11). In the laboratory this takes place
through unsized paper. The pores of the paper are sufficiently small
to retain particles of the dimensions usually met with, while permit-
ting the passage of the water with its dissolved matter. On a large
scale, beds of gravel are employed. In the household the Pasteur
filter is more compact and efficient. The water is forced by its own
pressure through the pores of a closed tube made of unglazed porce-
lain. Care must be taken to clean these tubes at frequent intervals,
so that organic and perhaps putrescent matters may not accumulate
upon them. If the cleansing is not carried out, the Pasteur filter-tube
becomes a breeding-place for bacteria,, and may add greatly to the con-
tamination of the water instead of diminishing it.
Matter in solution cannot be removed by filtration, and is eliminated
by distillation. Since the water is converted into steam and is con-
densed in platinum or tin pipes, only gases or volatile liquids dis-
solved in it can pass into the distillate.
Physical Properties of Water. — When we view a white object
through a deep layer of water we find that the liquid has a blue or
greenish-blue color. At a pressure of 760 mm., it exists as a liquid
between 0° and 100°. Below 0° it becomes solid, above 100° a gas. Of
all chemical substances it is the one which we use most, so that famil-
iarity with its properties is indispensable to the chemist. It will serve
also as a typical liquid, since it differs from others only in details.
WATER 115
The weight of a cubic centimeter of water at 4° gives us our unit, the gram. A
kilogram of water at 0° occupies 1.00013 liters, or 0.13 c.c. more than at 4° C. A
kilogram of ice at 0° occupies 1.09083 liters, or 90.7 c.c. more than an equal weight
of water. The volume of the same weight of water at 100° is 1.0432 liters.
Ice. — The raising or lowering of the temperature of a gram of
water through one degree corresponds to the addition or removal of
one calorie of heat. The conversion, however, of a gram of water at
0° to a gram of ice at 0° required the removal of 79 calories of heat.
The mere melting of a gram of ice causes an absorption of heat to the
same amount. This is called the heat of fusion of ice. At 0°a mix-
ture of ice and water will remain in unchanged proportions indefinitely.
Any cause which tends permanently to lower or raise the temperature
by a fraction of a degree, however, will bring about the disappearance
of the water or of the ice respectively. This temperature is called
Jbhe melting or the freezing point. Properties of this kind, marked
by transition points from one state to another, are much used in chem-
istry for keeping other bodies or systems at a constant temperature
during measurement or observation. A mixture of ice and water sur-
rounding a body, when kept in constant agitation, will automatically
maintain the body at a fixed temperature (0°) so long as both compo-
nents hold out.
Steam and Aqueous Tension. — At atmospheric pressure, water
passes into steam rapidly at 100°, but at lower temperatures, and even
when frozen, it does the same thing more slowly. The best way to
define the quantity of the vapor at various temperatures is by the
gaseous pressure it exercises. This is proportional to the concentra-
tion (p. 81) of the aqueous material in the space above the water,
and has a definite value for each temperature. Its amount may be
shown by allowing a few drops of water to ascend into the vacuum at
the top of a barometric column (Fig. 39). The tube on the left shows
the mercury when nothing presses on its surface. The tube on the
right shows the result of admitting the water. The tension of the
atmosphere being the same for both, the smaller height of mercury
which now suffices to counterbalance it shows that something, which
can be nothing but the water vapor, is pressing on the surface of the
mercury, and makes up the rest of the total stress needed. The differ-
ence in the height of the two columns gives the value of this pressure,
which we call the vapor pressure of the water. The jacket surround-
ing the tube on the right enables us, by adding ice or warm water, to
116
INOKGANIC CHEMISTRY
keep the water that is admitted to the vacuum, and the parts of the
apparatus immediately in contact with it, at any temperature be-
tween 0° and 100°.
When ice is used outside, and a piece of it is introduced into the
vacuum, the vapor it gives off quickly reaches a pressure of 4.5 mm.
The vapor pressure of the ice takes the place of 4.5 mm. of mercury in
balancing the atmospheric pressure, and
so the mercury column falls by this
amount. Similarly, water at 10° causes
a fall of 9.1 mm. and at 20° of 17.4mm.,
so that these represent the mercury-
height values of the vapor pressure at
these temperatures. The quantity of
water used makes no difference, so long
as a little more is present than is re-
quired to fill the available space with
vapor. Of course, if a large amount
is admitted, its dead weight will take
the place of an equal weight of mer-
cury in balancing the pressure of the
air. If there is a measurable column
of water, its height must be divided
by 13.6 (the sp. gr. of mercury), and
counted as if it were part of the mercury.
An expression is needed to describe
the varying power of the water at
different temperatxires to maintain dif-
ferent pressures of its vapor. This we
call the aqueous tension of the liquid.
Its magnitude at any given temperature
is measured by the maximum pressure
reached by the vapor (see Kinetic-molecular hypothesis applied to
liquids).
With water at higher temperatures the fall of the mercury column
becomes much greater. At 50° it is 92 mm., at 70° it is 233.3 mm., at
90° it is 525.5 mm., and at 100° it is 760 mm. At the boiling-point,
therefore, the aqueous tension takes the place of the whole barometric
column, and is equal to the average air pressure. At 121° the aqueous
tension is two atmospheres, at 180° it is ten atmospheres.
There is another standpoint from which these phenomena may be
PIG. 39.
WATER 117
viewed. Water vapor can exist at 10° only if the pressure upon it is
9.1 mm. or less. If we imagine the water placed in a cylinder closed
by a frictionless, weightless piston (Fig. 40), then at 10° the piston
will remain at rest whether we place it high or low, provided it is
loaded with a weight exactly equal to that of a layer of mercury 9.1 mm.
thick covering its whole area. We speak of such a system as being in
equilibrium.* With a less weight the piston will move slowly upwards,
as the vapor continually given off by the water presses upon it, until
it reaches the top or the water all evaporates. Conversely,
if it bears a greater load, it will move down and the vapor
will condense on the walls and bottom of the cylinder until
the piston comes in contact with the water itself and the
vapor is all abolished. These conceptions will find con-
stant application not only to physical but also to chemical
phenomena (see Kinetic-molecular hypothesis). The ex- I L_
pression :
Aq. (liq.) <=> Aq. (vap.)
is used to represent the state of equilibrium in a system like FIG. 40.
the above.
One other phase of this subject is recognized by special phraseology.
When water at a certain temperature has given the full amount of
water vapor to the space above it that its aqueous tension permits, we
say that the space is saturated with vapor. That concentration of
vapor which constitutes saturation varies with the temperature of the
water and depends therefore solely on the power of the water to give
off vapor. It has nothing to do with the size of the space, and is even
independent of other gases the space may already contain (p. 88).
The space immediately above the surface of the ground, which is
mainly occupied by atmospheric air, is, on an average, less than two-
thirds saturated with water vapor. That is to say, such air, when
inclosed in a vessel containing water, will take up about one-half more
than it already contains. At 100° the water vapor displaces the air
entirely, and the liquid is said to boil.
The water present in the air plays an important part in many
chemical phenomena, as we shall see. All our substances and appara-
tus have traces of water condensed on their surfaces. This water is,
in a sense, in an abnormal condition, for it does not evaporate even in
* In chemistry we do not speak of stable, unstable, and indifferent equilib-
rium, as is done in physics, but only of the first. The term " inetastable, " how-
ever, is employed (see Chap. x).
118 INORGANIC CHEMISTRY
dry air. It is observed to pass off in vapor, however, when we have
occasion to heat any material.
In passing into vapor, water absorbs heat without changing its
temperature. A gram of water at 100°, for example, in turning into a
gram of steam at 100°, takes up 537 calories. This is called its heat
of vaporization. Steam, in fact, contains much more internal energy
than an equal weight of water at the same temperature, just as water,
in turn, contains more energy than ice.
The temperature of 100° is, like the melting-point of ice, an im-
portant transition point, and has the same properties, mutatis mutan-
dis, as the latter. It is less exactly recoverable by simply keeping a
vessel full of water in ebullition, however, because variations in the
pressure of the atmosphere affect it more markedly than they do the
melting-point of ice. Near to 100°, the boiling-point rises or falls
about 0.037° for 1 mm. change in pressure (cf. p. 116). On the top
of Mont Blanc water boils at 84°.
Water as a Solvent. — One of those physical properties of water
which are most used in chemical work is its tendency to dissolve many
substances. This subject is so important and extensive that we shall
presently devote a complete chapter to some of its simpler and more
familiar aspects.
Chemical Properties of Water. — Water is so very frequently
used in chemical experiments in which it is a mere mechanical ad-
junct, that the beginner has difficulty in distinguishing the cases in
which it has itself taken part in the chemical interaction. The rather
limited list of kinds of chemical activity it can show should therefore
receive careful notice : It (1) is a relatively stable substance. It (2)
combines directly with many substances. In the commoner of the two
kinds of this action, the formation of hydrates (see below), the com-
pound exists in the solid form only, however, and is decomposed in
solution. Finally, it (3) interacts with some substances in a way which
we describe as hydrolysis. We shall not discuss this last property
until some substance which is markedly affected is encountered (see
Preparation of hydrogen chloride).
Perhaps we should add that steam at a high temperature oxidizes
elements which readily combine with oxygen. For example, it turns
iron into the magnetic oxide (p. 99). At such high temperatures,
however, the water is partially resolved into a mixture of hydrogen and
WATER 119
oxygen, and, the latter being the more active of the two elements, the
oxidizing effects predominate. Even other compounds containing
oxygen will give exactly the same results. Hence this cannot be re-
garded as a property of water itself.
Water a Stable Compound. — Whether the substance is relatively
stable or unstable is, in the case of a compound, the first chemical
property to be given. Usually the specification is in terms of the
temperature required to decompose it. Thus, potassium chlorate gives
off oxygen at a low red heat. Water, even at 2500°, is but little decom-
posed, and reunion occurs as soon as the temperature is lowered.
Union of Water with Oxides. — When sodium combines with
oxygen under certain conditions we obtain sodium oxide (Na20). The
product unites violently with water to form sodium hydroxide :
Na/) + H20 ~* 2NaOH.
The slaking of quicklime is a more familiar action of the same kind :
CaO + H20 -> Ca(OH)2.
No other products are formed. The clouds of steam produced in the
second instance are due to evaporation of a part of the water by the
heat produced in the formation of calcium hydroxide. The aqueous
solutions of these two products have a soapy feeling, and turn red lit-
mus blue, and the substances therefore belong to the class of alkalies
(p. 98) or bases. Very many hydroxides which are of the same na-
ture, for example ferric hydroxide (Fe(OH)8) and tin hydroxide
(Sn(OH)2), are formed so slowly by direct union of the oxide and water
that they are always prepared in other ways.
Some oxides, although they unite with water, give products of an
entirely different character. Phosphoric anhydride and sulphur dioxide
(p. 71) are of this class and yield acids.
These two classes of final products are so different that we make
the distinction the basis of classification of the elements present in the
original oxides. The elements, like sodium and iron, whose oxides
give bases, are called metals ; those, like phosphorus, whose oxides give
acids, are called non-metals. The distinguishing words are selected
because the division corresponds, in a general way at least, with the
separation into two sets to which merely physical examination of the
elementary substances would lead.
120
INORGANIC CHEMISTRY
Formerly the hydroxides of metals were termed " hydrates," and the
word is still used familiarly by chemists in a few cases, such as potas-
sium « hydrate " (KOH) and sodium « hydrate " (NaOH). These sub-
stances, however, have nothing in common with the compounds properly
known as hydrates whose nature is discussed in the next section.
FIG. 41.
Hydrates. — Many substances when dissolved in water and re-
covered by spontaneous evaporation of the solvent are found to have
entered into combination with the liquid. The products, which are
solids, are called hydrates. These compounds show definite chemical
composition expressible in terms of chemical unit weights of the con-
stituents. Often much heat is given out in their
formation. Thus, in the case of washing soda,
the decahydrate of sodium carbonate (Na^COg,
lOHgO), the heat of the union (p. 76) is 8800 cal.
The hydrates have physical properties entirely
different from those of their components. Thus,
cupric sulphate, often called anhydrous cupric
sulphate to distinguish it from the compound
with water, is a white substance crystallizing
in shining, colorless, needle-like prisms. The
pentahydrate (blue-stone or blue vitriol) is blue in color, and forms
larger but much less symmetrical (asymmetric or triclinic) crystals
(Fig. 41):
CuS04 + 5H20-> CuS04, 5H20.
When heated, the hydrates, as a rule, lose none of the constituents of
the original compound, but only those of the water. They do so, in
general, rather easily. Hence, to avoid the disguise of the fundamental
substance which would occur if we wrote, for example, H10CuS09, we
segregate (as above) the elements of the water in the formula. The
aqueous solutions made from the anhydrous substances and from the hy-
drates have identical physical and chemical properties. Hence the
cheaper of the two forms is generally purchased, and many of the
chemicals used in laboratories are in the form of hydrates.
Some of these hydrates decompose very readily. The decahydrate
of sodium sulphate, Na2S04, 10H20 (Glauber's salt), gives up all the water
it contains when simply kept in an open vessel. At 100° blue vitriol
very quickly loses 4H20 and the rest of the water more slowly. The
last equation might therefore have been written to show a reversible
WATER 121
action (p. 64). A decomposition which proceeds at high temperatures,
while at lower temperatures recombination of the constituents can take
place, as with these hydrates, is called a dissociation. The decomposi-
tion of potassium chlorate (p. 64) is not a dissociation because it is not
reversible ; oxygen will not under any circumstances reunite with potas-
sium chloride.
The condition which controls such actions is not, however, the tem-
perature alone. When Glauber's salt is kept in a closed bottle, a very
little of it loses water, and then the decomposition ceases. When the
bottle is left open, the dissociation proceeds until no decahydrate re-
mains. The cause of this we discover when a crystal of the hydrate
is placed above mercury, like the ice or water in Fig. 39 (p. 116). It
shows a definite aqueous tension. At 9° the value of this is 5.5 mm.
As its temperature is raised, the tension increases. When the tem-
perature is lowered, on the other hand, the tension diminishes, the
mercury rises, and more of the water enters into combination again.
Different hydrates show different aqueous tensions at the same tempera-
ture. For example, at 30°, water itself is 31.5 mm., strontium chloride
(SrCl2, 6H20) 11.5 mm., cupric sulphate (CuS04, 5H20) 12.5 mm., barium
chloride (BaCl,, 2H20) 4 mm.
It appears, therefore, that the water in these compounds evaporates
as do ordinary waters. Those which, like washing soda, have a vapor
tension approaching that of water itself, lose their water at ordinary
temperatures at a rapid pace. In this connection we have to remember
that atmospheric air is always less than two-thirds saturated with water
vapor, and the partial pressure of this vapor opposes the dissociation.
Thus at 9°, the vapor tension of water being 8.6 mm., the average vapor
pressure of water in the atmosphere will be about 5 mm. Any hydrate
with a greater aqueous tension than 5 mm., at 9°, such as Glauber's
salt, will therefore decompose spontaneously in an open vessel. But
those with a lower vapor tension, such as the peiitahydrate of cupric
sulphate with a tension of 2 mm. at 9°, will not do so (see p. 135).
_The behavior of hydrates does not indicate, as might seem at first
sight to be the case, that the water is contained in them in some way
in the free state. The fact is that the above statements, with corre-
sponding changes in the wording, might be made of all dissociations in
chemistry. Oxides give a different pressure of oxygen at each tem-
perature, carbonates of carbon dioxide, and so forth.
The measurement of the vapor tension of hydrates gives definite information
in regard to whether there are other hydrates, say of cupric sulphate, with less
122 INOKGANIC CHEMISTRY
than the normal number of formula-weights of water. If there were only two
substances, CuS04 and CuS04, 5H20, with no compound of intermediate composi-
tion, then a partially decomposed specimen would be made up partly of the one
substance and partly of the other. But if there were an intermediate compound,
say CuS04, 3H2O, then desiccating a specimen of the pentahydrate would give
nothing but mixtures of CuSO4, 3H20 and CuSO4, 5H20 until all the latter was
decomposed. Then, and only then, the trihydrate would begin to lose water. Now
the trihydrate, being a definite and different substance, would have a vapor tension
of its own, and experimental study would show its presence.
Experiment shows that there really are several hydrated cupric sulphates.
The pentahydrate, at 50°, has a vapor tension of 47 mm., and this vapor tension is
observed so long as any pentahydrate remains to be decomposed. As soon as the
proportion of water goes down to CuSO4, 3H20, the vapor tension suddenly drops
to 30 mm. As the desiccation continues, this tension is maintained until the com-
position has reached CuS04, H2O. At this point the vapor pressure falls to that
of the rnonohydrate, 4.6 mm., and remains at this value until all the rest of the
water has been removed. Had there been no intermediate compound with 3H2O
the tension would have dropped at once from 47 mm. to 4.5 mm. If, conversely, we
try to combine water as vapor with anhydrous cupric sulphate, at 50°, a vapor
pressure of at least 4.5 mm. is required to cause union to take place. The union
stops when one formula-weight of water has undergone combination. To intro-
duce more, the concentration of the water vapor must be increased to nearly seven
times its first value, namely, to 30 mm. pressure. This enforces combination up to
CuS04, 3H?O. For further hydration, a still higher pressure of water vapor is
needed (47 mm.), and the absorption ceases when CuS04, 5H26 has been formed.
There are thus three distinct reversible actions which succeed one another as
the hydration proceeds :
CuS04 + H2O T=> CuSO4, H2O
CuS04, H20 + 2H20 <=t CuSO4, 3H2O
CuS04, 3H2O + 2H2O <=> CuSO4, 6H2O.
The first represents a greater affinity than the second, and the second than the
third.
The graphic representation of these facts (Fig. 42) will make the behavior of the
compounds clearer. The proportion of water combined with one formula-weight
of cupric sulphate is laid off along the horizontal axis. The pressures at which it
enters or leaves the compounds at 50° are the ordinates. As far as 1H20 the pressure
is constant (4.5 mm.). Beyond that point and up to 3H2O it is constant but much
higher. Between 3H20 and 5H2O it is constant again but higher still.
The tension of free water at the same temperature is 92 mm. It is constant
irrespective of the amount of water, and would therefore be on a single continuous
line parallel to the horizontal axis and twice as high above it as the uppermost one
in the diagram. If, at 50°, a vessel of water were put under a bell jar alongside of
anhydrous cupric sulphate, its vapor would be more than sufficiently concentrated
fully to hydrate the compound. Again, while 4.5 mm. pressure of water vapor
will cause water to combine with anhydrous cupric sulphate at 50°, a pressure of
92 mm. will be required to liquefy the water vapor at the same temperature.
The above discussion shows that the last formula- weight of water, in hydrates
WATER
128
50
47—
40
30
20
0,10
which dissociate by stages, is not different in kind from the others. It differs only
in the degree of tenacity with which it is held. It is therefore unnecessary,
merely on this account, to dignify it by the separate name of water of constitu-
tion, as has been done by some chemists.
Water of hydration is frequently called •water of crystallization, on account
of the fact that when water is driven off by heating, the substance usually crumbles
to pieces (effloresces) . The term is not particularly appropriate for several reasons.
It suggests that water and crys-
tallization are related in some
way, which is not the case.
Sulphur, galena, potassium
chlorate, and thousands of other
crystallized substances, do not
contain the elements of water.
Nor do the substances which
combine with water remain
amorphous in its absence.
They all crystallize from the
molten condition or from some
non-aqueous solvent, although,
as substances different from the
hydrates, their crystalline form
is different. Iceland spar, or
any other crystallized carbonate
which can be decomposed by
heating, becomes opaque and
porous or 'falls to powder when the carbon dioxide is driven out. But it has
not occurred to any one to call this carbon dioxide of crystallization! The fact is
that all pure chemical substances, in solid form, when in a stable physical condi-
tion, are crystalline. Amorphous substances are always supercooled liquids.
The term arose from a misconception, and, when used, always succeeds in
transmitting the misconception along with the name. The ease with which some
of the hydrates decomposed suggested the idea that they contained water as a dis-
crete substance. There is no more justification for this idea, however, than for the
notion that carbonates contain ready-made carbon dioxide. The hydrates contain
the elements of water just as sugar and alcohol do, and there is no evidence that
they "contain water" in any other sense than that in which the phrase might be
used of these organic bodies.
In consequence of their decomposition into and formation from substances
capable of separate existence, the hydrates are classed with molecular compounds
(q.v. ). The behavior of the compounds of salts with ammonia (like 2AgCl, 3NH3),
with nitric oxide, and with each other (double salts), is quite similar.
4.5
0
1234
Formula-weights of water
FIG. 42.
Composition of Water. — The measurement of the proportions
by weight and volume in which hydrogen and oxygen combine to form
water has been the subject of a larger number of elaborate investiga-
tions than any other single problem of this kind. The difficulty in
124
INORGANIC CHEMISTRY
making the former measurement arises from the fact that both constit-
uents are gases, and therefore difficult to weigh.
The determination of this proportion which, until very recently, held its place
in all chemical works, was that made by the French chemist, Dumas. His experi-
ments gave the ratio of hydrogen to oxygen 2 : 15.96.
It was not until 1887 that Reiser obtained a figure for the oxygen appreciably
smaller than this, and soon determinations by other observers showed that Dumas'
proportion was probably too large. The investigation which
finally settled this question was that of Edward Morley. The
most striking of his experiments consisted in a series of syn-
theses of water, in which- he weighed the hydrogen as well as
the oxygen, and afterwards weighed the water produced from
them. The hydrogen was confined by absorption in palladium
(p. 107), and could thus be contained in large quantity in a
small, elongated bulb. During the progress of the experiment
it was driven out by a suitable heating arrangement. The
oxygen was contained in large globes holding 15-20 liters. The
losses in weight of the palladium tube and of the globes gave
the hydrogen and oxygen consumed. The manipulator in
which the gases were combined and the water collected is
represented in Fig. 43. The gases entered through two small
tubes marked A. Just above them, between two platinum
wires, a discharge of electricity started the union and when
necessary maintained it. The vessel was first filled by admit-
ting oxygen, and the hydrogen was burned at the mouth of
the tube from which it issued. This part of the apparatus
was immersed in a vessel of water with transparent walls
through which the union could be watched, and the steam
formed was condensed and collected in the bottom of the ves-
sel. The vacuum thus produced enabled the oxygen continually
to flow into the manipulator from the globes. In this way
forty-two liters of hydrogen and twenty-one liters of oxygen
could be combined in about an hour and a half.
At the end of the experiment this part of the apparatus
was disconnected and placed in a freezing mixture which con-
verted the water into ice and practically condensed the whole
of its vapor. The uncombined gas in the apparatus was withdrawn and its
nature and quantity determined. The increase in weight of the manipulator
gave the quantity of water formed. The success of each experiment could be
tested by comparing the sum of the weights of oxygen and hydrogen with that
of the water obtained from them. The manipulation was so skilful, and the
various corrections used were so adequate, that this difference was almost
negligible. The ratio of hydrogen to oxygen in water in this series of experi-
ments was 2 : 15.879, a result which agreed with the other methods of determin-
ing the same ratio which Morley used. It agrees also exactly with the average
of the numbers obtained by several other observers.
The most probable value of the ratio by weight, taking his own
PIG. 43.
WATER
125
and other trustworthy measurements into account, is given by Morley
as 2 : 15.879 or 2.015 : 16. The proportion by volume is 2.0027
volumes of hydrogen to 1 volume of oxygen.
That the proportion by volume is very close to 2 : 1 may easily be
shown. We may use a U-shaped tube closed at one end by a stopcock
and graduated (Fig. 44). At first, the left limb of the tube, called a
eudiometer, is filled with mercury. One of the gases is admitted so as to
fill a portion of the tube and, the levels having been
equalized (cf. p. 82), the volume of the gas is read.
Then some of the other gas is introduced and the
leveling and reading repeated. Let us suppose
that 15 c.c. of hydrogen and 10 c.c. of oxygen
have thus been taken. The right limb is then
filled with mercury and closed firmly with the
thumb. A spark from an induction coil passing
between the two short platinum wires near the top
of the tube explodes the mixture. The steam pro-
duced by the union condenses almost immediately
and occupies practically no volume worth con-
sidering. When the thumb is removed, the mer-
cury rises on the left and fills up the space left
by the disappearance of part of the gases. Unless
the proportion taken happens to have been exact,
some of one or other of the gases will remain.
Its volume is measured by equalizing the levels
and reading as before. In the case we have im-
agined, the residual gas is oxygen, and there are
almost exactly 2.5 c.c. of it. It is evident, there-
fore, that 15 c.c. of hydrogen united with 7.5 c.c. of oxygen ; in other
words, the proportion by volume is 2 : 1.
Gay-Lussac's Law of Combining Volumes. — The almost
mathematical exactness with which small integers express this propor-
tion is not a mere coincidence. Whenever gases unite, or gaseous
products are formed, the proportion by volume (measured at the
same temperature and pressure) of all the gaseous bodies concerned
can be represented very accurately by ratios of small integers. This
is called Gay-Lussac's law of combining volumes (1808). Thus, when
the above experiment is carried out at 100°, in order that the product,
watei'j may be gaseous also, it is found that the three volumes of the
FIG. 44.
126
INORGANIC CHEMISTRY
T '
• ' *•
constituents give almost exactly two volumes of steam. For example,
15 c.c. of hydrogen and 7.5 c.c of oxygen give 15 c.c. of steam.
Of course the hydrogen, oxygen, and steam must be measured at
the same pressure, and the temperature
must remain constant (100°) during the
experiment. Proper manipulation secures
the former, and a jacket filled with steam
(Fig. 45) the latter condition. Strips of
paper, 1, 2, and 3, are pasted on the jacket
in such a way that equal lengths of the
eudiometer, in this case a straight one,
are laid off. The three divisions having
been filled with a mixture of hydrogen
and oxygen in the proper proportions, the
gas, after the explosion, shrinks so as to
occupy, at the same pressure, only two
of them.
From this universal truth in regard
to the combination of gases, we draw the
important inference that the chemical
unit-weights of simple substances, and the
formula- weights of compounds, in the
gaseous condition, occupy at the same
temperature and pressure volumes •which
are equal or stand to one another in the
ratio of small integers (see Molar weights).
FIG. 45.
The chemical behavior of the other compound
of hydrogen and oxygen, hydrogen peroxide
(<?.».), is difficult to comprehend until further experience has been gained.
Then, too, its formula (H2O2) cannot be justified until the means of deter-
mining molar weights in solution have been discussed. In view of these facts,
and because it furnishes the simplest illustration of the meaning of a structural
formula, a conception which would be quite out of place at this stage, it will be
taken up later.
Exercises. — 1. Name some other transitions from one physical
state to another which are familiar (p. 115).
2. What evidence is there in the common behavior of ether,
alcohol, and chloroform tending to show that these liquids have high
vapor tensions ?
WATER 127
3. If the pressure of the steam in a boiler is ten atmospheres, at what
temperature is the water boiling (p. 116) ?
4. How many grams of water could be heated from 20° to 100° by
the heat required to melt 1 kgm. of ice at 0° ?
5. What do you infer from the fact that alum and washing soda
lose their water of crystallization when left in open vessels, while
gypsum does not (p. 121) ?
6. Which facts show conclusively that hydrates are true chemical
compounds? Is there any fact which throws doubt on this con-
clusion ?
7. In what way does a hydrate differ from (a) a solution, (£) an
hydroxide ?
8. Should you expect to find any difference, in respect to chemical
activity, between the three forms of water (ice, water, and steam) ?
If so, arrange them in the order of probable increasing activity (pp.
26-28). Have we had any experimental confirmation, or the reverse,
of this conclusion ?
9. Which contains more chemical energy, and is therefore more
active, the anhydrous substance, or the corresponding hydrate ?
THE KINETIC-MOLECULAR HYPOTHESIS
As soon as we have constructed a law (p. 7) we desire immediately
to find out the basis of the constant mode of behavior it epitomizes.
If no explanation, that is, more detailed description, is forthcoming
as the result of closer observation, we proceed to imagine one (p. 10).
This always takes a mechanical form, often crude at first, and later
undergoing refinement. Thus, at first, the phenomena of light were
explained by the conception of clouds of fine corpuscles emanating
from the luminous body. The chances of hitting upon an objective
reality by guess-work like this is obviously remote. Whether such
particles did really fly about was not the main question, however.
Their value lay in the fact that they could be pictured concretely and
gave a basis for further thought and perhaps suggestions for new
experiments. Such a structure of the imagination is called an hypoth-
esis. If it furnishes an explanation of more than one law, so much
the better.
The Molecular Hypothesis. — The only mechanical basis we can
imagine to account for the physical properties of matter is a discontin-
uous structure of some description. The fact that all kinds of matter
can be compressed (gases to an enormous extent, solids and liquids to a
measurable extent) may be explained, either by a diminution in the
volume of the material itself, or by the closer packing together of the par-
ticles into which this material is divided. It is evident that the latter
is much more in harmony with our experience. Compression, we
imagine, therefore, does not diminish the actual volume occupied by
matter, but crowds the particles closer together and diminishes the
space between them. The same hypothesis will furnish a concrete
description of how one kind of matter will frequently absorb a large
quantity of another. Thus we picture the hydrogen gas taken up by
iron and other metals as being packed away in the spaces between the
particles of the metal. So also, in solution, the volume of the liquid
does not usually increase by an amount equal to that of the substances
128
THE KINETIC-MOLECULAR HYPOTHESIS 129
dissolved. Hence we imagine the particles as possibly incompressible
and the interstices between them as furnishing a part of the accommo-
dation for the foreign material. A particle is a fragment of matter
which can be seen and measured directly, and handled separately.
Since the particles of this hypothesis have none of these qualities, we
distinguish them by the name molecules. Molecules are the imaginary
units of which bodies are aggregates.
Since the behavior of gases has been considered most fully in the
preceding chapters, and is in any case capable of more exact and
simple description than that of solids or liquids, we shall apply this
hypothesis first to them.
Kinetic- Molecular Hypothesis Applied to Gases- — Let us
first build up our hypothesis to fit the qualitative properties of gases.
The most remarkable thing about a gas, considering the looseness with
which its material is packed, is the total absence in it of any tendency
to settling or subsidence. Since the molecules cannot be at rest upon
one another, as the great compressibility shows, we are driven to
suppose that they are •widely separated from one another, and that
they occupy the space by constantly moving about in all directions.
But a moving aggregate of particles which does not even finally settle
must be in perpetual motion. We must, therefore, imagine the mole-
cules to be wholly unlike particles of matter in having perfect elasticity,
in consequence of which they undergo no loss of energy' after a
collision. They must continually strike the walls of the vessel and
one another and rebound, yet without loss of motion. The diffusibility
of gases and their mutual permeability require no additional assump-
tions. The fact that each gas is homogeneous, efforts to sift out lighter
or heavier samples having failed, requires the supposition that all the
molecules of a pure gas are closely alike.
Passing now to Boyle's law (p. 84), the thing to be accounted for
is that when a sample of a gas diminishes in volume, its pressure
increases in the same proportion. Let the diagram (Fig. 46) represent
a cylinder with a movable piston, upon which weights may be placed
to resist the pressure. Now the pressure exercised by the gas cannot
be like the pressure of the hand upon a table, since we have just
assumed that the particles are not even approximately at rest, and the
spaces between them are enormous compared with the size of the
molecules themselves. The gaseous pressure must therefore be attri-
buted to the colossal hailstorm which their innumerable impacts upon
130 INORGANIC CHEMISTRY
the piston produce. If this is the case, the compressing of a gas must
consist simply in moving the partition downwards so that the particles
as they fly about are gradually restricted to a smaller and smaller
space. Their paths become on an average shorter and shorter. Their
impacts upon the wall become more and more frequent. So the
pressure which this occasions becomes greater and greater, and is pro-
portional to the degree of crowding of the molecules.
There are two other points which must be added. When
we diminish the volume to one-half, we find from experience
that the pressure becomes exactly, or almost exactly, twice
as great. This must mean that although the particles are
T becoming crowded they do not interfere with one another's
motion, excepting of course where actual collision causes
a rebound. Only in the absence of interference would
doubling the number of particles per unit of volume give
exactly double the number of impacts on the walls. Hence
FIG. 46. the particles must have practically no tendency to cohesion.
Again, the molecules must move in straight lines, because,
if they moved in orbits of some kind, many of the orbits would not be
intersected by the wall of the vessel until great reduction in the
volume had taken place, and thus, as the volume diminished, the fre-
quency of the impacts, and therefore the pressure, would increase
faster than the concentration.
Boyle's law therefore adds four more conceptions to our molecular
hypothesis, namely, that the impacts of the particles produce the press-
ure, that the crowding of the molecules represents the concentration
(p. 80), and that the particles move in straight lines and show almost
no cohesion, since pressure and concentration are very closely propor-
tional to one another.
It will be seen, on consideration, that if the molecules are assumed to repel
one another, they would do so more violently the more closely they were packed
together. This assumption would therefore suit the case of a gaseous body in
which the pressures increased according to some power of the concentration 'other
than the first, and therefore much more rapidly than in known gases. In spite of
its inapplicability, this notion is supposed by many people to be part of the
kinetic hypothesis.
Charles' law (p. 87), that a gas receives equal increments in volume
or pressure for equal elevations in temperature, reqxiires but one addi-
tion to the hypothesis. Concretely, if our specimen of gas (Fig. 46) is
at 0°, and we permit its pressure to remain constant by leaving the
THE KINETIC-MOLECULAR HYPOTHESIS 131
same weight on the piston, then when the temperature of the gas is
raised to 1°, the volume will gain ^\^ of the original volume. If, on
the other hand, we restrict the gas to the original volume, the pressure
will evidently increase, and the augmentation will be 5|^ of the
original pressure. Now, how can we account for an increase in
pressure as the result of heating a mass of rapidly moving molecules ?
The action of a particle colliding with a surface is measured in physics
in terms of its mass and its velocity. It is evident that heating a cloud
of molecules would not increase the mass of each, and it must there-
fore increase the velocity of each since the kinetic energy of all
becomes greater. This conclusion is in harmony with our experience
that violently rubbing a solid raises its temperature, and such a mode
of treatment might plausibly be supposed to communicate motion to
the minute parts of the body.
The fact that the combining volumes of gaseous substances are equal,
or stand to one another in the ratio of small whole numbers (cf. Gay-
Lussac's law, pp. 125-126), suggests two ideas : First, that chemical
combination, considered in detail, and arranged to harmonize with this
hypothesis, would involve unions of a few particles of more than one
kind to form composite molecules.* And, second, that a simple inte-
gral relation must be assumed to exist between the numbers of mole-
cules in equal volumes of different gases, at the same temperature and
pressure. Avogadro (1811), the professor of physics in Turin, put
forward the hypothesis that these numbers might be equal. A more
strict study of the assumptions we have been making, and of some addi-
tional facts, has since shown that no other conjecture than Avogadro's
would be consistent with them. Thus it now bears the relation of a
logical deduction from the kinetic-molecular hypothesis and the
properties of gases, and is known as Avogadro's hypothesis. It may
also be put in the form : At the same temperature and pressure, the
molecular concentration (cf. p. 80) of all kinds of gases has the same
value. _
The law of diffusion (p. 108) harmonizes with the kinetic-molecular
hypothesis without further modification of the latter. The strict
deduction of this law, as well as of the preceding ones, from our series
of assumptions, will be found in any work on physical chemistry.
Finally, we have referred (p. 85) to the fact that at low pressures
* This is essentially the idea used by Dalton, before Gay-Lussac's law was
known, however, for the explanation of the laws of chemical combination. He
called it the atomic hypothesis (q.v.).
132
INORGANIC CHEMISTRY
the concentration increases more, and at high pressures much less than
Boyle's law indicates. The former effect is brought into accord with
our hypothesis when we remember that the matter even of gases can
cohere, as is shown plainly when they are solidified. The tendency of
the molecules to cohere must therefore show itself in the gaseous con-
dition by pulling the gas together and producing somewhat greater
concentration than is strictly consistent with the value of the pressure.
Althougn this effect of cohesion is usually insignificant, the modern
method of liquefying gases (q.v.) depends upon it almost entirely.
The abnormally small reductions in volume which occur when the
volume of the gas has already been greatly reduced remind us that, ac-
cording to our hypothesis, it is only the space between the molecules
that is diminished as pressure rises, and not the space occupied by the
molecules. Hence, when the molecules have become so crowded to-
gether that this irreducible space begins to form an appreciable frac-
tion of the whole, a doubling of the pressure will diminish to one-half
its value only a part (the vacant part) of the volume the gas occupies.
If the incompressible space occupied by the molecules is called 6, and that of the
whole gas «, then the amended form of Boyle's law reads p (u— 6) = constant. Sim-
ilarly, if the cohesive tendency is taken into account, it is plain that its effect will
be numerically greater at small volumes, although not so easily observed. It is in
fact inversely proportional to the square of the volume. If it is expressed in the
same units as the pressure by a, the total of the compressing tendencies becomes
p -\ — | • Hence Boyle's law, for constant temperatures, as amended by Van der
Waals, reads I p -\ — - 1 (» — 6) = constant, a formula which describes the actual
behavior of most gases with remarkable accuracy. Hydrogen alone, at ordinary
temperatures, shows uo excessive compressibility at low pressures. The cohesion
(a) is so slight that the effect of the constant (6) counterbalances it from the very
first.
We may summarize the facts about gases, appearing in italics
above, with the corresponding fictions, in heavy type, which we have
added one by one m manufacturing our hypothesis, as follows :
FACTS.
HYPOTHESIS.
Non-settling "]
f The fictitious particles called molecules are,
Compressibility !
J at 0° and 760 mm. , at great average distances
Diff usibility [
| from one another ; they are in constant
Permeability J
[_ motion and have perfect elasticity.
Homogeneity
The molecules of the same substance are
closely alike.
THE KINETIC-MOLECULAR HYPOTHESIS
133
FACTS.
HYPOTHESIS.
Relation of pressure and concen-
tration (Boyle's law).
Relation of volume (or pressure)
and temperature (Charles' law).
Relation of atomic weights and
volumes (Gay-Lussac's law) .
Law of diffusion.
Abnormal incompressibility, es-
pecially at high pressures.
Abnormal compressibility, es-
pecially at low pressures.
The effect of pressure is produced by the im-
pacts of the molecules, and is proportional to
the degree to which they are crowded to-
gether ; the molecules move in straight lines,
and have almost no tendency to cohesion.
A rise in temperature increases the velocity
and therefore tha kinetic energy of the mole-
cules.
Chemical union consists in fusion of different
kinds of molecules (Dalton's hypothesis), of
which there are equal numbers in equal vol-
umes of different gases at the same tempera-
ture and pressure (Avogadro's hypothesis).
The molecules themselves are incompressible.
The tendency to cohesion becomes evident
under some circumstances.
Critical Phenomena. — We may now use the terms of the kin-
etic-molecular hypothesis in describing a property of gases not yet dis-
cussed. When the concentration of a gas at ordinary temperatures is
greatly increased by compression, the cohesive forces have an oppor-
tunity to produce liquefaction. In many cases, as with sulphur diox-
ide and carbon dioxide, when the approximation of the molecules has
reached a certain point, the liquid begins to form on the sides of the
vessel. The condition is then exactly the same as that of aqueous
vapor and water (p. 117), and no further increase in pressure is required
to complete the liquefaction of the whole. The only difference be-
tween steam, at a pressure below the aqueous tension of water at 10°,
and carbon dioxide at the same temperature, is that4tot more than 9.1
mm. of pressure is required to liquefy the steam, while about 50 atmos-
pheres are needed to liquefy the carbon dioxide.
There are some gases in which, at the ordinary temperature, even
with the closest approximation of the molecules, the cohesion is unable
to overcome the motion of the molecules and draw the material together
into the more compact liquid form. Such gases are hydrogen, oxygen,
nitrogen, and air, which is a mixture of the last two. The remedy is
obvious. We know of no way to increase the intrinsic cohesiveness of
134 INORGANIC CHEMISTRY
the material, but we can reduce the kinetic energy of the molecules by
lowering the temperature of the gas.* When this has been done suffi-
ciently, compression is followed by liquefaction. Now it is found that
there is a critical value for each individual gas to or beyond which the
kinetic energy must be reduced by lowering the temperature, before
the cohesive tendency of that particular gas can become effective to pro-
duce liquefaction. The highest temperature below which liquefaction
is possible is called the critical temperature. For oxygen this temper-
ature is —118°, for hydrogen about —234°, for nitrogen —146°. For
carbon dioxide it is 31.35°, for sulphur dioxide 156°, for water 358°.
The temperature of a room being below the critical point of the last
three substances, they are all liquefiable without cooling, and more
easily the farther the ordinary (say 20°) lies below the critical tem-
perature.
The foregoing is an example of how the images furnished by an
hypothesis assist us in understanding facts, and how the figurative
language it suggests enables us to explain the relations of the facts
briefly and vividly. In interpreting statements like the above, how-
ever, we must carefully distinguish between the facts they contain,
which are permanent, and the fictions and figures of speech in which
we have clothed them. The wrappings will doubtless go out of fashion
and change in the future, as they have done in the past, in response to
our preference in emphasizing some different inter-relation in the
growing body of truth. The kinetic hypothesis is already considered
out of fashion in some quarters.
Kinetic Hypothesis Applied to Liquids. — The phenomena con-
nected with surface tension, such as coherence into drops, show that
cohesion plays a larger part in liquids than in gases. The formation
of vapor from cold liquids, however, requires us to suppose, using the
terms of the hypothesis, that motion of the molecules has not been
annihilated by cqhesion. To be consistent, we have also to imagine
that the vapor above the liquid, for example the water in the
barometer tube in Fig. 39 (p. 116), is riot composed of the same
set of molecules one minute as it was during the preceding minute.
Their motions must cause many of them to plunge into the liquid,
while others emerge and take their places. When the water is first
introduced, there are no molecules of vapor in the space at all, so that
emission from the water predominates. The pressure of the vapor
increases as the concentration of the molecules of vapor becomes
THE KINETIC-MOLECULAR HYPOTHESIS 135
greater, hence the mercury column falls steadily. At the same time
the number of gaseous molecules plunging into the water per second
must increase in proportion to the degree to which they are crowded in
the vapor. Hence the rate at which vapor molecules enter the water
must eventually equal that at which other molecules leave the liquid.
At this point, occasion for visible change ceases and the mercury comes
to rest. We are bound to think, however, of the exchange as still
going on, since nothing has occurred to stop it. The condition is not
one of rest but of rapid and equal exchange. Such, described in
terms of the hypothesis, is the state of affairs which is characteristic
of a condition of equilibrium (p. 117). The condition is kinetic, and
not static.
To get a clear notion of a state of equilibrium, we must distinguish
two opposing tendencies, which when equilibrium is reached balance
one another. In the barometer at rest, for example, there are the two
pressures, that of the air and that of the mercury. Here, one is the
hail of molecules leaving the liquid which is constant throughout the
above experiment. It represents the vapor tension of the liquid
(p. 116). The other is the hail of returning molecules which increases
steadily, at first, as the concentration of the vapor becomes greater.
This is the vapor pressure of the vapor. These have the effect of op-
posing pressures, and when the latter becomes equal -to the former,
equilibrium is established. Both are still at work, but neither can
effect any visible change in the system. The kinetic point of view we
have here used is employed so continually in chemistry that it should
be studied attentively.
When the temperature of a liquid is raised, the kinetic energy of
its molecules is increased, the rate at which they leave its surface
becomes greater, the vapor tension increases, and, hence, a greater
concentration of vapor can be maintained. This puts into the lan-
guage of our hypothesis, the change in vapor tension with tempera-
ture (p. 116).
When the liquid is placed in an open shallow vessel, there is prac-
tically no return of the emitted molecules. Hence complete evapora-
tion takes place. Elevation of the temperature hastens the process.
A draft insures the total prevention of all returns, and has therefore
the same effect. The two methods of assisting the forward displace-
ment of an equilibrium, and particularly the second, in which the
opposed process is weakened and the forward process triumphs solely
on this account, should be considered attentively (see Chap. xv).
136
INORGANIC CHEMISTRY
Diffusion in Liquids. — We are further compelled, in applying our
hypothesis, to suppose that there is motion of the molecules inside the
liquid. When alcohol, as the lighter liquid, is floated upon water in a
cylinder (Fig. 47), the plane separating the liquids is at first easily
visible. But soon it becomes obliterated. The water diffuses upward,
and the alcohol downward, each sifting its way through the other in
spite of gravity. The complete mixing of
the liquids takes a much longer time than
in the case of two gases. It may take months.
But even here the hypothesis helps us by
pointing to the vast impediment which the
close packing of the molecules must place in
the way of the progress of any one mole-
cule. Further, once the mixture is formed,
no tendency to spontaneous separation is
ever observed. Here again, the hypothesis
shows that none is to be expected. If it oc-
curred,, it would be immediately undone by
diffusion.
Kinetic Hypothesis Applied to Solids.
— The properties of solids differ from those
of liquids chiefly in the fact that the solid
has a definite form of which it can be de-
prived only with difficulty. This we may
explain in accordance with the kinetic hypothesis by the supposition
that the cohesion in solids is very much more prominent than in
liquids. We obtain solids from liquids by cooling them; in other words,
by diminishing the kinetic energy and therefore the velocity of the
particles. The cohesive tendency of the latter is thus able to make
itself felt to a greater extent. If, conversely, we heat a solid, or,
according to the hypothesis, if we increase the speed with which the
particles move, the body first melts and gives a liquid, and this finally
boils and becomes a gas. The intrinsic cohesion of the particular sub-
stance can unde-rgo no change, but the increasing kinetic energy of the
particles steadily and continuously obliterates its effects. Yet some
motion still survives in a solid. Thus we find that xhen the layer of
silver is stripped from a very old piece of electroplate the presence of
this metal in the German, silver or copper basis of the article is easily
demonstrated.
FIG. 47.
THE KINETIC-MOLECULAR HYPOTHESIS 137
Roberts Austin has found that if bars of lead are prepared, in one end of which
an alloy containing a certain proportion of gold has been used, while the remainder
of the bar is composed of pure lead, the gold has a tendency to wander slowly into
the pure lead. The process is greatly aided by keeping the bars at a fairly high
temperature, but one much below the melting-point is amply sufficient. After a
suitable interval of time the bar may be sawn into fragments of equal length, and
its parts analyzed. The quantity of gold in a section is found to increase as we
approach the portion of the bar to which originally the whole of the gold was
confined.
The tendency of all solids to assume crystalline forms, which show-
definite cleavage and other evidences of structure, distinguishes them
sharply from liquids. The force of cohesion in liquids is exercised
equally in different directions. In solids it must differ in different
directions in order that structure may result. Since each substance
shows an individual structure of its own, these directive forces must
have special values in magnitude and direction in each substance.
A crystal arises by growth. When the process is watched, as it
occurs in a melted solid or an evaporating solution, the slow and sys-
tematic addition of the material in lines and layers, as if according to a
regular design, is one of the most beautiful and interesting of natural
phenomena. The fern-like patterns produced by ice on a window-pane
show the general appearance characteristic of crystallization in a thin
layer. A larger mass in a deep vessel gives forms which are geometri-
cally more perfect. From its very incipiency the crystal has the same
form as when, later, its outlines can be distinguished by the eye. Hence
the outward form is only an expression of a specific internal structure
which the continual reproduction of the same outward form on a larger
and larger scale leaves as a memorial of itself in the interior.
Crystal Forms. — Crystalline form is so continually used in iden-
tifying (p. 36) the substances produced in chemical actions that a list of
the kinds of forms which occur will assist in giving definite meaning
to our descriptions.
The classification of crystalline forms is carried out according to
the degree of symmetry of the crystals. Thirty-two distinct classes
are distinguished, but for our purpose a rougher division into six
groups will suffice. These groups are known by the following names:
1. Kegular system.
2. Quadratic, or square prismatic system.
3. Hexagonal system.
4. Rhombic system.
138
INORGANIC CHEMISTRY
FIG. 48.
FIG 49.
5. Monosymmetric, or monoclinic system.
6. Asymmetric, or triclinic system.
The regular system presents the most symmetrical figures of all.
Some forms which commonly occur are the octahedron
(Fig. 48) shown by alum, the cube (Fig. 6, p. 13)
affected by common salt, and the dodecahedron
(Fig. 49) frequently assumed by the garnet.
The square prismatic system in-
cludes less symmetrical forms than
the previous one, since the crystals
are lengthened in one direction. Fig.
50 shows the condition in which zir-
con (ZrSi04), which furnishes us with the basis of
certain incandescent illuminating arrangements, occurs
in nature. The form of ordinary hy-
drated nickel sulphate (NiS04, 4H20) is similar to this.
The hexagonal system, like the preceding, frequently
exhibits elongated prismatic forms, but the section of
the crystals is a hexagon instead of a square, and the
termination is a six-sided pyramid. Quartz (Fig. 51), or
rock crystal, is the most familiar min-
eral in this system. Calcite (CaC03),
which is chemically identical with
chalk, or marble, takes forms known as the scaleiio-
hedron (Fig. 52) and rhombohedron
(Fig. 9, p. 14), which are classified in
a subdivision of this system. Indeed,
recently it has become common to
erect this into a separate system (the
trigonal), in which both quartz and
calcite are included.
The rhombic system includes the
natural forms of the topaz, and of
sulphur (Fig. 1, p. 11), as well as that
of potassium permanganate (Fig. 53), potassium nitrate
(Fig. 98), and many other substances. These crystals
FIG. 62. exhibit a good deal of symmetry, but their section is
always rhombic, and hence the name.
The monosymmetric system exhibits forms which have but one
plane of symmetry. Gypsum (Fig. 54), which is hydrated calcium
FIG. so.
FIG. 51.
THE KINETIC-MOLECULAR HYPOTHESIS
139
sulphate (CaS04, 2H20), and felspar are minerals possessing forms of
this kind. Tartaric acid, rock candy (Fig. 55), potassium chlorate, and
hydrated sodium carbonate (washing soda) belong to this system.
The asymmetric system includes forms
which have no plane of symmetry whatever.
Blue vitriol (Fig. 41, p. 120), (CuS04, 5H,0),
is one of the most familiar substances of this FIG> 53.
kind.
The forms of crystals which we may actually make seldom corre-
spond exactly with the figures. If we allow a crystal to grow upon
the bottom of a vessel, for example, it will usually
have a tendency to spread itself out parallel to the
surface of the glass, and when taken up for examina-
tion will be found to present a somewhat distorted
form. By changing, at frequent intervals, the face
on which the crystal stands, however, uniform growth
in all directions is secured. Hanging a small
crystal by a thread insures almost ideal develop-
ment. Yet the form even of distorted crystals can
readily be recognized by suitable means. The
shape of the faces may indeed be extremely mis-
FIG. 54. leading. We find, however, that the angles at which
the faces meet are always the same, whatever dis-
proportionate growth may have' occurred in the development of the
crystal.
Since, in general, each substance has a form of its own, no other
substance, as a rule, can be used even partially in building up the
crystal (see, however, Isomorphism). This
fact is taken advantage of in order to separate
chemical substances from impurities. The
impure body is first dissolved in some solvent.
The preponderating substance in the mixture,
unless it is very much more soluble than
the impurity, will then usually give pure
crystals while the foreign body remains in solution.
The shapes of gems must not be confused with crystalline forms.
The original crystals are cut and polished to a new form specially
adapted to increase the ornamental value of the stone by causing it to
reflect more light (see Diamond). Again, glass is really a very viscous
fluid (amorphous body, p. 123), and has no structure or form of its own.
FIG. 55.
140 INORGANIC CHEMISTRY
The word " crystal " applied to cut glass would therefore be misleading
if taken literally.
Crystal Structure. — As the above would lead us to expect, the study of
crystallized substances shows that their peculiarities are not confined to the out-
side layer. The outline represents a certain structure which permeates the whole
mass. In crystals of the regular system, many of the ordinary physical properties
are the same as those of an amorphous substance, like glass. For example, if we
turn a sphere out of crystallized salt and hang it in pure water, we find that solu-
tion takes place at a uniform rate all over the surface. This is not the case, however,
with substances from any of the other systems. Spheres cut from substances be-
longing to the second and third systems would dissolve more, or less, rapidly in the
direction of the chief axis than in any other direction, and so ellipsoids of revolu-
tion would quickly be produced. In the other three systems, more complex forms
would result.
The tenacity of crystals, to whatever system they may belong, is different in
different directions. Thus, a crystal of salt has a cleavage parallel to any of the
faces of the cube, and, therefore, splits most easily in one of three directions at
right angles to each other. Calcite, whatever the outward form of the crystal,
always cleaves so as to give a rhombohedron. Fluorite (CaF2), although almost
always cubical in form, splits when broken so as to give an octahedron.
The behavior of crystals towards light is also extremely interesting. The rate
at which light moves through crystals of the regular system is the same in all
directions. In other crystals, however, we find that it moves with a different
speed in different directions, the variations in speed being likewise related to the
outward form. Finally, if a thin slab of rock salt is covered with wax and the
point of a heated cone of metal is placed in the center, the wax melts uniformly
in a circle around the point, indicating that the heat is conducted with equal speed
in all directions. The way in which the slab has been cut with reference to the
surface of the crystal, in the case of substances of the first of the above systems,
has no effect upon this result. In all other cases, however, the zone of melting
wax is in general elliptical, or even more complex in form, according to the system
to which the substance belongs and the direction in which the slab has been cut.
Molecular Magnitudes. — Taking the properties of matter on
the one hand and the assumptions of the molecular hypothesis on the
other, we can estimate the degree of mutual proximity which must be
assigned to the molecules in order that the hypothesis may be consist-
ent with the facts. By considering the thickness, or rather the thin-
ness, of thin films, certain properties of gases, and other phenomena,
fairly coincident, independent values have been obtained. According
to Lord Kelvin, 1 c.c. of gas at 0° and 760 mm. contains not less than
1020 (that is, 1 followed by twenty ciphers) molecules. A globe of
water as large as a football, if magnified to the size of the earth, the
component molecules being magnified at the same time, would show a
THE KINETIC-MOLECULAR HYPOTHESIS 141
structure coarser than a heap of small shot but less coarse than a
heap of footballs. The former estimate, in regard to gases, leads by
calculation to the conclusion that one molecule of hydrogen weighs
0.9 x 10~24 g., or 9 over 1 followed by twenty-five ciphers of a gram.
The speed of a hydrogen molecule between collisions must be about
1840 meters per second. That of an oxygen molecule must be one-
quarter of this, being, according to the law of diffusion (p. 108),
inversely as the square roots of the densities.
Formulative and Stochastic Hypotheses. — The nature and
use of an hypothesis like the one we have been discussing will now be
evident. It is a structure existing in the imagination. It cannot
exist anywhere else, because it includes novelties like perfectly elastic
bodies in perpetual motion. In making it, we are well aware of the
inverity of some of its elements. But then, as will have been
observed, we do not attempt to verify the hypothesis itself. We did
not make it in order to have before us the actual structure of matter,
but in order to have a sort of mechanical moving diagram which
should assist us in following the behavior of matter. It is like a scaf-
folding, constructed to enable us to examine or work upon a difficultly
accessible part of a building, which we never for a moment think of
as being a part of the building. It is a sort of formula. The algebraic
formula represents magnitudes ; the geometrical, directions and dimen-
sions. The formula in physics, by the use of mathematical conven-
tions, pictures, for example, some mode of behavior of matter. For
so concrete a subject, however, the mathematical mode of expression
is intensely abstract. And so a representation in terms of mechanism,
which is still a formula, is frequently resorted to. The molecular
hypothesis is therefore a formula consisting of imaginary machinery.
Its object is simply to help us in organizing or formulating knowledge
on a certain subject. Hence we name it a formulative hypothesis.
A formula, in the very general sense in which we have used the word above,
is anything which has certain properties of the nature of form in common with
some other thing or relation of things. For example, the plan of a city and a
blackboard diagram are formulae. In chemistry, a "model" of some carbon
compound, made up of balls of wood and wires, is just as much a formula as are
the written symbols showing the constitution of the substance. It is intended to
explain the behavior of the substance. From this sort of mechanical formula it is
but a step to one consisting of moving particles and intended to explain the be-
havior of gases. The fact that this model cannot actually be constructed, because
perfectly elastic materials are not available, does not alter the case one whit.
142 INORGANIC CHEMISTRY
The value of the present formulative hypothesis, and of formulative hypotheses
in general, is shown by the history of science. Dalton, who first worked out a
clear conception of the independent behavior of mixed gases (p. 88), used this hypoth-
esis constantly in that work. Clear understanding of the separate existence of
aqueous vapor in the air could be reached by him only by thinking of each ma-
terial as being made up of independent molecules. Any other mode of conceiving
the mixture that might readily occur to one would involve some adhesion or inter-
ference of the two substances. His recognition of the independent solubilities of
mixed gases, in proportion to the partial pressures of each, would have been de-
layed or prevented altogether if, in studying the results of his experiments, he had
not reached it by way of this hypothesis.
Other formulative hypotheses, showing the same contradiction of plain facts
in their fundamental assumptions that we have noted in the molecular hypothesis,
have been and are in common use. The conception of light as consisting of cor-
puscles was such an hypothesis in its day. The imponderable-matter (note the
contradiction) view of heat was another. The undulatory theory of light, which
postulates a perfectly elastic ether, weightless, frictionless, and lacking every
trace of impenetrability, is an hypothesis showing the same inverifiable elements.
It is handled characteristically also, for we do not try by its means to learn
more about ether, but more about light.
An essentially different kind of hypothesis is constantly used at
every step in investigation, although it is seldom mentioned in books.
When Mitscherlich discovered that Glauber's salt (p. 121) gave a definite
pressure of water vapor, he at once formed the hypothesis, that is,
supposition, that other hydrates would be found to do likewise. Ex-
periments showed this supposition to be correct. The hypothesis was
at once displaced by the fact. This sort of hypothesis predicts the
probable existence of certain facts or connections of facts, hence, reviv-
ing a disused word, we call it a stochastic hypothesis * (Gk. O-TOXCIO-TIKOS,
apt to divine the truth by conjecture). It differs from the other kind
in that it professes to be composed entirely of verifiable facts and
is subjected to verification as quickly as possible. In the case of
a formulative hypothesis we have no expectation, or at best a very re-
mote one, of verifying the hypothesis, because many of its essential
constituents are contrary to experience. At all events, our efforts are
bent, not to verifying the hypothesis itself, but to verifying the rela-
tions between facts which it suggests.
We may define a formulative hypothesis as follows : A structure, the
essential parts of which are assumed facts or connections of facts more or less
inconsistent with known facts, used in formulating other known facts. It
achieves this by virtue of certain logical and formal correspondences which exist
* The author owes to Professor Paul Shorey, head of the department of Greek
in the University of Chicago, the suggestion of this word.
THE KINETIC-MOLECULAR HYPOTHESIS 143
between its abstract qualities and those of the facts it is employed to explain.
The verification of the assumed facts is not in question, since their inverity is
one of the premises, but that of the relations between the ascertained facts which
emerge, is the step to which the making of the hypothesis was only a preliminary.
This sort of hypothesis, therefore, is a sort of formula, or has the properties of a
formula.
A stochastic hypothesis is : A supposition of the existence of certain facts,
or connections of facts. It is reached by deduction from other facts or con-
nections of facts already ascertained, and is intended to be subjected to ultimate
verification.
The word theory is applied to both kinds of hypotheses. A study of the
usages of the word and of the definitions given by the best authorities does not
reveal the existence of any essential qualitative difference between a theory and
an hypothesis, so that a sharp distinction cannot be drawn between them. To
avoid confusion, we shall use the word " theory" as a rule only in the sense in which
it occurs in phrases like "the theory of heat," or "theory and practice." Here
it refers to an aggregate of conceptions and largely or wholly verified generaliza-
tions and laws which constitute the abstract (as distinct from the concrete) state-
ment of the content of some branch or phase of knowledge, so far as this has
undergone successful organization.
Lest it should be imagined that the discovery of a sufficiently extensive and
complete correspondence between a formulative hypothesis and the facts with
which it deals constitutes, ipso facto, a proof of the verity of the hypothesis itself,
it must be stated explicitly that such an inference is entirely illogical. Yet this
fallacy is one of the commonest into which the student of science falls. This
inference would be equivalent to asserting the impossibility of devising any other
hypothesis which should correspond equally well with the facts. Apart from the
formal illogicality of such an inference, experience has taught us that sooner or
later we always encounter some feature in the behavior of the subject of the
hypothesis which is different from that which the hypothesis would have led us to
expect. Thus the corpuscular hypothesis to explain light, after long and useful
service, had finally to be discarded (see also Atomic hypothesis).
The philosophy of the molecular hypothesis is discussed by Stallo, Concepts
and Theories of Modern Physics, Chap, vii ; Pearson, Grammar of Science, Chap,
vii; James Ward, Naturalism and Agnosticism, Lectures iv and v, and many
other writers.
The value of the kinetic-molecular conception as a formulative
hypothesis has been illustrated in connection with critical phenomena,
and will appear again when we deal with solutions and chemical equi-
librium. But indeed its field of application is coextensive with the
science itself.
The conceptions of the kinetic hypothesis are of especially great assistance in
rationalizing chemical manipulation and so hastening the acquisition of an intelli-
gent control of it. The constructive imagination, on the use of which experi-
mental work depends so much for its success, must have something to work with,
and this hypothesis furnishes a tool such as it requires. Methods taught by rule
144 INORGANIC CHEMISTRY
of thumb are slowly learned and constantly fail in application. We may be told
a dozen times that using reagents in finely powdered, or metals in granulated con-
dition hastens all interactions, and still never think of this abstraction when
working. But if it is suggested that, in terms of the kinetic hypothesis, molecules
must meet freely in the same medium to react easily, and that, therefore, the
larger the surface the more copious will be the supply of molecules dissolving, we
are likely to form a conception of the reason for the procedure that will be
When a student is told to concentrate a solution and set it aside to crystallize,
why does he evaporate the liquid rapidly to dryness, and still expect the residue
to appear in large, well-formed crystals ? Because he has no notion of the necessary
slowness of the process. But if he has the idea that it is like building a house,
one stone at a time, and that there are far more units to be laid down according to
plan in making the smallest visible crystal than there are bricks in building the
largest factory, his procedure may promptly become more rational.
CHAPTER X
SOLUTION
WE have frequently made use of the fact that certain substances
form with others homogeneous systems which we call solutions.
Sometimes this property is taken advantage of for separating ma-
terials, as in the case of the removal of sulphur from admixture with
iron and ferrous sulphide (p. 11). In other cases we carry out the
interaction of chemical substances, by first dissolving them in some
liquid and then mixing the solutions. The liquid, commonly water, is
used as a vehicle for one or more of the substances, and takes no part
in the chemical change. Thus some knowledge of the properties of
solutions is absolutely necessary in order that we may employ them
intelligently. In what follows, we shall give a preliminary account of
some of the simpler facts about solution.
General Properties of Solutions. — A solid may be distributed
through a liquid either by being simply suspended (p. 113) in the
latter (mixture) or by being dissolved in it (solution). Similarly a
liquid may be suspended in droplets in another, as in milk (emulsion),
or it may be dissolved. It is usually easy to distinguish between the
two cases, for a suspended substance settles or separates sooner or
later, while a dissolved one shows no such tendency. The cases are
exceptional where the subdivision of a suspended substance is so
minute (colloidal solution) as to make its retention by filter paper
impossible. If a liquid is opalescent or opaque, then we have a case
of suspension. A solution is a clear, transparent, perfectly homo-
geneous liquid, in which the dissolved substance seems to have been
dispersed so completely that the liquid cannot be distinguished by the
eye from a pure substance.
There is no limit to the amount of dissipation which may thus be
produced. A single fragment of potassium permanganate, for ex-
ample, which gives a very deep purple solution in water, may be dis-
solved in a liter or even in twenty liters of water, and the purple
tinge which it gives to the liquid will still be perfectly perceptible in
145
146 INORGANIC CHEMISTRY
every part of the larger volume. The characteristics, therefore, of
solution are absence of settling, homogeneity, and extremely minute
subdivision of the dissolved substance.
The Scope of the Word. — The word is used for other systems
than those containing a solid body dissolved in a liquid. Thus, liquids
also may be dissolved in liquids, as alcohol in water. Again, if we
warm ordinary water, bubbles of gas appear on the sides of the vessel
before the water has approached the boiling-point. They are found
to be air. Further study of the subject shows that agitation of any
gas with water results in the solution of a large or small quantity of the
gas, and heat will usually drive the gas out again. It appears there-
fore that solids, liquids, and gases can equally form solutions in liquids.
The absorption of hydrogen by palladium (at all events after a
certain point) and by iron takes place in accordance with the same
laws as the solution of solids in liquids, and the results may be
described therefore as true solutions. Liquids are in some cases
absorbed by solids, and homogeneous mixtures of solids with solids
are perfectly familiar. The sapphire is a solution of a small amount
of a strongly colored substance, in a large amount of colorless aluminium
oxide. It may therefore be stated that solution of gases, liquids, and
solids in solids appears to be possible.
Limits of Solubility. — The next question which naturally occurs
to us is as to whether, the mingling of two substances in this manner
has any limits. We find that the results of experiment in this
direction may be divided into two classes. Some pairs, of liquids
particularly, may be mixed in any proportions whatever. Alcohol and
water is such a pair. On the other hand, at the ordinary laboratory
temperature, we can scarcely take a fragment of marble (CaC03) so
small that it will dissolve completely in 100 c.c. of pure water. Under
the same conditions any amount of potassium chlorate up to 5 g. will
almost completely disappear after vigorous stirring, while 90 g. of
ordinary Epsom salts (hydrated magnesium sulphate), but not more,
may be dissolved in about the same amount of water. In fact, most
solids may be dissolved in a liquid only up to a certain limit, which
with different solids may range from a scarcely perceptible to a very
large amount. No substance is absolutely insoluble. But for the
sake of brevity we call marble, for example, " insoluble " because in
most connections it may be so considered.
SOLUTION
147
Recognition and Measurement of Solubility. — The only
method of recognizing with certainty whether a solid is soluble in a
liquid or not is to filter the mixture and evap-
orate a few drops of the filtrate on a clean
watch-glass. For learning how much of the
body is contained in a given solution, a weighed
quantity of the solution is evaporated to dry-
ness and the weight of the residue deter-
mined. When the dissolved substance is
volatile, its presence is often shown by some
chemical test (p. 99).
Ether and water is a case typical of the
behavior of two liquids, each somewhat soluble
in the other. After being shaken together,
they seem to separate again completely into
two layers (Fig. 56) with the ether uppermost.
If, however, the water is withdrawn from
beneath the ether, we find that, when heated,
it gives off quantities of ether vapor which can
be set on fire. Conversely, the addition of
anhydrous cupric sulphate to a sample of the
ether shows the presence of water in the latter,
for the blue hydrated form of the substance is at once produced. In
some common cases the maximum solubilities at 22° are as follows :
FIG. 56.
SUBSTANCE.
GRAMS OF SUBSTANCE
IN 100
GRAMS OF WATER.
GRAMS OF WATER
IN 100
GRAMS SUBSTANCE.
Alcohol
No limit
No limit
Ether
2.16
11.02
Chloroform
0 64
0 10
Carbon disulphido
1.24
0 13
It must be stated explicitly that in going into solution, as we have
used the term, a compound dissolves as a whole, and, if the compound
is pure (p. 34), any residue has the same chemical composition as the
part which has dissolved.
Terminology. — In order to describe the relations of the compo-
nents of a solution, certain conceptions and corresponding technical
expressions are required. It is customary to speak of the substance
148 INORGANIC CHEMISTRY
which, like water in most cases, forms the bulk of the solution, as the
solvent. To express the substance which is dissolved, the word solute
is frequently used, and will be employed when we wish to avoid cir-
cumlocution. The term "strength" is too indefinite for scientific
purposes. It may imply activity, or power of resistance, or pungency
(in an odor), or, as in the case of solutions, it may be a measure of
quantity. The amount of the substance which has been dissolved
by a given quantity of the solvent is therefore described as- the
concentration of the solution. A solution containing a small propor-
tion of the dissolved body is called dilute ; it has a small concentration.
One which contains a larger amount is more concentrated. Very
" strong " solutions are frequently spoken of simply as concentrated
solutions. The partial removal of the solvent by evaporation is called
concentrating, its total removal evaporating to dryness. Finally, since
there is a limit to the solubility of most substances, a solution is
described as saturated when the solute has given as much material to
the solvent as it can. This state is reached after prolonged agitation
with an excess of the gas, the liquid, or the finely powdered solid, as the
case may be. The larger the excess, the sooner saturation is attained.
The maximum concentration attainable in this way is called the
solubility of the substance in a given solvent.
The distinction between solute and solvent is made merely for convenience.
Theoretically there is no distinction between the components of a solution.
The concentrations of solutions, saturated and otherwise, are some-
times expressed in physical, and sometimes in chemical, units of
weight. When physical units are employed, as in the above table, we
give the number of grams of the solute held in solution by one hundred
of the solvent, or, occasionally, the number of grams in one hundred
of the solution.
When chemical units of weight are employed, two different plans are
possible, and both are in use. Either the equivalent (p. 49) or the atomic
weights may be taken as a basis of measurement. In the former case,
the solutions are called normal solutions, and in the latter, for a reason
which will appear later (Chap, xii), molar solutions.
A normal solution contains one gram-equivalent of the solute in one
liter of solution. The word "equivalent" has been used hitherto only of
elements, and this application of the expression involves an extension
of its meaning. An equivalent -weight of a compound is that amount
of it which will interact with one equivalent of an element. Thus, a
SOLUTION
149
formula-weight of hydrochloric acid HC1 (36.5 g.) is also an equivalent
weight, for it contains 1 g. of hydrogen, and this amount of hydrogen is
displaceable by one equivalent weight of a metal. A formula-weight
of sulphuric acid H2S04 (98 g.), however, contains two equivalents of
the compound, and a formula-weight of aluminium chloride A1C13
(133.5 g.) three equivalents. Hence normal solutions of these three
substances contain respectively 36.5 g. (HC1), 49 g. (H2S04), and 44.5
g. (AlClg) per liter of solution. The special property of normal solu-
tions is, obviously, that equal volumes of two of them contain the exact
proportions of the solutes which are required for complete interaction.
Solutions of this kind are much used in quantitative analysis.
Solutions of different concentrations all prepared on the above
basis are named as follows, and are often indicated by the abbreviations
appended :
QUANTITY" OF SOLUTE
PER LITER.
NAME.
A BBREVIATION.
One hundredth of one gram-)
equivalent j
Centi-normal
m or -01 N
One tenth of one gram-equiv- j
alent j
Deci-normal
N
To or -IN
One half of one grain-equiva- j
lent I
Semi-normal
N * *r
-3- or .5 N
One gram-equivalent
Two and a half gram-equivalents
Normal
Two and a half normal
N
2J N or 2.5 N
A molar solution contains one mole (gram-molecular weight) of the
solute in one liter of solution. When molecular formulae (see Chap, xii) are
used, this means one gram-formula weight per liter. Iii the cases cited above, the
molar solution contains 36.5 g. (HC1), 98 g. (H2SO4), and 133.5 g. (A1C13) per liter.
As will be seen, the concentration of molar and normal solutions are necessarily
identical when the radicals are univalent. Other concentrations are described as
deci-molar (M/10 or .IJf), two and a half molar (2.5 .M), and so forth, on the same
plan as before.
There is also a chemical unit of volume (see Chap, xii) which is the volume
occupied by a mole (gram-molecular weight) of a gas (or dissolved substance) at
0° and 760 mm. pressure (gaseous or osmotic). This volume averages 22.4 liters, and
is called the gram-molecular or molar volume (G. M. V.). The unit of concentra-
tion for many theoretical purposes is, therefore, that of one mole in 22.4 liters.
Solution One of the Physical States of Aggregation of Matter.
— When a solid body dissolves in a liquid, the properties of the body
undergo a very marked change, which to all appearance might be
150 INOEGANIC CHEMISTRY
chemical. Yet, besides the ease with which a liquid may be removed
by evaporation and the solid recovered unchanged, we note particularly
that the concentration of a saturated solution cannot be expressed in
terms of integral multiples of the chemical combining weights. If,
therefore, the process were to be regarded as chemical, several impor-
tant generalizations would have to be revised or discarded (cf. p. 48).
We shall see also, in a later paragraph, that the quantity of a solid
which a liquid may take up varies with the slightest change in tem-
perature. Now we do not find the composition of chemical compounds
so to vary. The solution of a solid may therefore be likened to a
change in state, such as the conversion of a liquid into a gas or a solid.
As there is danger of confusion arising, we may repeat that a com-
pound is homogeneous and its composition is expressible in chemical
units of weight ; a saturated solution is homogeneous but its concentra-
tion varies with temperature so that chemical units cannot be used to
describe its composition ; a mixture, as of two solids or two liquids, is
neither homogeneous nor in any way definite in composition.
Kinetic- Molecular Hypothesis Applied to the State of Solu-
tion. — Accepting solution as a physical state, we may now apply the
same formulative hypothesis to the explanation of the behavior of a
substance in, solution as to matter in the gaseous or liquid states. We
saw that a solid body, which is ordinarily condensed in a small space,
can be disseminated by the use of a solvent through a very large one.
The molecules of the solid become scattered like those of a gas or vapor
through a much greater volume. We may regard the dissolved sub-
stance as being, practically, in a gaseous or quasi-ga,seou.s condition.
The molecules are torn apart from one another, their cohesion is over-
come, and their freedom of motion is in a measure restored. It is true
that they could not continue to occupy this large volume for a moment
in the absence of the solvent. But we may bring this into relation
with the case of a vapor by saying that a solid body, like common salt,
can only evaporate appreciably at the ordinary temperature, and
occupy a large space, when that space is already filled with a suitable
liquid. The latter acts as a vehicle for the particles of the solid. A
volatile liquid, on the contrary, can dissolve in an empty space and fill
it with its particles without any vehicle being required.
This conception of the quasi-g&seous condition of a dissolved sub-
stance would be simply fantastic if it did not lead us to a better under-
standing of the behavior of solutions. Now there are certain proper-
SOLUTION
151
ties of gases already discussed which we should naturally proceed to
look for in the case of dissolved bodies if this theory were correct.
These are the properties described under the law of diffusion, Boyle's
law, and Charles' law.
It is easy to show that, if we place a quantity of the pure solvent
(Fig. 57) above a concentrated solution of a substance, and then set
the arrangement aside, the
dissolved body slowly makes
its way through the liquid
(Fig. 58), obliterating the
original plane of separation.
Eventually the dissolved
body scatters itself uniform-
ly through the whole. In
other words, the particles of
the dissolved substance ex-
hibit the property of diffu-
sion in the same way as do
those of gases.
When the diffusion of a
gas is resisted by a suitable
partition, we find pressure is
exercised upon the walls of FIG. 57. FIG. 58.
the vessel and upon the
partition. It is possible to show that the particles of a dissolved sub-
stance exercise a pressure of a very similar kind. This pressure is
spoken of as osmotic pressure (g.v.'). In order to get evidence of this
pressure, we must place a partition between the solution and the pure
solvent (Fig. 57), and we must choose the material of this partition so
that its substance gives free passage to the solvent while it resists the
exit of the dissolved material. The particles of the latter, then, by
their impacts upon it, produce the effect of pressure exactly as in the
case of gases. This pressure is found to be proportional to the con-
centration of the solution. In other words, it depends upon the degree
of crowding of the molecules of the dissolved substance. It corre-
sponds therefore exactly to the pressure of gases in this respect, and
Boyle's law (p. 81) expresses the relation equally for both.
It may be added, also, that osmotic pressure increases when a
solution is warmed, and it gains ^3 of the value it had at 0° for every
degree through which the temperature is raised. In other words,
152
INORGANIC CHEMISTRY
Charles' law (p. 87) describes the change in osmotic pressure with
change in temperature just as correctly as it does the change in gaseous
pressure.
Kinetic-Molecular Hypothesis Applied to the Process of
Solution. — We may now apply the same hypothesis to the process of
dissolving, with a view more especially to explaining why the process
of dissolving ceases, in spite of the presence of excess of the solute,
when a certain concentration has been reached. If some of the mate
rial dissolves, why not more ?
Let us suppose that it is the dissolving of common salt in water
(Fig. 59) which we wish to explain in detail. We believe that in the
solid substance the molecules are somewhat
closely packed together, while in the solution
they are rather sparsely distributed. If there
were no water over the salt, practically none
of the particles of the latter would be able to
leave the solid and enter the space above. Thus,
the process of solution must consist in the
loosening of the molecules on the surface and
their passage into the liquid. By diffusion, the
free molecules will gradually move away from
the neighborhood of the surface of the solid and
make room for others, and thus, if the system
remains undisturbed, the liquid will eventually
become a solution of uniform concentration. If
a large enough amount of the solid has been
provided, the ultimate condition will be that of
a saturated solution with excess of the solid
beneath. If we had proper means of measuring
it, the tendency of the molecules to leave the solid in the presence
of a given liquid would give the effect of a kind of pressure. This is
spoken of as solution pressure.
Now the molecules, after having entered the liquid, move in
every direction, and consequently some of them will return to the
solid and attach themselves to it. The frequency with which this
will occur will be greater as the crowding of particles in the liquid
increases, so that a stage will eventually be reached at which the
number of molecules leaving the solid will be no greater than that
landing upon it in a given time. If the whole of the liquid has mean-
FlG. 59.
SOLUTION 153
while become equally charged with dissolved molecules, there will be no
chance that the field of liquid immediately round the solid will lose
them by diffusion, so that a condition of balance or equilibrium will
have been established : NaCl (solid) <z±NaCl (diss'd). The motion of
the particles in the liquid produces what we have called osmotic press-
ure ; and when the osmotic pressure, by the continual increase in the
number of dissolved molecules, becomes equal to the solution press-
ure, increase in concentration of the solution ceases. It is at this
point that we speak of the solution as being saturated with respect to
the particular substance dissolving. The analogy to vapor tension
and vapor pressure (p. 135) is evident.
The necessity of distinguishing between the fictions used in thinking about and
describing the phenomena of solution and the facts themselves must be empha-
sized here as it was in the preceding chapter. The arrangement and expression of
the facts about solutions in terms of the kinetic-molecular hypothesis is known as
the theory of solutions.
Independent Solubility. — Just as two gases, when mixed, are
independent of one another (p. 88), and have severally the same
pressure, solubility, and so forth, as they would possess if each alone
occupied the same space, so is it with dissolved substances. In gen-
eral, a volume of water, in which a moderate amount of some substance
has been dissolved, will take up as much of a second substance as
would an equal volume of pure water. Thus, water containing some
sugar will dissolve as much sodium chloride as the same amount of
pure water. In the point of view of the kinetic-molecular hypothesis,
the dissolved molecules of sugar have no connection with, or influence
upon, the mechanism which determines the solubility of the salt,
namely, the exchange of salt molecules between the suspended, dis-
solving crystals and the solution.
Naturally this principle of independent solubility does not hold
with any degree of exactness when the concentration of the substance
already present is great. It fails also when the two solutes interact
chemically, as will usually be the case, for example, when each is an
acid, base, or salt (see Chap. xvi). The solubility is affected in an
especial degree when the two substances have one radical in common,
as when they are nitric acid and a nitrate, or two chlorides (see Ionic
equilibrium).
Solution of a Gas in a Liquid. — The same conceptions may be
used to explain any case of solution. Let us take that of oxygen con-
154
INORGANIC CHEMISTRY
ducted into a bottle which is partially filled with water (Fig. 60), no
other gas being present in the space above the liquid. As the mole-
cules of the gas impinge upon the liquid, some of them pass into it
and dissolve. The particles which have thus gained access to the
liquid move about in every direc-
tion, and, as they become more
and more numerous, a larger and
larger number will escape from
the surface and pass back into the
gaseous condition. At first, this
reaction will be slight, but event-
ually, as the solution increases in
concentration, it must become
equal in rate to the process of
solution itself. It is assumed
that the supply of gas is main-
tained at a uniform pressure, and
therefore uniform molecular con-
centration, during the whole pro-
cess. Once more we shall have
a state of balance or equilibrium,
FlG- 6°- and the liquid will be saturated,
this time with a gas : 0 (gas) «r± 0
(diss'd). It is found, as the hypothesis would lead us to expect, that
the concentration of the saturated solution of a gas is proportional to
the pressure at -which the gas is supplied. This is usually known as
Henry's law.
Henry's law is anticipated from the theory when we consider that the pressure
of the gas is proportional to the degree of crowding of its molecules. On the
other hand, the concentration that can be maintained in the saturated solution is
proportional to the frequency of the impacts of the molecules of the gas upon the
surface of the liquid, and this in turn will be proportional to the degree to which
they are crowded in the space above the liquid. Hence the concentration of the
saturated solution is proportional to the gaseous pressure.
The solubility of different gases varies much. One volume of
water will dissolve 1050 volumes of ammonia at 0° and 760 mm.,
while it will dissolve only about 0.02 volumes of hydrogen under the
same conditions. In one volume of alcohol, at 0° and 760 mm., 17.9
volumes of hydrogen sulphide or 0.07 volumes of hydrogen may be
dissolved. The law describes the behavior of the bodies with exact-
SOLUTION 155
ness only when low gaseous pressures and gases whose solubility is
small are in question. Great solubility must be due in part, probably,
to chemical union between the material of the gas and the solvent, or
to cohesive influences which the molecules of the dissolved gas exert
upon each other. The hypothesis, on the other hand, considers only
an ideal behavior involving complete chemical and physical independ-
ence of the molecules.
When more than one gas is in contact with the solvent, the hy-
pothesis enables us successfully to foretell what will happen. The
quantity of each gas which can remain dissolved must depend simply
upon the frequency with which its own molecules strike the liquid,
and must be independent of the presence of the other gas. Hence the
solubility of each gas is the same as if it were present alone at its
own partial pressure (p. 88). Dalton used these very considerations
in drawing this conclusion from his experimental data.
Air dissolving in water is an illustration of this principle. It does
not dissolve as a whole, but the oxygen and nitrogen dissolve each in
proportion to its intrinsic solubility and partial pressure.
It is easy, by the use of this law, to form an approximate estimate of the pro-
portion of oxygen to nitrogen in the dissolved gases. The air may be taken to be
at 760 mm., and its composition by volume roughly 1/5 oxygen and 4/5 nitrogen.
The separate solubilities of the gases at 760 mm. are, respectively, 4 and 2
volumes in 100 volumes of water. Their partial pressures being 1/5 and 4/5 of an
atmosphere, the amounts actually dissolved will be 4 x 1/5 = 0.8 and 2 x 4/5 =
1.6 in 100 volumes of water. 'The ratio of free oxygen to nitrogen in the water
will therefore be 1 : 2.
Two Immiscible Solvents : Law of Partition. — An interesting
application of the same ideas may be made to a case which occurs
very commonly in chemical work. If we shake up a small particle of
iodine with water, we find that it dissolves slowly, giving eventually a
saturated but very dilute solution. If now ether in sufficient quantity
be shaken with the aqueous solution, the greater part of the iodine will
find its way into the ether, and be contained in the brown layer which
rises to the top. The process of removing a substance partially from
solution in one solvent and securing it in another is called extraction.
We find in such cases that neither solvent can entirely deprive the
other of the whole of the dissolved substance, if the latter is soluble
in both independently : I (in Aq.) ^± I (in ether). The partition of
the substance takes place in proportion to its solubility in each sol-
vent. It is found that any amount of the solute, up to the maxi-
156 INORGANIC CHEMISTRY
mum the system can contain, provided this does not involve too high
a concentration in either solvent, is divided so that the ratio of the
concentrations in the two solvents is always the same. In the case of
iodine divided between water and ether, this ratio is about 1 : 200.
The aqueous solution of potassium iodide has a very great power of dissolving
iodine, and we find that in the presence of this salt the ether leaves a much larger
share of the element in the lower layer. A part of the iodine combines to form
KI3, however, so that this is not a case of simple solution, and the law of partition
does not hold. The chemical equilibrium (see Chap, xv) between free and com-
bined iodine in the aqueous layer has to be considered.
Influence of Temperature on Solubility. — The quantity of a
substance which we can dissolve in a fixed amount of a given solvent
depends very largely upon the temperature of both. Usually the solu-
bility increases with rise in temperature. Measurements may be made
by the method described before (p. 147), using excess of the finely pow-
dered solute with different portions of the same solvent in vessels kept
at different temperatures. The most useful way of representing the re-
sults is to plot them graphically. The diagram (Fig. 61) shows the
curves for a few familiar substances. The ordinates represent the num-
ber of grams of the anhydrous compound which is held in solution by
100 g. of water in each case. The abscissae represent the tempera-
tures. The concentration for any temperature can be read off at once.
Thus 100 g. of water holds 13 g. of potassium nitrate in solution at 0°
and 200 g. at 87°. The increase in solubility is here enormous. On
the other hand, the same quantity of water will hold 35.6 g. of sodium
chloride in solution at 0° and 40 g. at 100°. The difference is shown
at once when we examine the curves and observe that the line repre-
senting the solubility of sodium chloride scarcely rises at all between
0° and 100°, while that of potassium nitrate is extremely steep.
Cases in which the solubility decreases with rise in temperature are
less common. When cold water is saturated with calcium citrate, and
the solution is then warmed, a large part of the salt is quickly pre-
cipitated.
When triethylamine, an organic base, N(C2H5)3, liquid at ordinary tempera-
tures, is added to cold water until no more will dissolve, the solution, which is
perfectly clear and transparent, on being warmed with the hand at once becomes
clouded from the separation of the two liquids. A comparatively slight elevation
in temperature causes a separation into two distinct layers.
Phases. — We can frequently abbreviate our statements by using
two words of broad significance, one of which has already been em-
SOLUTION
157
20° 30° 40° 50° 60° 70° 80° 90° 100°
Temperature
FIG. 61.
158
INORGANIC CHEMISTRY
ployed. A set of materials in or tending towards a condition of equi-
librium is called a system. The discrete parts of an inhomogeneous
system are called its phases. Thus, a liquid with its vapor forms a
system with a liquid phase and a vapor phase. A saturated solution
is a system with three phases, the undissolved excess of the solute
(solid phase), the solution, and the vapor.
Equilibrium in a Saturated, Solution. — Once a solution has be-
come saturated, the dissolving substance remains thereafter unchanged
in amount no matter how long the materials are left in contact. In
technical terms, the quantity of each phase has no influence on the
concentration of any
of them. A greater
excess of the solute
forces no more mat-
ter into solution than
does a small excess.
It should be
clearly understood
that the kinetic hy-
pothesis requires us
to assume that an
exchange of mole-
cules (p. 135) is still
going on between the
solid and the solu-
tion. That this con-
ception is correct
may be shown in various ways. Thus, if a crystal, the edges or cor-
ners of which have been, broken, is suspended in a saturated solution
of the same substance, it neither increases nor diminishes in weight.
Yet we find that the imperfections are removed, and that this takes
place by the solution of a portion of the substance from the perfect
surfaces and its deposition upon the imperfect ones.
Another very striking proof of this may be obtained by saturating
water with ordinary Glauber's salt (hydrated sodium sulphate, NagSOj,
10H2O) at a temperature somewhat above the ordinary, say 30°. The
excess of the solid is carefully and completely separated from the
liquid, and the latter is allowed to cool in a flask loosely stoppered with
cotton. The solution now contains (Fig. 62) a much larger amount of
o so
10° 20° 30°
40° 50° 60°
Temperature
FIG. 62.
70° 80° 90° 100°
SOLUTION 159
sodium sulphate (N"a2S04) than at its present temperature it could
acquire from contact with Glauber's salt. Yet in the absence of a
crystal, with which the above described exchange could take place, no
deposition of the dissolved substance begins. The solution may be
kept indefinitely without alteration. The introduction, however, of the
minutest fragment of the decahydrate at once starts the exchange, and
this is necessarily very much to the disadvantage of the solution and
the advantage of the crystal: Na,jS04 (diss'd) +± Na,,S04, 10H20 (solid).
The latter therefore forms the center of a radiating mass of blade-
like processes, which sprout with astonishing rapidity through the
liquid.
Usually the cooling of a concentrated solution leads to the almost
immediate appearance of crystals spontaneously, and the substance is
deposited gradually as the temperature falls. But solutions of a nun>
ber of common substances, such as sodium thiosulphate (photogra-
pher's " hypo ") and sodium chlorate, behave like that of sodium sulphate.
They are said to have a tendency to give supersaturated solutions.
In general, crystallization can be started only by introduction of a
specimen of the same substance, or at all events of one isomorphous
(<?.?;.) with it. The smallest particle of the right material floating in
the air, if it gains accidental admission, will bring about the result.
This shows the importance of the interchange of molecules of which
we have spoken for establishing equilibrium.
Metastdble Condition. — The above phenomenon is not an iso-
lated or exceptional one in physical science. It is commonly the
case that, when the conditions for some physical change have been
reached, the beginning of the physical change is delayed or entirely
fails. The system is then said to be in a metastable condition. Un-
stable it is not. Yet it is not in the state of greatest stability, for the
element of eqiiilibrium is lacking. Thus, pure water may easily be
cooled three or four degrees below 0° without the appearance of any
ice. Agitation, however, in this case, results in the appearance of ice
sooner or later. In like manner water may be heated to a temperature
above 100° without boiling. Drops of water, suspended in some oil of
almost the same specific gravity, may even be raised to 175° before the
water turns into steam. Similarly, air which is saturated with moisture,
if it contains no dust, may be cooled without the appearance of fog.
Phenomena of supersaturation, of a temporary kind at least, are extremely
common in chemistry. Almost every delayed precipitation is a case of it. Barium
160 INORGANIC CHEMISTRY
sulphate, for example, is always slow in appearing in dilute solutions. So is sulphur,
set free from dilute sodium thiosulphate solution by the action of an acid. In the
latter case, instant reneutralization with a base does not prevent the ultimate
appearance of the sulphur, showing that the cause does not lie in slow interaction
of the salt with the acid.
Saturation. — When we have shaken a solid for a sufficient length
of time with a given amount of a liquid, we obtain a solution which is
saturated with respect to that substance. Having called this a satu-
rated solution, we are inclined to extract from the term a meaning
different from that which it was really intended to convey. We are
in danger of thinking that the solution itself is in some way peculiar
— that, for instance, it contains all of the solid which it is capable of
holding. This would be an entire misconception. If we desire to
make a solution of sodium sulphate (Na^SC^), for example, we may
present this substance to the water either in the form of Glauber's
salt (NagSC^, 10H2O) or of anhydrous sodium sulphate. Now the
anhydrous and the hydrated forms of a substance always behave like
entirely different substances. This hydrate cannot give more than 5
parts of sodium sulphate (Na^SC^) to 100 parts of water at 0°. When
the anhydrous compound is used, many times this amount (Fig. 62) is
dissolved at the same temperature: N^SC^ (solid) <=»Na.jS04 (diss'd).
The solution pressures of the two forms are entirely different. The
phrase, " a saturated solution of sodium sulphate," is therefore devoid of
definite meaning. We must describe the liquid as a solution saturated
by anhydrous, or by hydrated sodium sulphate as the case may be.
Being different substances, the hydrated and anhydrous forms of a compound
must be investigated separately as to their solubility at various temperatures. The
results must give different curves, as for distinct substances.
Before referring to the curves of the two sodium sulphates, it must be remarked
that hydrates decompose into the anhydrous, or some less hydrated form at a defi-
nite temperature. We cannot therefore continue the observation of the solubility
of the substance beyond the temperature at which it ceases to exist. Thus the
solubility curve of a hydrate comes to an abrupt termination at the decomposition,
or, as it is usually called, the transition point. Now the decahydrate of sodium
sulphate decomposes at 32.4°, so that its solubilities can be measured only from about
0° to 32.4°. The solubility of the anhydrous form, however, can be investigated
up to 100°, or beyond it, if necessary. Not only so, but measurements can be
carried out below 32.4°. The union with water to form the decahydrate is so slow
that there is time to saturate the solution with the anhydrous body, and decant the
liquid for analysis, before the hydrate begins to be produced.
The solubility of the decahydrate (Fig. 62) rises rapidly between 0° and 32.4°
from 5 to 65 parts in 100. The solubility of the anhydrous sodium sulphate de-
SOLUTION 161
creases steadily from more than 55 parts below 32.4° to 42.5 parts at 100°. The
character of the two bodies, in the matter of solubility, is therefore entirely differ-
ent. The solutions themselves, as has been said already, are identical in every
way, when they have the same concentration, whether they have been made from
the one substance or the other.
Let us return now to the proper use of the term " saturated solution." We might
say, correctly, that at 20° a solution made from hydrated sodium sulphate and
containing 19.4 parts of sodium sulphate in 100 of water was saturated. This
would not, however, be the maximum quantity which the same amount of
water could hold, for, with the help of the anhydrous comp'ound, we could add an
amount equivalent to prolonging the ordinate at 20° until it intersected the curve
of anhydrous sodium sulphate somewhere about the value 60. Nor can we be
sure that even then the water would contain all the sodium sulphate which it
could hold. It is conceivable that, by presenting the substance in some still other
form, even greater solubility might be observed.
A saturated solution, if we fix our minds upon it simply as a solution, is not
different from any other solution. There is no feature in the properties of such
a solution qua solution which distinguishes it in the least from one containing
slightly less or one containing slightly more of the dissolved substance. In con-
tact with a crystal of the substance used to produce the saturation, however, the
saturated solution is found to be in equilibrium, while the unsaturated solution
takes up more of the substance, and the supersaturated solution at once deposits
the amount which it contains in excess of the saturated solution. The words un-
saturated, saturated, and supersaturated convey, therefore, no meaning unless we
add that the solution is so towards some specific form of material. These are quali-
ties of the system including the undissolved body and not of the solution by itself.
The fact that the hydrated and anhydrous sodium sulphates give saturated
solutions of the same concentration at 32.4°, so that the curves intersect at this
point, is a very significant one. This is the temperature at which the former sub-
stance turns into the latter. At transition points like this the values of solubility,
vapor pressure, and some other properties, are always the same for both forms (see
Freezing-points of solutions).
Properties of Solutions Proportional to Concentration :
Vapor Tension. — Besides osmotic pressure (p. 151), there are several
properties of solutions which are proportional to the concentration of
the solution.
If, instead of water, we introduce aqueous solutions of the same
substance successively into the barometric vacuum (Fig. 39, p. 116), we
find that the vapor pressures of the solutions are less than that of
water at the same temperature. The diminution in the fall of the
mercury column, which measxires the lowering in vapor pressure, is
proportional to the concentration of each solution. The limit is
reached with the saturated solution, although, if this is rather con-
centrated, the proportionality does not hold strictly down to that
point (see Chap. xvii).
162
INORGANIC CHEMISTRY
mm.
760
Temperature
FIG. 63.
This lowering in the vapor pressure of water is often considerable.
Thus at 100° a 7.5 per cent solution of potassium chloride shows a
vapor pressure of only 734.1 mm., while that of water is 760 mm. The
difference is 25.9 mm. Hence the solution has to be raised to a higher
temperature (100.96°) before it boils. This is almost exactly .037°
per 1 mm., which is the value for pure water (p. 118).
This conclusion may also
be reached graphically (Fig.
63). The ordinates represent
the vapor tensions correspond-
ing to the temperatures shown
by the abscissae. They in-
crease in length with rise in
temperature. The horizontal
dotted line shows the vapor
tension of 760 mm. at which
any liquid will boil. The
boiling-point of the solvent is
therefore the temperature at
which its curve intersects this
line. Since the vapor tensions
of the solution are all below
those of the solvent, its curve lies below that for the solvent but ascends along with
the latter. Hence it also cuts the 760 mm. line, but at a point beyond the boil-
ing-point of the pure solvent. Since, for short lengths, the curves are very nearly
straight lines, the distances from boiling-point to boiling-point are very nearly pro-
portional to the vertical distances between the curves. That is to say, the eleva-
tions in the boiling-point are proportional to the depressions in the vapor
tension, and therefore to the concentration of the different solutions of
any one substance. When different substances or solvents are compared, the
scale alone is different (see Chap. xvii).
The same effect may be observed in isomorphous mixtures of solid bodies,
which in many ways resemble solutions. Thus the vapor tension of water in the
alums is greater than the average vapor pressure of water in the air, and hence
they lose their water of hydration spontaneously (cf. p. 121). But if mixed crystals
of two alums, say ordinary alum and iron alum (K2SO4, A12(S04)3, 24H2O and
(NH4)2SO4, Fe2(SO4)s, 24H2O), are prepared, they keep perfectly. The vapor tension
of the water in each has been lowered by the influence of the other alum dissolved
in it. Similarly, calcium formate (Ca(CH02)2, 4H2O) effloresces (p. 123), but loses
this tendency when crystallized with some of the isomorphous barium or strontium
salt.
This principle explains the habit of very soluble compounds
(deliquescent substances) to become moist when exposed to the air
and finally to dissolve in the water they seem to attract from it.
Since, on account of their solubility, the moisture first present on
SOLUTION
163
their surfaces (p. 117) is a highly concentrated solution, its vapor ten-
sion is lower than the vapor pressure of water in ordinary air (p. 161).
Hence further condensation takes place (pp. 121-122) until the sub-
stance is so diluted that the vapor tension of its solution becomes equal
to the aqueous pressure of the atmosphere. (Concentrated sulphuric
acid, calcium chloride, and magnesium chloride (p. 34) are substances
of this kind.
Freezing-Points of Solutions. — Every pure liquid has a defi-
nite temperature at which it freezes. Thus, pure water freezes at 0°
and benzene at 6°. The presence of a foreign, dissolved body, how-
ever, lowers the freezing-point. Thus, sea-water is harder to freeze
than fresh water. The freezing-points of solutions of the same sub-
stance are found to be depressed below that of the solvent in pro-
portion to the concentration of the solute* (see Chap. xvii). This
may be shown graphi-
cally (Fig. 64). Theordi- mm.
nates represent vapor c
tensions corresponding •§
to the temperatures £
shown by the abscissas. o
The rate at which the £
former rise is greater for 4.6 — •*
ice than for water, hence
the ice curve is steeper.
At 0°, ice and water can Temp. o°
coexist permanently
(p. 115). By measure-
ment, they have the same vapor tension (4.6 mm.) at this point.
Theoretically, if they had not, they could not coexist indefinitely,
for the one with the greater vapor tension would evaporate, and its
vapor would condense on the other until one of them alone remained.
Now the vapor tension of a solution is, at all temperatures, lower
(Fig. 64) than that of water. Hence, for a solution, the curve must
cut the ice curve below 4.6 mm. and therefore behind 0°. In other
words, ice and the solution cannot have equal vapor tensions, and
therefore coexist indefinitely, except at some temperature below 0°.
* The ice which separates during the freezing consists, as a rule, of the pure
solvent, and the solute does not enter into it. Only when this is the case does this
law represent the facts.
FIG. 64.
164 INORGANIC CHEMISTRY
But the temperature at which ice can exist indefinitely in a solution
is the freezing-point. Hence, freezing-points of solutions are always
lower than those of the pure solvents. ,
By measuring the vapor tensions of the solution at several temper-
atures, in order to see how far the curve for the solution is below that
of water, and then producing the curve for the solution backwards, the
intersection with the ice curve, and therefore the freezing-point, may be
obtained graphically. Direct measurement always confirms the result.
At transition points like that of ice and water, the values of proper-
ties of both forms, such as the vapor pressure and the solubility in
some different substance, are always identical.
Since the ice curve is, for a short distance, almost a straight line, it
follows that the depressions of the freezing-point are proportional to
those of the vapor tension, and these in turn are proportional to the
concentrations. From this relation we get the statement with which
this section opened. Thus, solutions of sugar containing 11.4, 22.8, and
34.2. g. of sugar to 100 g. of water freeze at - 0.62°, - 1.23°, and - 1.85°,
respectively. Numerically, in the case of water, a lowering of the
vapor pressure by T£7 of its amount at each temperature sets the
freezing-point back 1.05°.
In everyday life we scatter salt on ice to melt it. The salt dis-
solves in the moisture on the surface, and the ice cannot exist in pres-
ence of this solution at 0°. It melts, absorbing heat in doing so, until
the temperature of the mixture reaches that of a freezing, saturated
solution of salt (about — 21° = — 6° F.). When the existing temperature
is lower than this, the salt has no effect on the ice. Freezing mixtures,
being always mixtures of ice with various soluble substances, work
in accordance with the principle.
Densities of Solutions. — The densities or specific gravities of solutions
are also functions, although not simple ones, of the concentrations, and hence the
latter are commonly defined for commercial purposes by the former. We purchase
ammonium hydrate of "0.88 sp. gr.," meaning 35 per cent of ammonia, or sul-
phuric acid of " 1.84 sp. gr.," meaning 100 per cent of the acid.
The solution usually has a smaller volume than that of the sum of the constit-
uents. Thus, 58.5 g. of sodium chloride occupying 27.5 c.c., give, when dissolved
in 10 liters of water, a solution whose volume is 10.0166 liters. The solution is only
16.6 c.c. more bulky than the water. Similarly, a formula-weight of potassium
nitrate, occupying 44.7 c.c., adds only 38.5 c.c. to the bulk of 10 liters of water.
Heat of Solution. — When a body is dissolved in water, the solu-
tion may be either warmer or colder than the original materials. The
SOLUTION 165
sources or destiny of the heat given out or absorbed have not been
studied in such a way that definite statements can be made about the
theory of the subject. There are many factors which would have to be
considered. For example, the body, if a solid, goes into an essentially
liquid condition, and its heat of fusion is always negative. The changes
in its molecular condition involve either liberation or absorption of
heat. The change in volume must have a heat effect connected with
it. But almost nothing is known about these and other (see Heat
of ionization) essential parts of the phenomenon.
The first water used always causes a greater heat change than the ad-
dition of succeeding equal amounts. Heats of solution are measured for
the solution of one formula- weight of the substance in unlimited water.
The values in calories for some common substances are as follows :
(H2S04, Aq) = + 39,170 (NaCl, Aq) = - 1180
(HC1, Aq) = + 17,400 QStRfR, Aq) = - 3880
(KOH, Aq) = + 12,500 (KC1, Aq) = - 4440
(NaOH, Aq) = + 9780 (Na2C03, 10H20, Aq) = - 16,160
(Ka^COg, Aq) = + 5640 (Na^SO,, Aq) = + 460
(CaCL,, Aq) = + 3258 (Na^SO,, 10H20, Aq) = - 18,760
When a substance comes out of solution, the heat effect is equal and of
opposite sign to that occurring when the same substance goes into solu-
tion. Hence, since the decahydrate of sodium sulphate absorbs heat
in dissolving, a considerable development of heat is noticed when it
suddenly crystallizes from a supersaturated solution. Some ether in a
tube immersed in the solution may be boiled by this heat and its vapor
set on fire to make the fact evident at a distance. An important rela-
tion between heat of solution and solubility will be discussed under
van't Hoff's law of mobile equilibrium (</.v.).
Definition of a Solution. — We are now able to make a brief
statement which shall distinguish solutions from mixtures on the one
hand and from chemical compounds on the other. Solutions are
homogeneous mixtures of two or more substances -which are not sep-
arable into their constituents by mechanical means •without altering the
state of one of the substances, and -whose properties vary continuously
•with the proportions of the constituents bet-ween certain limits.
Application in Chemical Work- — The theory of this subject
has been given on account of its intensely practical interest, and it
166 INORGANIC CHEMISTRY
should be kept in mind in all ordinary chemical operations. It will
afford an explanation of many things which might otherwise be attrib-
uted to the wrong cause, or might remain entirely without explanation.
For example, why is the action of a metal upon an acid so slow ? We
must remember that an acid diluted with water is being used, and
only one molecule out of every dozen or hundred is a molecule of the
acid. So that the access of the latter to the metal is restricted at first,
and becomes more and more so as the molecules of the acid in the
immediate vicinity of the metal undergo chemical change.* On the
other hand, the metal, especially if it be in the form of sticks obtained
by casting, presents one of the elements in the action in a most com-
pact form. The only parts which are accessible to the acid are those
upon the surface, and, the metal not'being appreciably soluble in water,
the molecules can only pass off and expose a fresh layer very slowly.
It is no wonder that many chemical actions occupy a considerable time.
The wonder is that they should take place as rapidly as they do.
Their speed would seem to point to a most intense chemical activity,
even in the seemingly feebler instances. Various artifices are habitu-
ally employed for facilitating chemical action. Thus, the metal may be
reduced to a leafy form by pouring the molten substance into cold
water. Naturally, with metals, the maximum surface and the most
rapid chemical action are obtained by using a fine powder.
The most speedy interaction of all, other things being equal, must
be attainable by dissolving all the interacting substances in water.
Under these circumstances all the molecules of each substance must
simultaneously have many molecules of the others within easy reach of
them.
Exercises. — 1. Give other examples of limited solubility in vari-
ous solvents (p. 146).
2. If you were not permitted to evaporate sea-water to dryness, how
should you show that it was a solution and not a pure substance ?
3. Eeexpress Henry's law (p. 154) in terms of the volume of gas
dissolved at different pressures.
4. If hydrogen sulphide is diluted with ten times its volume of
hydrogen, what volume of it, estimated as pure gas, will be dissolved
by 20 volumes of alcohol at 0° and 760 mm. (p. 154) ?
* In spite of the continuous exhaustion of the acid, there is often a steady in-
crease in the rate at which a dilute acid interacts with a metal. This is due, at
first, to the dirt on the surface of the metal which temporarily obstructs the ac-
tion, and, later, to the rising temperature of the interacting bodies.
SOLUTION 167
5. If the dissolved air, after being removed from water by boiling,
were to be shaken with water once more, in what proportions by
volume would the gases now dissolve (p. 155) ?
6. Read from the curves (p. 157) the solubilities of potassium
nitrate at 15°, of potassium chloride at 30°, of potassium chlorate at
45°. What are the relative rates at which the solubilities of these
salts increase with rise in temperature ?
7. Express the concentrations of solutions of ammonium chloride,
saturated at 0° (sp. gr. 1.076), and of potassium sulphate K2S()4, satu-
rated at 10° (sp. gr. 1.083), in terms of a normal solution (p. 149).
8. Express the concentration of a five per cent aqueous solution
of phosphoric acid (sp. gr. 1.027), in terms of a normal and a molar
solution, respectively.
9. Name the phases (p. 156) in a system consisting of oxygen and
its aqueous solution, (a) above 0°, (b) below 0°.
10. When a solution of a very soluble substance, like zinc chloride,
is evaporated to dryness on a water bath, why is the escape of the
last portions of the solvent so much slower than is that of the first ?
CHAPTER XI
CHLORINE AND HYDROGEN CHLORIDE
CHLOBINE was first recognized as a distinct substance by Scheele
(1774). He obtained it from salt by means of manganese dioxide,
using the common method described below. It was for years sup-
posed to be a compound containing oxygen, but the work of Davy
(1709-1818) established the fact that it is an element.
Occurrence. — Chlorine does not occur free in nature. There are,
however, many compounds of it to be found in the mineral kingdom.
Sea- water contains a number of chlorides in solution. Nearly 2.8 of
the 3.6 per cent of solid matter in sea-water is common salt (sodium
chloride, Nad). During past geological ages the evaporation of sea-
water has led to the formation of immense deposits of the compounds
usually found in such water. Thus, at Stassfurt, such strata attain a
thickness of over a thousand feet. Certain layers of these strata are
composed mainly of sodium chloride, called by the mineralogist halite
(rock salt). In other layers potassium chloride (sylvite), and hydrated
magnesium chloride (bischofite), and other compounds of chlorine, oc-
cur. The chloride of silver (horn silver) is a valuable ore.
Preparation. — Chlorine cannot be obtained with the same ease
as oxygen. There are only a few chlorides, such as those of gold and
platinum, which lose chlorine when heated, and they are too expensive
or difficult to make for laboratory use. We employ therefore methods
like those used for the preparation of hydrogen. We may (1) decom-
pose any chloride by means of. electricity, just as, to get hydrogen, we
electrolyzed a dilute acid (pp. 63, 94). Or (2) we may take some inex-
pensive compound of chlorine, such as hydrogen chloride (HC1), and by
means of some simple substance which is capable of uniting with the
other constituent, — here oxygen serves the purpose, — secure the
liberation of the element (p. 99). Or (3)- — and this turns out to be
the most convenient laboratory method — we may use a more complex
action.
168
CHLORINE AND HYDROGEN CHLORIDE
169
Electrolysis of Chlorides. — Hydrogen chloride and those chlo-
rides of metals which are soluble in water are all decomposed when a
current of electricity is passed through the aqueous solution. They
yield chlorine at the positive electrode. The other constituent, the
hydrogen (Fig. 65), manganese, or whatever it may be, is liberated at
the negative wire. To decompose hydrochloric acid an electromotive
force of at least 1.31 volts is required. Since the chlorine is soluble in
water, the efferves-
cence due to its re-
lease is not notice-
able until the liquid
round the electrode
has become satu-
rated with the gas :
CLj (diss'd) <r± CL,
(gas). The shape
of the apparatus
keeps the two prod-
ucts from mingling.
The presence of the
chlorine in the liq-
uid at the positive
end may be shown
by a suitable test
(pp. 99 and 175).
In commerce
chlorine is now ob-
tained chiefly by
this method, sodium
chloride or potassium chloride being the source of the element. Elec-
trodes of artificial graphite are used, as most other conductors unite with
the chlorine. The potassium or sodium, as the case may be, travels
towards the negative electrode, but is not liberated (p. 99). Instead,
potassium or sodium hydroxide accumulates in the solution round the
plate and hydrogen escapes. The chlorine is released at the positive
electrode, as usual. The hydroxide and the chlorine both find chemi-
cal applications. The chlorine is either liquefied by compression in
iron cylinders or employed at once for making bleaching powder (q.v.~).
In the Acker process (q.v.), melted, in pure sodium chloride, with-
out any solvent, is electrolyzed,
FIG. 65.
170 INORGANIC CHEMISTRY
Action of Free Oxygen on Chlorides. — Oxygen does not inter-
act with sodium chloride even at a high temperature. Hence the
chlorine must first be transferred to some other form of combination.
By the interaction of acids, of which the most commonly used is sul-
phuric acid, with chlorides, of which the cheapest is sodium chloride,
we obtain hydrogen chloride. The details of this action are described
below (p. 178). In order to liberate chlorine from this compound, we
may combine the hydrogen with oxygen obtained from the air. The
action is in accordance with the equation :
2HC1 + 0 -> H20 + 2C1.
The two gases interact so slowly, however, that a catalytic agent must
be employed. The mixture of air and hydrogen chloride is passed
over pieces of heated pumice-stone or broken brick previously saturated
with cupric chloride solution. A temperature of 370°-400° is used.
In the resulting gas the chlorine is mixed with steam and with a very
large volume of nitrogen which entered with the oxygen, so that for
making the pure substance this method (Deacon's process) is quite un-
suitable.
Using the same principle, magnesium chloride may be heated in a stream of air,
when the oxide of magnesium is formed and chlorine is given off: MgCl2 + O — »
MgO + 2C1. The oxide of magnesium can then be treated with hydrochloric
acid to regenerate the chloride, which in turn may be subjected once more to the
action of oxygen. The process is thus a continuous one.
The above action is spoken of as an oxidation. It is true that no
oxygen is actually introduced into the hydrogen chloride as a whole.
The removal of hydrogen from combination with the chlorine is, how-
ever, the first step towards the introduction of oxygen into combination
with the latter, and is essentially an oxidation.
Action of Combined Oxygen upon Chlorides. — The usual
laboratory method of making chlorine is to mix a solution of hydrogen
chloride in water (hydrochloric acid, p. 93) with roughly powdered man-
ganese dioxide. The flask containing the mixture is heated by means
of a bath containing hot water (Fig. 66). The gas is passed through
a washing bottle containing water, in order to remove some hydro-
gen chloride which may be carried over. It may be dried, if neces-
sary, in a second washing bottle containing concentrated sulphuric
acid. It cannot be collected over water on account of its solubility, so
that jars are usually filled with it by downward displacement of air.
CHLORINE AND HYDROGEN CHLORIDE
171
The chemical change used here is somewhat complex. When an
acid (here HC1) interacts with an oxide (here MnO2), the hydrogen of
the former unites with the oxygen of the latter, giving water. We
perceive at once that to combine with 20, 4H, obtainable only by
taking 4HC1, will be required. Hence the equation might be :
Mn02 + 4HC1 _» 2H20 + MnCl4.
This is probably what happens in the first place. The products
actually obtained, however, are water, manganous chloride (MnCLj) and
FIG.
chlorine. The manganese tetrachloride is decomposed by the heating,
the chlorine escapes, and the other two products remain in the vessel.
MnO2 + 4HC1 -» 2H20 + MnCL, + 2C1. (1)
We owe the chlorine to the fact that the tetrachloride is unstable.
When the mixture is surrounded by ice and saturated with chlorine, it can be
shown with some degree of certainty that it contains the tetrachloride. If it is
quickly poured into water, hydrated manganese dioxide is precipitated (Wacker).
The decomposition of the tetrachloride is reversible :
MnCl4 ?z> MnCl2 + CL,,
and is driven back by the excess of chlorine. The tetrachloride is hydrolyzed by
water :
MnCl, + 2H2O + xII20 — > MnO2, zII,O
172 INORGANIC CHEMISTRY
The action (1) is of a type very common in chemistry. It is more
complex even than double decomposition (p. 99), and, unlike this, its
results cannot be anticipated by guessing. If we had used manganous
oxide (MnO), we should have had a double decomposition :
MnO + 2HC1 il» H20 + MnCL,, (2)
but we should have got no chlorine. Perhaps the simplest way to de-
scribe the difference between these two actions is in terms of the val-
ence of the manganese. In MnIV02u the element is quadrivalent.
This means that its atomic weight professes to be able to hold four
unit weights of a univalent element. The four valences of oxygen
(20 n) can do the same thing. In equation (1) the oxygen fulfils this
promise by taking 4H.1 But the MnIV can hold only 2C11 permanently
and lets the other 2G11 go free. In other words, the valence of the
unit weight of manganese changes in the course of the action. In
equation (2), on the other hand, the manganese is bivalent to start
with (Mn1^11), and is able to retain the amount of chlorine (2C11)
equivalent to O". Actions like that of manganese dioxide in (1) are
classed as oxidations. The hydrogen chloride, or rather half of it, is
oxidized. A graphic mode of writing may make this remark clearer :
2HC1-»H90 + MnJI012
2HC1^ HJO + 2C1.
The upper half is a double decomposition, the lower an oxidation by
half the combined oxygen of the dioxide.
In practice, instead of employing aqueous hydrochloric acid we
frequently use the materials from which it is prepared, namely, common
salt and concentrated sulphuric acid (p. 178), along with the manganese
dioxide. Under those circumstances, the action appears more com-
plex, but is simply a combination of the two chemical changes, and is
represented by the equation :
Mn02 + 2NaCl + 3H2S04 -» 2H20 + 2NaHS04 + MnS04 + 2C1.
Using the same principle, we find that lead dioxide, barium dioxide,
potassium dichromate (K2Cr207), potassium permanganate (KMn04),
and many other substances, when treated with aqueous hydrochloric
acid, likewise give chlorine.
Perhaps the best way to obtain a steady, easily regulated stream of
chlorine is to place some solid potassium permanganate in a flask,
arranged like that in Fig. 70 (p. 232) but without the U -tubes, and
CHLORINE AM) HYDROGEN CHLORIDE 173
allow concentrated commercial hydrochloric acid, diluted with half its
volume of water, to fall upon it drop by drop from the dropping fun-
nel. The action is very rapid, the acid is exhausted almost as fast as
it falls, and so the stream of gas can be stopped by simply closing the
stopcock. In this, and in all the cases mentioned above, the chemical
action follows the same plan as before (p. 171). The oxygen and hy-
drogen combine to form water, hence here 8HC1 will be needed ; the
metals yield the chlorides which are stable at the temperature of the
action, and the rest of the chlorine is liberated :
KMn04 + 8HC1 -» 4H20 -|- KC1 + MnCL, -f 5C1.
Kinetic-Molecular Hypothesis Applied to these Actions. — In
preparing chlorine by the usual laboratory method (p. 170), it will be
observed that the gas is produced rather slowly. Even heating does
not accelerate the interaction very greatly. The situation is that we
have placed together manganese dioxide in a granular form and water
which contains hydrogen chloride in solution. The dioxide is very
insoluble in water, and consequently its molecules, which must dissolve
before they can meet the acid, become available very slowly : Mn02
(solid) <=± Mn02 (diss'd). The finer the pulverization, the less will be
the delay from this cause. On the other hand, the acid contains
originally only about one molecule of hydrogen chloride for every four
of water, and as the former is used up the scarcity of the active sub-
stance becomes greater.
Again, we heat the mixture on a water bath so as to hasten the
process (p. 72) by raising the temperature to about 90°. When we
prepared oxygen, we forced the temperature up with a naked Bunsen
flame until, between 200° and 300°, a sufficiently rapid stream of the
gas was secured. The iron and sulphur (p. 11) we raised nearly to a
red heat. Here the conditions make stronger heating impossible. No
aqueous solution of hydrogen chloride can be raised above 110°, the
maximum boiling-point (p. 182). But we must not carry the heating
so far as 110°, because even below this point the concentrated acid gives
off gaseous hydrogen chloride freely. If we did, we should contaminate
our chlorine and at the same time lose a part of one of the ingredients on
which the action depends. Intelligent chemical work always demands
a careful consideration of purely physical facts of this description.
Physical Properties. — Chlorine differs from the gases we have
encountered so far in having a strong greenish-yellow tint (Gk.
174 INORGANIC CHEMISTRY
os, pale green), a fact which gave rise to its name, and having a
powerful irritating effect upon the membranes of the nose and throat.
Density (H = 1), 35.79 Boiling-point (liq.), - 33.6°
Weight of 1 1., 3.220 g. Melting-point (solid), - 102°
Solubility in Aq. (20°), 216 vols. in 100 Vap. tension (liq.) 0°, 3.66 atmos.
Crit. temp., + 146° Vap. tension (liq.) 20°, 6.62 atinos.
Since a liter of air weighs 1.293 g., chlorine is two and a half
times heavier. In solubility it stands between slightly soluble gases,
like oxygen and hydrogen, and those which are extremely soluble. It
can be collected over hot water or a strong solution of salt.
It was first liquefied by Faraday. The critical temperature (p.
133) is exceptionally high (146°), so that at all ordinary temperatures
the gas can be liquefied by compression alone. It forms a yellow
liquid which, contained in steel cylinders, is now an article of com-
merce. On being cooled below — 102°, it gives a pale-yellow solid.
Chemical Properties. — Chlorine is at least as active a sub-
stance as is oxygen. It presents a more varied array of chemical
properties than does that element.
1. When a strong solution of chlorine in water is cooled with ice,
crystals of chlorine hydrate (Cl, 4H20) are formed. This hydrate,
being the hydrate of a gas, has a high vapor tension of chlorine, as
well as a smaller one of water (p. 121). At 0° the partial tension of
the chlorine is 249 mm., at 9.6° it is 760 mm. The compound decom-
poses rapidly, therefore, at the latter temperature unless the pressure
of chlorine over it is greater than one atmosphere.
2. Chlorine unites directly with many elements. A jet of hydrogen
burns vigorously in chlorine, producing hydrogen chloride. The
presence of this product may be recognized at once, because, while
chlorine in contact with moist breath gives no cloud, hydrogen
chloride (q.v.) produces a dense fog. The union of the gases, when
a mixture of them is kept cold and in the dark, is too slow to be per-
ceived. On exposure to diffused light, however, they unite slowly,
while a sudden flash of sunlight or the burning of a magnesium ribbon
causes instant explosion. The function of the light here is entirely
different from that in the decomposition of silver chloride (p. 19).
In the latter case light was used to maintain the change, which comes to
a stop whenever the light is withdrawn. The action was endothermal
and consumed energy. The union of hydrogen and chlorine is highly
exothermal, and a minimum of light only is needed to start it (p. 74).
CHLORINE AND HYDROGEN CHLORIDE 175
Sodium burns in chlorine, producing a cloud of white particles of
sodium chloride. Copper in the condition of thin leaf commonly
used for gilding (Dutch-metal), catches fire spontaneously when thrust
into the gas. Phosphorus burns in it with a rather feeble light, pro-
ducing phosphorus trichloride, a liquid (b.-p. 74°) which condenses
readily. The proportions of the materials taking part in this change
show that one atomic weight of phosphorus (31) and three atomic
weights of chlorine form the product :
p _|_ 3d _» PCI,.
Almost all the more familiar elements unite directly with chlorine to
form chlorides. The exceptions are, nitrogen, oxygen, carbon, helium,
and argon (p. 68). Compounds with all but the last two may be ob-
tained as products of more complex interactions, however.
Some elements form more than one chloride. Thus, when phos-
phorus unites at a high temperature with chlorine, we get phosphorus
trichloride. When we pass chlorine gas into cooled phosphorus tri-
chloride, however, a considerable absorption takes place, and finally
a solid body containing two more atomic weights of chlorine, phos-
phorus pentachloride (PC15), is formed.
Carefully dried substances unite slowly or, to all appearances, not
at all with chlorine. Thus, perfectly dry chlorine does not unite with
copper, even when the system is warmed. The introduction of a drop
of water (p. 75), however, into a remote part of the apparatus at
once supplies the trace of moisture which seems to be necessary to
start the chemical change. Chlorine, as ordinarily made, unites with
iron with great vigor/ But dried chlorine is quite indifferent and
does not attack the steel cylinders in which it is now sold. The water
is a catalytic agent, and we regard it as simply hastening an action
which otherwise is vanishingly slow (p. 73).
3. Chlorine, like sodium (p. 97), may also displace elements which
are already in combination. Thus, when turpentine (C10H16) is poured
over a strip of paper and is then immersed in a jar of chlorine, a
violent action takes place, and an immense cloud of finely divided
carbon bursts forth. The heat of the action, as it starts, vaporizes
the turpentine, and the hydrogen in the latter unites with the chlorine
forming hydrogen chloride, while the carbon is set free : (C10H16 +
16C1 -» 16HC1 + IOC).
The action of chlorine on potassium iodide, dry or in solution, is of
this kind, and furnishes the commonest test (p. 99) for free chlorine.
176 INORGANIC CHEMISTRY
The chlorine simply takes the place of the iodine (KI + Cl —> KC1 + I),
and the latter is liberated. Iodine when moist is deep brown in color,
and the amount liberated by a large quantity of chlorine may easily
be seen. Yet, it is an advantage so to arrange a test that it may be as
delicate as possible, that is, may give a plainly visible result with the
minimum of material. In this case much starch emulsion and a little
potassium iodide are employed. Strips of filter paper dipped in this
mixture show a deep-blue color (see Iodine) when brought into a gas
containing even a trace of free chlorine. Combined chlorine, as in a
chloride, has no effect.
4. When actions like the above are moderated by proper means,
the decomposition is not so complete. If methane (marsh gas, CH4) is
mixed witb. chlorine and exposed to sunlight, a slower action occurs,
of which the first stage consists in the removal of one unit weight of
hydrogen and the substitution of chlorine for it according to the
following equation :
CH4 + 2C1 -> CH8C1 + HC1.
The process may continue further by the substitution* of chlorine for
the units of hydrogen one by one until carbon tetrachloride (CC14) is
finally formed.
A most interesting and important action of this class occurs when
chlorine is dissolved in water. A small part of the chlorine interacts
with a little of the water, the element being substituted for one-half
of the hydrogen :
2C1 + HHO -* HOC1 + HC1. (1)
Only traces of hydrogen chloride and hypochlorous acid are produced,
however. The change comes quickly to a standstill, because the prod-
ucts interact even more vigorously to reproduce chlorine and water :
HC1 + HOC1 «=» 2C1 + H20, (2)
the interaction being reversible (p. 64). This interaction of chlorine
and water *(1), slight as it is, is of importance on account of the insta-
* Substitution resembles displacement (p. 99) in that an element and a com-
pound interact, and the element takes the place of one unit in the composition, of
the latter. In the above action, one unit of chlorine takes the place of one unit
of hydrogen. But the latter is not liberated : it combines with another unit of
chlorine. Double decomposition the action is not, because elements do not decom-
pose. The name used is intended to fix the attention on the compound and on the
fact that one unit has been substituted for another in it. This conception is a
favorite one in the chemistry of compounds of carbon.
CHLORINE AND HYDROGEN CHLORIDE 177
bility of the kypochlorous acid (q.v.) which it produces. When the
solution is exposed to sunlight, the hypochlorous acid decomposes and
oxygen gas is produced : HC10 — > HC1 + 0. Since this removes the
substance on whose interaction with the hydrogen chloride in (2), the
reversal of (1) depends, the latter action proceeds under continuous
illumination gradually to completion. Hence the aqueous solution of
chlorine must be kept in the dark, since otherwise a dilute solution
of hydrogen chloride alone remains.
The so-called bleaching action of "chlorine" is almost always the result of
oxidation of the coloring matter by hypochlorous acid. Chlorine and the dye in
the cloth, even when only moderately dry, show no tendency to interaction. This
may be demonstrated by collecting some chlorine in a stoppered bottle in the
bottom of which a little concentrated sulphuric acid stands. A piece of colored
calico may be attached by a pin to a cork stuck in the bottom of the stopper and
so suspended in the gas. After twenty-four hours no action will be found to have
occurred. Yet if the rag is first moistened, the bleaching is almost instantaneous.
Hence this bleaching action is treated under the properties of hypochlorous acid.
5. Chlorine may simply add itself to a compound. Thus, one of the
oxides of carbon, carbon monoxide (CO), when mixed with chlorine
and exposed to sunlight gives drops of a volatile liquid (b.-p. 8.2°)
known as phosgene (COCLj).
Chemical Relations of the Element.* — In the formation of
chlorides, an atomic weight of chlorine is equivalent to one atomic
weight of hydrogen or of sodium. The element is, therefore, univalent
(p. 103). It never shows any higher valence than this save in its oxygen
compounds (see Chap. xvi). The oxides of chlorine interact with
water to give acids, and the element is, therefore, to be classed as a
non-metal (p. 119). It belongs to that group of the non-metals called
the halogens (q.v.\ as a consideration of some others of its relations
will show (see Chap. xiv).
Uses of Chlorine. — Large quantities of chlorine are manu-
factured for the preparation of bleaching materials and disinfecting
agents.
* In accordance with the distinction that must be drawn (p. 32) between the ele-
ment as a variety of matter in combination, and the elementary substance or free
form of the element, and to avoid a common source of confusion, we shall always
give only the behavior of the elementary substance under the title chemical proper-
ties. The characteristics which distinguish the compounds of the element, as a
class, from, or relate them as a class to the compounds of other elements will then
appear in a separate section under the above title (see Chap, xiv, first section).
178 INORGANIC CHEMISTRY
In disinfection, the minute germs of disease and putrefaction are
acted upon either by the chlorine or by the hypochlorous acid formed
by its interaction with water, and instantly their life is destroyed.
One of the processes for the extraction of gold involves an action of
chlorine gas upon the material after it has received preliminary treat-
ment. A chloride of gold is formed, which can be dissolved out of
the matrix by means of water, and the metallic gold is afterwards
precipitated from the solution.
HYDROGEN CHLORIDE.
Preparation from Sodium Chloride' — As we have seen, the
direct union of hydrogen and chlorine produces a gas, hydrogen chlo-
ride (HC1). For the purpose of preparing this gas, however, some
more easily managed method will naturally be employed.
In commerce large quantities of hydrogen chloride are obtained as
a by-product in connection with the manufacture of soda. The same
materials are commonly employed in the laboratory. Common salt is
treated with concentrated sulphuric acid at a gentle heat. Efferves-
cence is seen and hydrogen chloride is given off, while a compound
known as sodium hydrogen sulphate (sodium bisulphate) remains
behind. The action is represented by the equation :
NaCl + H2S04 -» HC1 f + NaHS04. (1)
The apparatus used in the preparation of chlorine (Fig. 66) may be
employed. On account of its extreme solubility the gas is not washed
in water, however. For the same reason it must be collected by down-
ward displacement of air, or over mercury.
The above statements apply to the action of a large amount of sul-
phuric acid upon a limited amount of salt, where no high temperature
is employed. If, however, a larger proportion of salt is used and a
sufficiently high temperature produced by artificial heating, then com-
mon sodium sulphate is formed according to the equation :
2NaCl + H2S04 -, 2HC1 + Na2S04. (2)
The former action is that which occurs under the conditions used in
the laboratory. The latter is the action which is employed in com-
merce, since it is for the purpose of making sodium sulphate (Na^SO.,),
from which sodium carbonate is afterwards to be prepared, that the
operation is undertaken. The hydrogen chloride passes through a
tower, down which water trickles over lumps of coke, and is dissolved.
CHLORINE AND HYDROGEN CHLORIDE 179
The aqueous solution is called hydrochloric acid, or, in commerce,
muriatic acid.
Interaction of Acids and Chlorides. — It should be noted that
the above is simply an illustration of a perfectly general method.
Almost any chloride of a metal might have been used instead of sodium
chloride, and almost any acid which could be obtained free from any
large amount of water (see below) could have taken the place of sul-
phuric acid. Thus, concentrated phosphoric acid with any chloride
will give a change parallel to the above. With sodium chloride this
action would be NaCl + H8P04 — > HC1 \ -f- NaH2P04 (primary sodium
phosphate).
If various chlorides are used with the same acid, it will be found
that the vigor of the actions is very different. In some, hydrogen
chloride will be produced copiously without the assistance of heat.
In others, there will be difficulty in showing that hydrogen chloride
gas is produced at all. We must not hastily assume that this is owing
to any greater chemical affinity in one case than another. More ex-
tensive experimentation will show that the more soluble chlorides as a
rule give more vigorous effects than those which are less so (pp. 166, 173).
Ammonium chloride with sulphuric acid would represent the former
variety, while mercuric chloride with the same acid would represent
the latter.
The Kinetic Hypothesis Applied to the Interaction of Sul-
phuric Acid and Salt. — One who has used the above method for mak-
ing hydrogen chloride without reflection would not realize the complex-
ity of the machinery by which the result is achieved. The means are
apparently very simple. Yet the mechanical features of this experi-
ment, when laid bare, are extremely curious and interesting. A single
fact will show the possibilities which are concealed in it.
If we take a saturated solution of sodium hydrogen sulphate in
water and add to it a concentrated solution of hydrogen chloride in
water (concentrated hydrochloric acid), we shall perceive at once the
formation of a copious precipitate. This is composed entirely of
minute cubes of sodium chloride :
NaHS04 + HC1 -* H2S04 + NaCl |. (3)
Now this action is nothing less than the precise reverse of ^(1), yet it
proceeds with equal success. In fact, this chemical interaction is not
only reversible (p. 64), but can be carried to completion in either direc-
180 INORGANIC CHEMISTRY
tion. It is only in presence of a large amount of water that it stops
midway in its career and is valueless for securing a complete trans-
formation in either direction :
NaHS04 + HC1 <z± H2S04 + NaCl.
In an action which is reversible, if the products remain as perfectly
mixed and accessible to each other as were the initial substances, their
interaction will continually undo a part of the work of the forward
direction of the change. Hence, in such a case the reaction must, and
does, come to a standstill while as yet only partly accomplished
(cf. p. 176) ; but this was not the case with actions (1) and (3). Let
us examine the means by which the premature cessation of each was
avoided.
In (1) the salt dissolved to some extent in the sulphuric acid,
NaCl (solid) <=> NaCl (diss'd), and so, by contact of the two kinds of
molecules, the products were formed. On the other hand, the hydrogen
chloride, being insoluble in sulphuric acid, escaped as fast as it was
formed : HC1 (diss'd) <=± HC1 (gas). Hence, in that case, almost no
reverse action was possible, and the double decomposition went on to
completion. With all the sodium hydrogen sulphate in the bottom of
the flask, and most of the hydrogen chloride in the space above, the two
products might as well have been in separate vessels so far as any
efficient interaction was concerned. This plan, in which water is pur-
posely excluded, forms therefore the method of making hydrogen
chloride.
In (3), on the other hand, the hydrogen chloride was taken in
aqueous solution, and was kept permanently in full contact with the so-
dium bisulphate. It had therefore in this case every opportunity to
interact with the latter and no chance of escape. Every molecule of
each ingredient could reach every molecule of the other with equal
ease. Furthermore, the sodium chloride produced as a result of their
activity is not very soluble in concentrated hydrochloric acid (far
less so than in water), and so it came out as a precipitate : NaCl
(diss'd) +± NaCl (solid). But this was almost the same as if it had
gone off as a gas. It meant that the greater part of the salt was in the
solid form. It was in a state of fine powder, it is true. But, in the
molecular point of view, the smallest particle of a powder contains
millions of molecules, and most of these are necessarily buried in the
interior. Thus the sodium chloride was no longer able to interact
effectively molecule, to molecule with the other product, the sulphuric
CHLORINE AND HYDROGEN CHLORIDE 181
acid. Hence, there was little reverse action to impede the progress of
the primary one. Thus (3) is nearly as perfect a way of liberating sul-
phuric acid as (1) is of liberating hydrogen chloride.
This discussion is given to show that, in many chemical actions,
the affinity is entirely subordinated by the effects of a purely mechan-
ical arrangement (cf. pp. 28, 110, 166). If the latter is well devised,
an action propelled by a feeble affinity may prevail against a reverse
action involving a very powerful one (see Chemical equilibrium).
The egregious misconception that sulphuric acid is shown by this action to be
"stronger" than hydrochloric acid was disposed of, so far as the science was con-
cerned, half a century ago. But it survives in suburban chemical circles with re-
markable tenacity. The fact, quaintly enough, is that the real relation in respect
to activity is just the reverse.
Other Ways of Obtaining Hydrogen Chloride. — Although
never used for generating hydrogen chloride on a large scale, there is
another important kind of action in which the substance is a product.
When water acts upon the chlorides of non-metallic substances like
sulphur, phosphorus, and iodine, a double decomposition occurs. Since
water is always one of the interacting substances, this kind of change,
— a double decomposition involving -water, — is called hydrolysis
(Gk. v8<ap, water, and Awns, the act of loosing). Thus, when a little
water is added to one of the chlorides of phosphorus, hydrogen chloride
is formed. Besides this, the trichloride gives phosphorous acid, and
the pentachloride, phosphoric acid :
-> 3HC1 + P(OH),,
5HC1 + H8P04.
A dissociation is a reversible decomposition of one substance into two. Hy-
drolysis is an ordinary double decomposition or metathesis where water is one of
the reagents. Yet it has been perversely named hydrolytic dissociation by
many writers. A whole chapter might be devoted to the ingenuity with which
chemists have misnamed many of the things with which they deal. Perhaps this
tendency is a survival of the habit the chemists had of using obscure and symbolical
names for their materials to prevent the penetration of their secrets by uninitiated
seekers after knowledge. Important facts and principles have been sedulously
labeled with misleading titles, like : Water of crystallization, which has no more to
do with crystallization than with color, density, or any other physical property ;
supersaturated solution, which, as a solution, is the same as any other ; mass ac-
tion, which has nothing to do with mass, but is concerned wholly with concentra-
tion ; strong acid, which refers to activity and not power of resistance ; reciprocal
proportions, a law in which reciprocals of numbers play no part ; downward dis-
placement of air, when the air is displaced upwards, and so forth. Here there
182 INORGANIC CHEMISTRY
is an opportunity to confuse hydrolysis with electrolytic dissociation, and the be-
ginner never fails to embrace it. Hydrolytic double decomposition would have
been a correct if somewhat clumsy term.
Often, when a steady stream of hydrogen chloride is required, con-
centrated hydrochloric acid is placed in a generating flask, and concen-
trated sulphuric acid is allowed to trickle into it from a dropping
funnel. The hydrogen chloride is less soluble in diluted sulphuric
acid than in water (see Product of solubility) and escapes.
Physical Properties. — Hydrogen chloride is a colorless gas,
which produces a suffocating effect when breathed.
Density (H = 1), 18.23 Crit. temp., + 52°
Weight of 1 1., 1.641 g. Boiling-point (liq.), - 83.7°
Solubility in Aq. (0°), 50,300 vols. in 100 Melting-point (solid), - 110°
The gas is one-fourth heaTier than air. On account of its great
solubility and the small vapor tension of its solution, it condenses at-
mospheric moisture into a fog of drops of hydrochloric acid. On
account of its high critical point, it may be liquefied by pressure alone.
Both in the gaseous and liquefied states it is a nonconductor of elec-
tricity. Its heat of solution (p. 164) is 17,400 calories.
On account of its high concentration, the solution may be looked
upon as a mixture of liquefied hydrogen chloride and water. At 15°
and 760 mm. 454.6 volumes of the gas dissolve in 1 volume of water,
or 746 g. in 1 1. The mixture weighs therefore 1746 g. (42.7 per
cent of HC1). Its sp. gr. is 1.215. The volume of the solution is
given by the proportion 1215 : 1 : : 1746 : x, in which x = 1.437 1.
Hence the addition of 454.6 liters of the gas has increased the volume
by only 437 c.c. Now at 15° the sp. gr. of liquefied hydrogen chlo-
ride is 0.8320, and the volume of 746 g. is therefore 746 x 0.832 =
620.7 c.c. So that even if the substances had been mixed in liquid
form a considerable shrinkage would still have occurred.
When the concentrated aqueous solution is heated, it is the gas and
not the water which is driven out for the most part. When the con-
centration has been reduced to 20.2 per cent the rest of the mixture
distils unchanged at 110°. If a dilute solution is used, water is the'
chief product of distillation (about 100°), but gradually the boiling-
point rises, and, when the concentration has reached 20.2 per cent once
more, the same hydrochloric acid of constant boiling-point, as it is
called, forms the residue, It is thus impossible to separate by distil-
CHLORINE AND HYDROGEN CHLORIDE 183
lation the components of mixtures which behave in this way. This
must necessarily be the case whenever, as here, the vapor tensions of
the components separately and those of all other mixtures are higher
than that of one particular mixture. When, as is more often the case,
one of the components has a vapor tension which is lower than that of
the other and lower than that of any mixture of the two, this compo-
nent will tend to remain behind, and separation can be effected. The
separation of petroleum products (q.v.) from one another illustrates the
common case (see under Alcohol for the third possibility).
The composition of the mixture having the minimum vapor tension
varies with the external pressure, and so does the boiling-point. At 300
mm. the constant boiling liquid contains 21.8 per cent of hydrogen chlo-
ride and boils at 84° ; at 1520 mm. it contains 19.1 per cent of the gas.
The common belief that hydrochloric acid of constant boiling-point is a definite
compound is without foundation. Compounds do not vary in composition with
changes in pressure in this manner. Aqueous solutions of hydrogen iodide, hydro-
gen bromide, and nitric acid behave in the same way. But solutions of oxygen, of
ammonia, and of many liquids (e.g. alcohol) in water belong to the second of the
two classes mentioned above, and the more volatile component often leaves the
water entirely before much of the latter has evaporated.
Chemical Properties. — This compound is extremely stable, as
we might expect from the vigor with which the elements of which it is
composed combine. On being heated to a temperature of 1800° it be-
gins, however, to dissociate into its constituents.
In the chemical point of view, it is on the whole rather an indif-
ferent substance. When water is saturated with the gas at — 22° a
hydrate (HC1, 2^0) crystallizes out. This decomposes into the same
constituents when allowed to warm up again to — 18°. Hydrogen
chloride (the gas) has no action upon any of the non-metals, such as
phosphorus, carbon, sulphur, etc. Many of the metals, however, par-
ticularly the more active ones, such as potassium, sodium, and mag-
nesium, decompose it. Hydrogen is set free, and the chloride of
the metal is formed (K + HC1 — >• KC1 -f- H). Hydrogen chloride
unites directly with ammonia gas to form solid ammonium chloride
(HC1 + NHg — > NH4C1). The liquefied gas has the same properties.
Composition. — The proportion of hydrogen to chlorine by weight
in this compound is 1 : 35.18. Taking the atomic weight of hydrogen
1.008, so as to harmonize that of chlorine with 0 = 16, the ratio
becomes 1,008 : 35.45,
184
INORGANIC CHEMISTRY
The proportion by volume in which the constituents unite, and
the relation of this to the volume of the resulting hydrogen chloride,
may easily be shown in several ways. The decomposition of the solu-
tion of hydrogen chloride in water by means of the electric current
proves that the gases are liberated in equal volumes.
The apparatus in Fig. 31 (p. 94) cannot be used to show this, because, under
the increasing pressure due to the displacement of the liquid into the higher bulb,
the chlorine becomes more and more soluble, and its volume therefore falls pro-
gressively more and more below what it should be.
A special form of apparatus (Fig. 67) was devised by Lothar
Meyer to demonstrate the volumetric proportion. The central part is
the same as in Fig. 31,
but the gases go to
right and left, and dis-
place the liquid in two
inverted tubes. The
equal rate at which
this takes place on
both sides proves that
the gases are generated
in equal volumes.
In order to ascer-
tain the relation be-
tween the volumes of
the constituents and
that of the product, we
may unite the gases
and find out whether
any change in volume occurs. A tube with thick walls (Fig. 68)
is filled with the mixed gases obtained by electrolysis. By dipping
one end of the tube under mercury and opening the lower stop-
cock, it is seen that no gas leaves and no mercury enters. After
the mixture has been exploded, by the light from burning magnesium,
the same test is repeated with the same result. The pressure has
therefore remained equal to that of the atmosphere. Hence there has
been no change in volume as the result of the union. It appears
therefore, that :
1 vol. hydrogen + 1 vol. chlorine — > 2 vols. hydrogen chloride,
a result in harmony with G-ay-Lussac's law (p. 125).
FIG. 67.
CHLORINE AND HYDROGEN CHLORIDE 185
Another way of demonstrating the equality in the volumes of the hydrogen
and chlorine is to fill a wide tube, closed at each end by a stopcock, with the mixed
gases arising from electrolysis of hydrochloric acid. When the air has been
entirely displaced, the stopcocks are closed. The gases which the tube then con-
tains are present very nearly in the proportions in which they are liberated from
the decomposition of the substance. By introducing a small amount of potassium
iodide solution, the chlorine is removed. It forms potassium chloride, which
remains dissolved in the water, and free iodine, which dissolves in the excess of
potassium iodide solution (Kl + Cl — >KC1 + I). Neither product is gaseous
under the circumstances, so that the volume of the mixed gases dimin-
ishes by the amount of chlorine removed. If the stopcock is now
opened under water, the latter enters and fills half the length of the
tube. The remaining gas is easily shown to be hydrogen.
To show that the volume of the hydrogen chloride is twice that of
either constituent in the free condition, an alternative method is like-
wise available. We may completely fill a long test-tube with hydrogen
chloride, introduce into it quickly some sodium dissolved in mercury,
and, after agitation, open the closed tube under mercury. The sodium
gives sodium chloride and free hydrogen, and it is found that the
mercury enters so as to fill one-half of the tube. Since from this
experiment we learn that the hydrogen occupies half the space of the
hydrogen chloride, and from the previous experiment we know that the
volume of hydrogen is equal to that of the chlorine, we conclude that
two volumes of mixed hydrogen and chlorine would give two volumes
of hydrogen chloride.
Chlorides. — The chlorides are described individually
under the other element which each contains. For the pres-
ent we simply add to the statement on p. 179 that the
majority of the chlorides which do not interact with water
(p. 181), that is to say of the chlorides of the metals, are
easily soluble in water. The only exceptions are silver
chloride (AgCl), mercurous chloride (calomel, HgCl), cuprous
chloride (CuCl), aurous chloride (one of the chlorides of gold,
AuCl), thallous chloride (T1C1), and ordinary lead chloride
(PbCl2). The last of these is on the border line as regards
solubility. An appreciable amount dissolves in cold water, and a con-
siderable amount in boiling water.
Chemical Properties of Hydrochloric Acid. — The solution of
hydrogen chloride in water is an entirely different substance in its
chemical behavior from hydrogen chloride. It is strongly acid, turn-
ing litmus red. The gas and liquefied gas have no such property.
The solution conducts electricity, as we have seen, very well, and is
decomposed in the process (cf. p. 194).
186 INORGANIC CHEMISTRY
Many metals, when introduced into hydrochloric acid, displace the
hydrogen (p. 95), and form the chloride of the metal. In the case of
zinc the action was represented by the equation :
Zn + 2HC1 -» ZnCl2 + 2H.
The liquefied gas has no action upon zinc, and even its solution in
many solvents shows little activity. The solution in alcohol behaves
like that in water. But the solutions in toluene, benzene, and other
compounds of carbon and hydrogen, in many of which the gas is
freely soluble, are hardly affected by the presence of zinc and other
metals. These, and many other facts which we shall notice later
(see Dissociation in solution), show that the condition of this sub-
stance in aqueous solution is peculiar.
The aqueous solution of hydrogen chloride interacts rapidly with
most oxides and hydroxides, as, for example, those of zinc :
ZnO + 2HC1 -> ZnCl2 + H20,
Zn(OH)2 + -2HC1 _> ZnCl2 + 2H20.
Here no free hydrogen is obtained, since the oxygen in the oxide, and
the hydroxyl in the hydroxide, unite with it to form water. In each
case, however, the chloride of the metal is obtained. It may be noted
in passing that all acids behave in a similar manner towards oxides
and hydroxides, giving water and a compound corresponding to the
chloride (cf. p. 171). Dilute sulphuric acid, for example, gives sulphates.
In the two preceding paragraphs, three kinds of actions, each con-
stituting a different way of obtaining chlorides, have been mentioned
incidentally. There are two others which we have already en-
countered. The simplest is the direct union of the element with
chlorine (Zn -f- 2C1 — * ZnCl2). The other method is illustrated in the
case of the precipitation of silver chloride (p. 13). Here the forma-
tion of the chloride occurred by exchange of another radical for
chlorine (AgN08 -f- NaCl -» AgCl | + NaM)3). The insoluble chlo-
rides (p. 185) can be made conveniently by this plan. The formation
of the precipitates, for example that of silver chlorifle, is used as a
test for the presence of a soluble chloride in the solution.
The solution of hydrogen chloride in water sold in commerce is
known by the name of muriatic acid (Lat. muria, brine). It is a
yellow liquid which contains a number of impurities. The most com-
mon are ferric chloride, which is responsible for part of the yellow
tint, some yellow organic coloring material, arsenious chloride, and
CHLORINE AND HYDROGEN CHLORIDE 187
free chlorine. The acid frequently gives a residue when evaporated,
and this must of course represent some impurity. It is sometimes
adulterated with calcium chloride, since the price obtained depends
upon the specific gravity of the solution, and this may be raised by
dissolving calcium chloride in it.
Classification of Chemical Interactions and Exercises
Thereon. — So far we have denned ten more or less distinct kinds
of chemical change : Combination (p. 14), decomposition (p. 15),
dissociation (p. 121), displacement (p. 99), substitution (p. 176),
double decomposition (p. 99), hydrolysis (p. 181), oxidation (pp.
72, 110, 172), reduction (pp. 72, 110), and electrolysis (p. 19). In one
or two of these classes all the actions are reversible, in others some are
reversible and some are not. Illustrations of every one of these will be
found in the present chapter. The classes are not mutually exclusive.
Some actions belong to one class or another according to our point of
view at the moment. The ability readily to classify each phe-
nomenon, as it comes up, requires precisely that grasp of the frame-
work of the science which the reader must seek speedily to attain.
For example, let him classify the following actions : 1. Action of
heat on chloroplatinic acid ; 2. of potassium on water ; 3. of heat on
potassium chlorate ; 4. of chlorine on metals ; 5. of chlorine on tur-
pentine ; 6. of chlorine on potassium iodide ; 7. of chlorine on
methane ; 8. of carbon monoxide and chlorine ; 9. of sunlight on
hypochlorous acid ; 10. of sulphuric acid on salt ; 11. of zinc oxide
and hydrochloric acid ; 12. of zinc on hydrochloric acid.
13. Expand the explanation of the tendency of hydrogen chloride
to fume in moist air (p. 182).
14. Explain the interaction of steam and iron (p. 110) on mechanical
principles similar to those used in describing how hydrogen chlo-
ride is formed from salt and sulphuric acid (p. 179).
15. What is the maximum that the temperature* of the action (p. 172)
of hydrochloric acid on potassium permanganate may attain ? Why
is this interaction so much more vigorous than that where manganese
dioxide is used (p. 173) ?
16. In view of the explanations given, can you define the general
nature of the " other substances " (p. 172) which may be used to oxi-
dize hydrochloric acid ?
188 INORGANIC CHEMISTRY
SUMMARY OF PRINCIPLES.
It may be useful at this point partially to summarize the principles (general
facts) of chemistry so far as they have been developed in the preceding chapters.
These principles are given under fourteen heads below. They are stated as far
as possible strictly in terms of facts, since hypotheses are not integral parts of
chemistry, but are scaffolding temporarily employed to facilitate the erection of
the structure of the science. In a later chapter (Chap, xv), some other impor-
tant principles will be summarized in like manner. To secure more strictly logical
arrangement than has seemed advisable in the text, two conceptions which have
already been dealt with are held over to the second half of the summary, namely,
17 (valence) and 21 (chemical relations of elements). The reader should give care-
ful thought to the various points, many of which, in a backward view, will be found
to have become susceptible of improved statement. We begin the series with the
most fundamental fact of all, — the one without which no chemical work would be
possible :
1. Each substance has its own set of specific physical properties. By means
of these it is recognized and, when necessary, separated from other substances (p. 5).
2. Substances are either simple (elementary), containing only one kind of
matter, or compound, containing more than one kind of matter (p. 30).
3. In all chemical phenomena (excepting " internal rearrangements "), changes
in the material composition of bodies occur (p. 16).
4. In chemical phenomena there is no change in the total mass of the system
(P- 17).
5. Each substance has a definite material composition by weight (p. 41).
6. The proportions by weight in which all chemical combinations take place
can be expressed in terms of small integral multiples of fixed numbers, which may
be called combining weights, one for each element. That weight of each element
which combines with 8 parts of oxygen is called the equivalent weight and has the
properties of a combining weight (pp. 48, 50).
7. The proportions by volume in which all chemical interactions involving
substances in the gaseous condition take place are expressible by small integers
(p. 125).
8. From 6 and 7 it follows that the equivalent weights of all substances occupy,
in the gaseous condition and at the same temperature and pressure, volumes which
are either equal or stand to one another in the ratio of small integers (p. 126).
9. In every chemical phenomenon a transformation of energy occurs. This
results in a redistribution of the chemical energy in the substances concerned, and
also in an increase or a decrease in the total chemical energy in the system (p. 26).
DEFINITIONS : The word substance is applied to a compound of one or more
kinds of elementary matter with a certain proportion of chemical energy (p. 82).
The word element is applied to a variety of simple matter which exists only in com-
bination with energy, and often with other kinds of elementary matter as
well (p. 31).
10. In chemical phenomena there is no actual loss or gain, but only transfor-
mation of energy (p. 23).
11. Interactions which proceed spontaneously are in general those in which
internal energy is transformed into some other variety or varieties of energy
(p. 27).
CHLORINE AND HYDROGEN CHLORIDE 189
12. Each substance has its own set of chemical properties, such as :
(a) Affinity : the given substance can or can not interact with such and such
elementary and compound substances.
(6) Relative activity of the systems in (a). This is measured quantitatively
by: (a) Relative speed under like conditions (see 13, 14, 18); ()3) Relative heat
developed, when actions compared can be carried out so that all conditions are
alike ; (7) Relative E.M.F. of cell when the action is so arranged as to give
electricity (pp. 28, 76, 111).
13. The speed of every interaction is increased by raising the temperature (p. 72).
14. The speed of interactions is increased or decreased by catalytic agents,
each of which is individual in the kind and amount of its effect (p. 74),
CHAPTER XII
MOLECULAR 'WEIGHTS AND ATOMIC WEIGHTS
AVOGADHO'S hypothesis (p. 131) has proved to be by far the most
suggestive and fruitful of all the conceptions of a hypothetical nature
in the science of chemistry. We are now in a position to discuss
several of its most important applications. To speak in terms of the
hypothesis, these concern more particularly the measurement of the
relative weights of the molecules of different gaseous substances, and
the determination of the most convenient magnitudes for the chemical
unit weights (atomic weights ; cf. p. 50).
Meaning of Avogadro 's Hypothesis. — First, we must under-
stand clearly what is implied in the statement that : In equal volumes of
all gases, at the same temperature and pressure, there are equal numbers
of molecules. It means that, for instance, at 100° and 760 mm., in all
specimens of gases the average spacing of the molecules is identical.
This condition is independent of the nature of the gas — for example,
whether it is a simple or a compound substance, like oxygen and carbon
dioxide respectively, or a mixture, like air. It means that when, at
some fixed temperature, we fill the same vessel with a number of
different gases or gaseous mixtures successively, the number of mole-
cules that it will hold at a pressure, say, of one atmosphere will
always be the same. If we take care to keep temperature and pressure
the same, the equality in the number of molecules that will enter
the jar will take care of itself automatically. In what follows, to
avoid continual repetition, it is to be assumed that temperatures and
pressures are equal unless the contrary is expressly stated.*
This statement would be strictly true only in the case of gases, if such existed,
which behaved in ideal accord with the laws of Boyle and Charles. Since, how-
ever, in all gases, with the exception of hydrogen, a certain tendency to cohesion
between molecules is distinctly noticeable and its amount varies from gas to gas, the
density with which the molecules are packed is not precisely the same in any two
of them. Hence, Avogadro's hypothesis is not perfectly realized in any known
* At the first reading, the beginner is advised to omit the paragraphs set in
smaller type, excepting that on pp. 194, 105.
190
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 191
gases. In the case of hydrogen, for example, a divergence from the behavior of an
ideal gas exists, but is barely measurable, and in the case of chlorine, it amounts to
about 1^ per cent, and is quite conspicuous. This slight irregularity in the pack-
ing of the molecules, however, does not interfere with the application of this
hypothesis in chemistry.
Two Kinds of Laws in Science. — It will have been observed that the
laws of science may be divided into two kinds. Some, like those of conservation
of mass and of definite proportions, express the facts with perfect exactness. The
divergence between our experimental data and these laws we find to become
smaller and smaller the more carefully our experiments are made. The difference
between our best determinations and the ideal described by the law, is always less
than the known errors of observation. There is, however, a second class of laws
in which the opposite is the case. The more carefully our measurements are made,
the more clearly is it recognized that the second kind of law does not state the
fact with exactness for any single example. The known errors of measurement
are in these cases less than the discrepancy between the observed fact and the ideal.
The laws of Boyle, of Charles, and of Gay-Lussac are examples of this class.
The value of such laws is not impaired by the fact that actual substances are in no
case accurately described by them. The law gives us a norm of behavior to which
most of the substances conform with a fair degree of closeness. Such a law
resembles a limit in mathematics, toward which some expression tends to converge
although it does not actually reach it. The first kind of law represents the actual
behavior of materials, the second kind of law an average behavior to which no one
material adheres with perfect exactness.
Now Avogadro's hypothesis describes in figurative terms some general fact
about gases whose exact nature is unknown. Using the molecular hypothesis, we
state that this fact has to do with the uniform packing of the molecules in all
gases. But, whatever it is, no real gases conform to it with perfect strictness.
Avogadro's hypothesis is allied therefore to the second class of laws, to some
of which indeed it owes its origin.
The Relative Weights of the Molecules. — According to Avoga-
dro's hypothesis, vessels of equal size filled with different gases contain
equal numbers of gaseous molecules. Now equal volumes of different
gases differ very markedly in weight, or, in other words, the densities
of various known gases cover a wide range of values. Thus, hydrogen
is the lightest of all, chlorine is more than thirty -five times, mercuric
chloride (corrosive sublimate) vapor over one hundred and thirty-four
times as heavy. Since these different weights of equal volumes repre-
sent the weights of equal numbers of molecules, the difference must be
due to the differing weights of the molecules themselves. The densi-
ties of gases, therefore, may be taken as measures of the relative
weights of their individual molecules. The extreme significance of
this inference in chemistry will appear as we elaborate upon it.
192
INORGANIC CHEMISTRY
The various scales on which the densities of gases may be calcu-
lated, such as the weights of one liter of each gas, or the weights of
volumes equal to that of one gram of air, or of one gram of hydrogen,
are illustrated in the first three columns of the following table : *
"W^EIGHT
OF ONE
LITER,
0° AND 760
MM.
DENSITY,
AIR = 1.
DENSITY,
HYD. = l.
MOLECU-
LAR
WEIGHT,
Ox. = 32.
UNITS
COMPOSING THE
MOLECULE.
Hydrogen . .
0.090
0.0696
1.00
2.016
2xH = H2
Oxygen ....
Chlorine . . .
1.429
3.166
1.105
2.449
15.88
35.18
32.00
70.90
2xO = 02
2xCl=Cl2
Hydrogen chloride,
1.628
1.259
18.09
36.458
H + C1 = HC1
Carbon dioxide .
1.965
1.520
21.83
44.00
C+2xO = CO2
Water ....
0.8045
0.622
8.94
18.016
2xH + O = H,O
Mercury . . .
8.932
6.908
99.25
200.0
Hg
Mercuric chloride,
12.097
9.354
134.43
270.90
Hg+2xCl=HgCl2
1.293
1.00
14.37
28.955
[Mixture]
The values for water (b.-p. 100°), mercury (b.-p. 357°), and mer-
curic chloride (b.-p. 300°) are measured at high temperatures and re-
duced by rule (p. 91) to 0° and 760 mm. All the numbers in the first
three columns, as they stand, are purely physical in derivation. Those
in the second column are obtained from those in the first by using the
proportion :
1.293 (wt. 1 1. air) : 1.00 (air = 1) : : wt. of 1 1. any gas : x (dens, of that gas).
To get those in the third column we employ the proportion :
0.090 (wt. 1 1. hyd.): 1.00 (hyd. = 1) :: wt. of 1 1. any gas :x (dens, of that gas).
The last two columns will be explained presently. Since the num-
bers in the first column apply to equal volumes (1 L), and those in the
other two stand in constant ratios to them, the weights in each of these
two columns represent equal volumes also. In the second, the volume
is -1^ L, and in the third, ^ 1. The values in any one of the
columns represent the relative weights of the molecules of the various
substances (see Exercise 1 in this chapter).
* To speak strictly, the density of a gas is the weight of 1 c.c. at 0° and 760 mm.
Its value for each gas is, therefore, obtained by dividing the numbers in the first
column by 1000. The numbers thus obtained are, however, inconveniently small,
and, besides, the beginner usually measures the weight of a liter of several gases in
the laboratory, and so is more accustomed to the unit employed above.
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS
193
In order to avoid the creation of unnecessary confusion in the mind of the
beginner, the weights of one liter of gas in the above table are, with the exception
of that of oxygen, all ideal numbers. They are calculated back from the correct
molecular weights made up from the atomic weights. This enables us to show the
process by which the molecular weights are derived from, the weights of one liter
without the exhibition of arithmetical discrepancies which might obscure the
principle being explained. The weights of one liter of the various gases, as we
have given them, are based on the assumption that the molecules are always
packed uniformly in accordance with Avogadro's hypothesis. The actual values
are in most cases somewhat different from these, and we attribute the divergencies
to the varying degrees of cohesion between the molecules of different substances.
Even the weight of one liter of the same gas, after reduction to 0° and 760 mm.,
is found to vary with the temperature and pressure at which it was examined.
This is but natural, since changes in these conditions alter the effects of cohesion.
The following table gives the actual weights of one liter of the same gases, with a
few additional ones, and a comparison will show the extent of the divergencies.
The most interesting case perhaps is that of oxygen and hydrogen. The chemical
combining weights of these substances are in the ratio of 15.88: 1.00, while a
slight excess of cohesion in oxygen gives the ratio of their densities the value
15.90 : 1.00. These numbers are in both cases based upon Morley's results.
WEIGHT
OF ONE
LITER,
0° AND 760
MM.
DENSITY,
Am = 1.
DENSITY.
HYD. = I.
OB-
SERVED
MOLECU-
LAR
WEIGHT.
ADJUSTED
MOLECULAR
WEIGHT.
UNITS
COMPOSING THE
MOLECULE.
Hydrogen .
0.08987
0.0695
1.00
2.012
2.016
2xH = H2
Oxygen . .
1.429
1.105
15.90
32.00
32.00
2xO = O2
Nitrogen . .
1.251
0.967
13.92
28.00
28.08
2xN = N2
Chlorine .
3.220
2.490
35.83
72.01
70.90
2xCl = Cl2
Hydrogen
chloride .
1.641
1.269
18.26
36.69
36.458
H + C1 = HC1
Carbon
dioxide .
1.977
1.529
22.00
44.22
44.00
C+2xO = CO2
Hydrogen
sulphide .
1.537
1.189
17.10
34.43
34.076
2 xH + S = H2S
Ammonia .
0.772
0.697
8.59
17.29
17.064
3xH+N=NH3
Sulphur
dioxide
2.927
2.264
32.57
65.55
64.06
S + 2 x 0 = SO2
Water . .
0.8045
0.622
8.95
18.016
18.016
2 x H + O = H O
Mercury .
8.87
6.86
98.70
198.4
200.0
Hg.
Air . .
1.293
1.00
14.39
28.955
[Mixture]
[Mixture]
The Relation betiveen the Weights of Molecules and the
Combining Weights : Atoms. — When we now scrutinize these num-
bers (p. 192) in the chemical point of view an unexpected complication
is disclosed. Using the values in any of the columns, we observe that a
molecule of hydrogen chloride actually weighs much less than a mole-
194 INORGANIC CHEMISTRY
cule of chlorine plus a molecule of hydrogen. In fact, the total weight
of the last two molecules (3.256, for example, in the first column) is
exactly twice that of the molecule of the compound formed by the
union of these substances (1.628). Evidently, to bring the chemical
behavior into harmony with the molecular hypothesis, and particularly
with Avogadro's hypothesis, on which the assumption that these num-
bers represent the weights of the molecules is based, one further sup-
position will have to be made. This is to the effect that the molecule
of hydrogen and of chlorine splits into two portions before combination,
and the resulting molecule of hydrogen chloride is composed of a half
molecule of each.
This conclusion is of importance when we are confronted, as event-
ually we must be, with the task of harmonizing our combining weights
of elements with the weights we assign to molecules. It shows us that
our units of molecular -weight for the above two elements will have to
include at least two chemical unit weights in order that this splitting
may be allowed for, without subdivision of our fundamental combining
(atomic) weights. In other words, we may not make 35.18 (column
3) the molecular weight of chlorine, and 1 that of hydrogen, for then
the molecules of hydrogen chloride would be made up of the half-
values, namely, 17.59 of chlorine + 0.5 of hydrogen = 18.09 of hydro-
gen chloride, and this would involve using 0.5 as the chemical unit of
weight for hydrogen. We prefer, therefore, to employ 2 and 70.36 for
the molecular weights of hydrogen and chlorine respectively, and 1 and
35.18 for their units of combining weight (later to be renamed atomic
weights). Or rather we employ the corresponding slightly larger num-
bers of the scale Oxygen = 16, which are 2.016 and 70.9 for the
molecular weights and 1.008 and 35.45 for the units of combining
weight.
The data for water lead to a similar conclusion for oxygen. The
proportion by weight of hydrogen and oxygen in water is 2 : 15.88
(p. 125), and this, according to the table (third column), would mean a
ratio of two molecules of the former to one of the latter. But the
molecule of water weighs only 8.94, or half the sum of these numbers.
Hence but half a molecule of free oxygen is required to form a mole-
cule of water. When, therefore, we take 16 for the unit of combining
weight in the case of oxygen, the value 32 must be assigned as the
weight of the molecule of the free element.
That the molecules of hydrogen, oxygen, and chlorine split when undergoing
combination may be shown in another way. Let us represent by small squares
195
minute volumes of the various gases. The diagram below shows the volume of the
hydrogen chloride and that of the two constituents. The two squares side by side
represent two volumes of hydrogen chloride containing, say, 100 molecules each,
according to this hypothesis. These came from (cf. p. 184) one volume of hydrogen
containing 100 molecules and one volume of chlorine containing the same number
of molecules.
HYDKOGEN CHLORIDE HYDROGEN CHLORINK
100
100
came from
100
100
The 200 molecules of hydrogen chloride must have contained altogether at least
200 fragments of chlorine, since there was a sample of it in each molecule. There
were therefore 200 combining units of chlorine, at least. There might of course
have been 400 or 600, or even more, since there might be more than one unit of
chlorine in each molecule of hydrogen chloride, but there certainly could not have
been less than 200. Now the 200 units of chlorine came from a volume containing
100 of the physical subdivisions which this hypothesis postulates. Each of these,
therefore, must have contained at least two units. Hence we arrive at the conclu-
sion that the physical particle or molecule of chlorine contains two combining units.
We note particularly that this conclusion is due entirely to the fact that we have
assumed that there is the same number of molecules in a volume of chlorine that
there is in a volume of hydrogen chloride. This conclusion is based therefore
upon a certain specific assumption, namely, Avogadro's addition to the molecular
hypothesis. At first sight we seem to have sacrificed simplicity by accepting this
enlargement of the hypothesis. We shall find, however, that its use greatly sim-
plifies the study of many things in chemistry, and so it is adopted consistently.
To apply the same reasoning to hydrogen, the hydrogen chloride must have
contained at least 200 units of hydrogen. These came from one volume of hydro-
gen gas containing 100 molecules, and there must therefore be two units of hydro-
gen in each molecule. Regarding this from the other point of view, the 100 mole-
cules of hydrogen have to supply 200 molecules of hydrogen chloride with the
element, and must therefore be subdivided in order that this may be accomplished.
Hence the molecules of hydrogen cannot be identical with the combining units.
Take now the case of water. Two volumes of water vapor are formed from
two volumes of hydrogen and one volume of oxygen. Using the same diagram-
matic method, we have two squares representing 100 molecules of water, formed
from two squares each containing 100 molecules of hydrogen and one square con-
taining 100 molecules of oxygen.
WATER
HYDROGEN
OXYGEN
100
100
came from
Two hundred molecules of water must have contained at least 200 combining
units of oxygen. These came from 100 molecules of oxygen, since they came from
one volume, and each molecule of oxygen must therefore contain two units.
Tn the foregoing we have seen that the molecules of chlorine and of
hydrogen must be held to contain two units each of the respective
196 INORGANIC CHEMISTRY
elements. These units are of the same weight as the units of the
same elements contained in the molecule of hydrogen chloride. Now,
just as the molecule is hypothetically a discrete body, so we may
regard its constituent parts as distinct entities. These parts are
spoken of as atoms, and the word is used so as to be equivalent to our
word " units." There is nothing in the facts before us that requires us
to suppose that the units or atoms are single particles : they might be
clusters of many particles, but they must have definite relative
weights, namely, the atomic weights. All that we may infer so far is
that, if the atom of hydrogen or of chlorine in a molecule of hydrogen
chloride, or the atom of oxygen in a molecule of water, is a cluster of
particles, the m,olecule of each of these elements contains twice as
many such particles. We shall use the term " atom " sparingly pend-
ing fuller explanation in the next chapter.
The Choice of the Value 16 for the Atomic Weight ofOocygen.
— When the choice of a final unit of combining weights was under dis-
cussion (p. 50), the value 16 was assigned arbitrarily to oxygen in pref-
erence to 8, the equivalent weight. The reason for this choice may be
found in the preceding sections. If the unit of combining weight had
been 8, then 16 woxild have now to be taken as the molecular weight of
the element. But this scale may easily be shown to involve a number
less than 1 for the combining unit of hydrogen. For example, the
molecule of hydrogen, being only ^^ as heavy as that of oxygen
(p. 192), must receive the value 1.008, and, since this molecule splits
in uniting with chlorine and many other elements, the unit of combin-
ing weight for hydrogen must be only 0.504 on this scale. Again,
and to put the argument differently, the weight of hydrogen chloride
occupying the same volume as 16 g. of oxygen is 18.229 g., and in this
weight of the compound the amount of hydrogen is only og-^Fg °f
the whole, or 0.504 g. Hence, to avoid numbers less than 1, the combin-
ing weight of oxygen is raised to 16 (see Exercise 4 in this chapter).
The Chemical Unit of Molecular Weight. — The considerations
discussed in the preceding sections enable us to establish a scale of
chemical molecular weights for all gaseous and volatile substances.
Since 0 = 16 is our basis of unit or atomic weights, Oxygen = 32
must be the basis of our molecular weights. The molecule of the free
element thus contains two units or atoms, and is represented by the
formula 02. The relative weights of the molecules of other substances,
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 197
on the new scale, may now be obtained by an arithmetical process
similar to those used in converting to the scale Air = 1 and Hydrogen
= 1 (p. 192). The results differ from the former ones, however, in
possessing a definite relation to chemical facts through the atomic
weights, which before was lacking.
In the fourth column of the table (p. 192) the results of this change
oi scale are given. From the proportion 1.429 : 32.00 : : wt. of 1 1. of
a gas : x, we calculate all the values in this column. That is to say,
we multiply the weight of one liter of the gas by 5-^- It will be
found that perfect harmony between atomic weights and molecular
weights has now been secured. Thus the molecule of hydrogen
chloride (36.458) is made up of half a molecule or one unit (atom) each
of chlorine (35.45) and of hydrogen (1.008). The last column gives
the composition of each molecule, the symbols being used in place of
the chemical units of weight for which they stand. It will be noted
that the molecule of mercury does not split. Hence its molecular and
atomic weights are identical, and its formula in the free condition is
Hg. Thus all the molecules of elements do not contain two units :
each case must be studied separately (see below).
The whole process, condensed in one statement, is as follows :
To obtain the molecular weight of a substance, weigh a known volume
of the substance in the state of vapor, reduce the volume by rule to 0°
and 760 mm., calculate the weight of one liter by proportion (this
gives the value in the first column), and convert to the scale oxygen
= 32 by multiplying the result by 32/1.429 (for a simpler form of this
stat3inent, see below). To record the fact that the gram is the unit
of weight employed, the above term may be expanded to the form,
gram-molecular weight. It is not very appropriate to speak of
oxygen = 32 as the unit of the system. The real unit is an imagi-
nary gas which has a density one-thirty-second of that of oxygen.
Since the process by which molecular weights are obtained is
purely one of calculation from actual measurements, its application
involves no use of the molecular hypothesis. At this point, there-
fore, we emerge from the region of fiction and reach a conception
which can be stated in terms of experimental facts alone.
The term gram-molecular weight being somewhat ponderous, we
abbreviate it to molar weight, and still further to mole. Thus, a mole
of chlorine is 70.9 g. of the element, and a mole of hydrogen chloride is
36.458 g. of the compound. The mole of a substance is the formula
weight when molecular formulae (see below) are employed.
198 INORGANIC CHEMISTRY
When the above method of calculating the molecular weight from the weight
of one liter of a substance is applied to the actual experimental values, the result-
ing molecular weights necessarily diverge somewhat from the ideal ones which we
have given. Thus, since a liter of. hydrogen chloride actually weighs 3. 221, this
number when multiplied by 32 and divided by 1.429 gives the value 72.03. This
value of the molecular weight, however, does not contain the atomic weight of
chlorine, as found by measuring combining proportions, an even number of, times.
Consequently, the molecular weights actually determined by experiment have
always to be adjusted by a slight numerical change to the nearest value which
contains even multiples of the atomic weights of the constituents. In the table on
page 193 the observed and the adjusted molecular weights have both been given.
The mole is always the adjusted molecular weight and not the observed one.
This process of adjustment does not really involve any uncertainty, because it
never requires changes large enough to introduce actual confusion. The adjust-
ment of the observed molecular weight to correspond with the measured combin-
ing weights is made, in preference to the converse operation because the latter
values are always susceptible of exact determination while the former are not.
We cannot expect the molecular weights to be accurate, because, in the first place,
the measurement of the weight and volume of a gas or vapor is always difficult,
and often involves an error of one-half per cent, and, in the second place, the
way in which the behavior of gases departs from that of a perfect gas causes
the results to vary according to the temperature and pressure chosen for the
experiment.
The Chemical Molecule. — The definition of the hypothetical
chemical molecule differs slightly from that of a molecule in the gen-
eral physical sense (p. 129). Chemical molecules are the units of
•which gaseous bodies are aggregates. We have means of comparing,
roughly, at least, the dimensions of the physical units in gases,
liquids, and even solids, and find that in the liquid and solid states
the molecules are often multiples of the gaseous ones. The gaseous
molecules are, therefore, our standard in chemistry (see, however,
Chap. xvii).
It will be seen at once that it would be quite incorrect to define the molecular
weight of a substance as the weight of a molecule. We have no definite proof of
the existence of molecules as yet, and only the roughest ideas in regard to their
weight as individuals. The molecular weight, on the other hand, is a perfectly
definite and accurately determined value which would remain in use even if the
molecxilar hypothesis were to be discarded. The conception was first reached,
however, through the use of this hypothesis, and hence we might describe a molec-
ular weight as a weight originally fixed by the use of considerations involving
molecules.
The Gram-Molecular (Molar) Volume. — The weights in the
fourth column of the table (p. 192) must represent equal volumes of
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS
199
G.M.V.
22.4 LITERS
the different gases. This follows from the fact that they are derived
from the values in the first column by multiplying by a constant ratio
(32/1.429), and the volume in the first column is always 1 liter. The
actual dimension of this volume is evidently 32/1.429 liters, which is
almost exactly 22.39, or in round numbers 22.4 liters. This volume
at 0° and 760 mm. holds 32 g. of oxygen, 70.9 g. of chlorine, 44.00 g.
of carbon dioxide, or, in fact, the molar weight of any gaseous sub-
stance. It is called, therefore, the gram-molecular volume (G.M.V.)
or the molar volume (p. 149). It may be defined as that volume
which contains one mole (gram-molecular weight) of any gas at 0° and
760 mm. At other temperatures and pressures the G.M.V. has cor-
respondingly different values.
The G.M.V. gives us a concrete conception of a molar weight.
This volume is represented by a cube (Fig. 69) 28.19 cm. (or about
11.1 inches) high. Like any other vol-
ume, it holds the same number of mole-
cules of different gases. Its capacity at
0° and 760 mm. is the number of mole-
cules in 32 g. of oxygen. Hence, in
terms of the hypothesis, the weight of
any gas which fills it bears to 32 g. the -
same ratio as the weight of a molecule
of that gas to the weight of a molecule
of oxygen. We may, therefore, state
the method of finding the molar weight of a substance more simply
than before (p. 197) : Weigh a known volume of the substance, at
any temperature and pressure at which it is gaseous, reduce this
volume by rule to 0° and 760 mm., and calculate by proportion the
weight of 22.4 liters.
It is evident that the chemical molecular weights, adjusted so as to include
whole numbers of combining weights, will not all occupy exactly equal volumes.
They represent exactly equal numbers of molecules, and the slight differences in
the closeness with which the molecules are packed (p. 190) will cause the values of
the molar volumes to differ from gas to gas. The values of this volume calculated
from the actual weights of one liter of each gas are as follows: Hydrogen, 22.40;
oxygen, 22.39; nitrogen, 22.45; chlorine, 22.01; hydrogen chloride, 22.22; car-
bon dioxide, 22.26 ; water, 22.39 ; mercury, 22.55. The average value of this volume
in the case of the more nearly perfect gases is 22.4 liters, and this is, therefore,
the number which we have used in our definition (see Exercise 3 in this chapter).
Determination, of the Atomic Weight of Each Element. —
Having learned how the total weight of the molecule is expressed,
200
INORGANIC CHEMISTRY
using that of oxygen as the standard, we must now return to the com-
bining units of weight. The principle used is that the sum of the
units (atoms) represented in the formula of each substance must equal
the molecular weight of the latter. Thus we now approach the unit
weights from a new standpoint, first ascertaining the molecular
weights, and then dividing the latter between the constituents of the
compound. The following table shows the result of this process for a
number of compounds. The molar weights there given are adjusted
molecular weights (p. 193), and round numbers are used in some
cases.
The method of calculation is as follows : Suppose that we find the
molecular weight of phosphorus oxychloride to be 153.35. We then
determine by analysis the proportions of the constituents and express
them as usual in percentages. The results are, phosphorus, 20.215 :
chlorine, 69.35 ; oxygen, 10.435. The problem is to distribute these
elements in the same proportion in such a way that the total is equal
to the molecular weight. We find the value for phosphorus, for ex-
ample, by the proportion, 100 : 20.215 : : 153.35 : x, in which we
get the value x = 31. The proportions of chlorine and oxygen are
calculated in the same way, and give the numbers shown in the fourth
line of the table.
SUBSTANCE.
MOLAR
WEIGHT.
WEIGHTS OF CONSTITUENTS IN MOLAB WEIGHT.
Hydro-
gen.
Chlo-
rine.
Oxygen.
Phos-
phorus.
Car-
bon.
Mer-
cury.
Hydrogen chloride
Chlorine dioxide .
Phosphorus trichloride .
Phosphorus oxychloride
Phosphoric anhydride .
Phosphine
36.45
67.45
137.35
153.35
284
34
32
48
18
16
26
28
30
74
00
200
235.45
270.9
1
35.45
35.45
106.35
106.35
32
'l6
160
'31
31
124
31
12
24
24
12
36
24
200
200
200
3
Oxygen
32
48
16
16
32
32
Ozone
"Water
2
4
2
4
2
6
4
Methane
Acetylene
Ethylene
Formaldehyde ....
Methyl acetate . . .
Acetic acid
Mercurv
Mercurous chloride .
Mercuric chloride . .
• •
->5.45
70.9
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 201
Study of such a table as this, if it were extended to include all
known volatile substances, would enable us promptly to pick out the
most convenient values for the unit weights of the individual elements.
Thus, no compound containing chlorine is known whose total molecular
weight contains less than 35.45 parts of this element, although many con-
tain multiples of this amount. Hence this will serve as the unit quantity
for describing the proportion in any compound, and will be represented
by the symbol Cl. Similarly, 31 is the smallest amount of phosphorus
in the table, and is an integral factor of the larger amount. Hence
P = 31. For the same reason we choose 12 for carbon, 200 for mer-
cury, and 1 (more exactly 1.008) for hydrogen.
It will now be seen why the equivalents (p. 50) were multiplied by
various integers in making the chemical units. The equivalent of
carbon was 3. But there is no compound whose molecular weight con-
tains less than 12 parts. Now, for reasons that will appear later,
chemists always adjust the formula of each volatile substance so that
the total formula-weight (see below) is equal to the molar weight. To
accomplish this they multiply, when necessary, all the symbols in the
formula by some integer (e.g., N204). It would thus lead to needless
complication to take 3 as the unit amount of carbon, for every molecule
would then contain four units, or some multiple of four, and every
formula C4 or some multiple of C4. Having decided that molar weights
are to determine the dimensions of the total formula-weight of every
compound, we choose the largest units of combining weight that we
can, in order that all the chemical transactions of each element may
be formulated by means of the smallest possible numbers of units.
The mode of derivation makes all the quantities of any one ele-
ment in the above table integral multiples of the smallest. Of course,
an element might be found of which no volatile compound contain-
ing the unit weight was yet known. If methane and formaldehyde
were still undiscovered, and only the other four compounds contain-
ing carbon were available, we should then take the greatest common
measure of 24 and 36, namely, 12, as the unit.
The chemical unit weight or atomic -weight of an element may thus
be defined as : The G.C.M. (greatest common measure) of the weights
of that element found in the molar weights of all its volatile compounds
so far as these have been examined. It is always a multiple of the
equivalent by unity or some other small integer.
This definition is simply a description of the experimental method of determin-
ing atomic weights. The same thing may be expressed iu several other ways,
202 INORGANIC CHEMISTRY
For example, we may say that the atomic weight is the smallest weight of the
element found in the molar weight of any of its volatile compounds. This, how-
ever, is a less satisfactory definition, because it fails to cover the possibility of no
compound containing a single unit (atom) being known.
We may also define the atomic weight of any element as the largest even mul-
tiple of the equivalent which can be contained in the molecular weights of all the
volatile compounds of the element. That the atomic weight, determined by the
above method, must necessarily be a multiple of the equivalent weight, hardly re-
quires demonstration. In compounds containing oxygen, Gay-Lussac's law shows
that the quantity of oxygen in the molecular weight of the compound must be an
even multiple or sub-multiple joi. 32 and therefore an even multiple of 8. Hence
the weight of the other constituent combined with this amount of oxygen must
be an even multiple of the quantity which would combine with 8 parts of oxy-
gen. That is to say, it must be the combining weight, or an even multiple of it.
If the compound does not contain oxygen, then, according to Gay-Lussac's law,
the volumes of the compound and of oxygen which would interact must stand to
one another in the ratio of two integers. From this fact, by similar reasoning,
we reach the same conclusion as before.
The application of the criterion which Avogadro's hypothesis furnishes for the
determination of chemical unit weights was neglected for many years. Although
it was first suggested by Avogadro in 1811, and almost immediately afterwards, in-
dependently, by Ampere, it was not until Cannizzaro, still professor of chemistry
in Rome, recalled attention to its usefulness in 1858, that the values to which it
led began finally to be adopted by chemists. Consequently, during the first half
of the nineteenth century, through the absence of any definite standard for chemi-
cal unit weights, much confusion existed. This was not remedied very promptly,
even after 1858, for the unit weights based upon Avogadro's principle were not
accepted by all chemists simultaneously. Thus, when O = 8, the formula of water
is HO, and this mode of writing its composition will be found in many chemical
works published during the first two-thirds of the nineteenth century.
There is one feature of the definition to which attention must be drawn spe-
cifically. The atomic weight, according to the definition, is chosen from amongst
the values given by known volatile compounds. When these are few in number,
there is always a possibility that the actual minimum has not been found
amongst them. There are, however, as we shall see, other independent ways of
fixing the values of the atomic weights, and we have now good reason to believe
that few of those at present in use will require to be altered in consequence of
future discoveries.
The Accepted Atomic Weights of the Elements. — The most
trustworthy measurements have been selected in making up the fol-
lowing list of unit weights. It is the one approved by an international
committee (1906) consisting of F. W. Clarke, T. E. Thorpe, Karl
Seubert, and Henri Moissan.
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS
203
NAME OP ELEMENT.
SYM-
BOL.
ATOMIC
WEIGHT.
ELEMENTARY SUBSTANCE
FIKST ISOLATED BY:
DATE.*
Aluminium ....
Al
27.1
Wohler
1827
Antimony (Stibium). .
Argon
Sb
A
120.2
39.9
Basil Valentine ....
liayleigh and Ramsay .
15th Cent.
1894
Arsenic
As
75.0
Albertus Magnus
13th Cent
Barium
Ba
137.4
Davy
1808
Bismuth . . • . • .
Bi
208.5
Ment. by Basil Valentine
15th Cent
Boron
B
11.0
Gay-Lussac and The"nard .
1808
Bromine .....
Br
79.96
Balard
1826
Cadmium
Cd
112.4
Stromeyer
1817
Caesium
Calcium
Cs
Ca
132.9
40.1
Bunseu and Kirchhoff .
Davy
1861
1808
Carbon
C
12.00
Prehistoric
Cerium
Ce
140.25
Mosander
1839
Chlorine
Cl
35.46
Scheele
1774
Chromium
Cr
52.1
Vauquelin
1797
Cobalt
Co
59.0
Brand
1735
Columbium (Niobium) .
Cb
94
Hatchett
1801
Copper (Cuprum)
Cu
63.6
Prehistoric
Erbium
E
166
Mosander
1843
Fluorine
F
19
Moissan
1886
Gallium
Ga
70
Lecoq de Boisbaudran .
1875
Germanium ....
Ge
72.5
Winkler
1886
Glucinum (Beryllium) .
Gold (Aurutn)
Gl
Au
9.1
197.2
Wohler and Bussy . . .
1828
Prehistoric
Helium
He
4
Ramsay and Cleve .
1895
Hydrogen .
H
1.008
Cavendish
1766
Indium
In
115
Reich and Richter
1863
Iodine
I
126.97
Courtois ....
1811
Iridium
Ir
193.0
Smithson Tennant .
1804
Iron (Ferrum)
Fe
55.9
Prehistoric
Krypton
Kr
81.8
Ramsay and Travers
1898
Lanthanum ....
La
138.9
Mosander
1839
Lead (Plumbum) . . .
Lithium
Pb
Li
206.9
7 03
Mentioned by Pliny .
Arfvedson ....
Prehistoric
1817
Magnesium ....
Manganese
Mg
Mn
24.36
55.0
Liebig and Bussy . . .
Gahn
1830
1774
Mercury (Hydrargyrum)
Molybdenum ....
Hg
Mo
200.0
96.0
First ment. by Theophrastus
Hjelm
300 B.C.
1790
Neodymium ....
Neon
Nickel ....
Nd
Ne
Ni
143.6
20
58 07
Auer von Welsbach . .
Ramsay and Travers .
Cronstedt
1885
1898
1751
Nitrogen
N
14.04
Rutherford
1772
Osmium
Os
191
Smithson Tennant .
1804
Oxygen
o
16.00
Priestley
1774
Palladium
Pd
106.5
Wollaston
1803
Phosphorus ....
P
31.0
Brand
1674
Platinum
Pt
194.8
16th Cent.
Potassium (Kalium) .
K
39.15
Davy
1807
* In cases where the elementary substance has not been isolated, the date of
admission to the list of elements is given.
204
INORGANIC CHEMISTRY
NAME OF ELEMENT.
S YlH-
EC!,.
ATOMIC
WEIGHT.
ELEMENTARY SUBSTANCE
FIBST ISOLATED BY :
DATE.*
Praseodymium . . .
Radium
Pr
Ra
140.5
225
Auer von Welsbach . . .
Curie
1885
1902
Rhodium
Rh
103.0
Wollaston
1803
Rubidium
Rb
85.5
Bunsen and Kirchhoff .
1861
Ruthenium
Ru
101.7
Claus
1845
Samarium
Sa
150.3
Lecoq de Boisbaudran .
1879
Scandium
Sc
44.1
Nilson and Cleve ....
1879
Selenium ....
Se
79.2
Berzelius . . - .
1817
Silicon
Si
28.4
Berzelius
1823
Silver (Argentum)
Ag
107.93
Prehistoric
Sodium (Natrium)
Na
23 05
Davy
1807
Strontium
Sr
87.6
Davy
1808
Sulphur
s
32.06
Prehistoric
Tantalum
Ta
183
Eckeberg
1802
Tellurium
Te
127.6
Muller von Reichenstein .
1782
Thallium
Tl
204.1
Crookes
1861
Thorium ....
Th
232.5
Berzelius
1828
Tin (Stannum)
Sn
119.0
Prehistoric
Titanium ....
Ti
48.1
Gregor
1789
Tungsten (Wolfram)
Uranium
W
u
184
238.5
Bros, d' Elhujar ....
Peligot
1783
1841
Vanadium
V
51.2
Berzelius
1831
Xenon
X
128
Ramsay and Travers
1898
Ytterbium > .
Yb
173 0
Marignac
1878
Yttrium
Yt
89.0
Wohler
1828
Zinc
Zn
65.4
Ment. by Basil Valentine .
15th Cent.
Zirconium
Zr
90 6
Berzelius
1825
Molecular Formulae of Compounds. — Since we have all along
been using the symbols to represent the atomic weights of the elements,
no change in this respect will be required.
The formulae of compounds are used to express the proportion by
weight of the constituent elements. To this must now be added the
condition that the total formula-weight must be made up so as to
equal the molecular weight, in all cases where the latter is known.
Thus-, in acetylene (p. 200) the proportion of carbon to hydrogen, ex-
pressed in the simplest terms, is 12 : 1 (= CH). But the molecular
weight is 26, hence the formula must be C2H2. Similarly in ethy"-
lene the expression CH, would correctly represent the proportion by
weight of the constituents, but not the molecular weight. The latter
compels us to use C2H4 as the formula. In like manner acetic acid is
* In cases where the elementary substance has not been isolated, the date of
admission to the list of elements is given.
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 205
C2H402 and not CH20. Here the usage is justified by the fact that,
although the composition of acetic acid is the same as that of formal-
dehyde (CH20), the materials themselves are entirely different. The
molecular formulae C2H402 and CH20 show that they are different (see
Exercise 2 in this chapter).
Molecular Formulas of Simple Substances. — Hydrogen in the
free condition has two chemical units in its molecule (p. 194), and hence
its molecular formula is H2. In other words, two unit weights (atoms)
of the element (cf. p. 32) combine with each other to form the free sub-
stance. In combination with other elements, as in PH8, HC1, CjH^
and H20, any number of unit weights of hydrogen may be used.
Similarly the molecular formula of free chlorine is C12, since its molec-
ular weight is 70.9. On the other hand, that of mercury is Hg, for
200 (p. 200) is at once its unit of combining weight and its molecular
weight. The molecules of many other elements, such as cadmium,
zinc, potassium, and sodium, contain but one chemical unit (atom).
On the other hand, phosphorus vapor has the molecular formula
P4 and sulphur vapor S8, although in these cases measurements at
higher temperatures give smaller values. Thus at 1700° the molecular
weight of phosphorus is 91, showing it to be composed partly of mole-
cules of the formula P2. Sulphur vapor at 800 becomes wholly S2.
When the temperature is lowered, S8 is re-formed. Iodine (q.v.) supplies
the most interesting example of this kind.
The molecular weight of many of the elements is by no means a fixed quantity.
Its value often depends, as the above examples show, upon the temperature at
which it is measured. Thus, while the molecular weight of oxygen does not vary to
any appreciable extent between its boiling-point (— 182.5°) and 1700°, many of the
other elements show very marked changes. This seems to indicate that while
there is only one stable molecular form of the substance oxygen, there are several
molecular arrangements of some of the elements which have at least a certain
degree of stability. While the data in regard to many of these cases are given
under their respective elements, it may be worth while to collect them together in
this place in order to show between what limits the variations occur. An examina-
tion of the following table exhibits the fact that, for example, bromine is all Br2
up to 750°, but beyond this there is a small but measurable amount of dissocia-
tion. Again, iodine is all I2 up to 448°. Beyond this temperature the density
diminishes, and when 1700° has been reached the molecules have all been com-
pletely dissociated. Phosphorus, even at 1700°, is only partially decomposed into
molecules of the formula P2, while arsenic at the same temperature is almost
completely Asr
206
INORGANIC CHEMISTRY
Molecular Weights of the Elements at Various Temperatures.
ELEMENT.
TEMP.
OBSERVED
MOL. WT.
NEAREST MULTI-
PLE OF ATOMIC
WEIGHT.
KEMARKS.
Antimony. .
1640°
283
Sb2 = 240.4
Some Sb4
Arsenic . . .
644°
308.4
As4 = 300
All As4
Arsenic .
1700°
159.0
As2 = 150
All As2
Bismuth .
1640°
293.0
Bi2 =417.0
Much Bi
Bromine . . .
750°
160
Br2 = 159.9
All Br,
Bromine . .
1050°
150.5
Br2 = 159.9
Slight dissociation
Cadmium .
1040°
114
Cd =112
AllCd
Iodine . ' .
448°
254.8
I, = 253.94
All I,
Iodine ....
1700°
127
I = 126.97
( Complete dissocia-
| tion (I)
Mercury .
448°
198.5
Hg =200
All Ha
Mercury . . .
1730°
198
Hg =200
Still Hg
Oxygen . . .
1700°
32
02 =32
A11O2
Phosphorus .
313°
128
P4 = 124
A11P4
Phosphorus .
1700°
91.2
P4 = 124
Much P,
Potassium . .
Red heat
37.6
K = 39.15
All K '
Selenium . .
860°
222
Se2 = 158.4
Selenium . . .
1420°
164.4
Se, = 168.4
Nearly all Se2
Sodium . . .
Red heat
25.4
Na = 23.05
All Na
Sulphur . . .
193°
251
S8 =256.5
(2.1 mm. press )
Sulphur . . .
800°
64.2
S., = 64.1
A118.
Sulphur . . .
1719°
63.6
S.; = 64.1
Still S2
Tellurium
1400°
260
Te2 = 255.2
.
Thallium . . .
1730°
412.4
Tl, = 408.2
All T12
Zinc ....
1740°
76.4
Zu" = 65.4
All Zn
The case of oxygen demonstrates clearly the necessity of using
molecular formulae, even for simple substances. The table of analyzed
molecular weights (p. 200) shows two substances containing nothing
but oxygen. Ozone (q.v.) has a molecular weight 48, being a gas
exactly one-half heavier than ordinary oxygen. Its formula, there-
fore, is 08, while that of oxygen is 02. Oxygen and ozone are entirely
different chemical individuals. The latter has, for example, much
greater activity, as is shown by the fact that silver rusts when exposed
to it.
The molecular constitution of the familiar gases, like hydrogen
and chlorine, suggests one reason for the fact that they do not unite
at once when mixed. If they were made iip of free chemical units, it
is difficult to see how these could continually encounter one another
as they moved about, without combining. But the molecular formulae
suggest that each gas is really in combination already (with itself),
and hence is indifferent to the presence of the other. Indeed, the fact
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 207
that hydrogen shows no evidence of dissociation into atoms when
heated, since its molecular weight is the same at all temperatures,
proves it to be a very stable combination. Both water and hydrogen
chloride show distinct signs of decomposition when treated in the
same way. When a mixture of hydrogen and chlorine is heated so as
to dissociate a few of the molecules, the union of the elements cross-
wise begins at once, and the heat developed by this union quickly car-
ries the action through the whole mass. Of course this explanation
is based upon our hypothesis and, as such, is of the same imaginary
description as everything connected with that hypothesis. It cannot
be verified by experiment.
While the use of the molecular hypothesis, in explaining the constitution of
elements like hydrogen and chlorine, and their indifference to one another when
mixed, is probably the most easy for the beginner to understand, it is not the
safest in the scientific point of view. It is preferable to employ the conception of
energy and to say that when the mixture is heated locally, the amount of energy
in the parts of the components which are affected is increased, and consequently
their chemical activity is raised. When the energy which they contain reaches a
certain potential, the chemical change begins.
To prevent misconception, it must be repeated that these formulae,
02 and 03, apply only to the forms of free oxygen, and at present
these are the only forms of the simple substance which are known
with certainty. In combination, any number of units of the element
may be contained in a composite molecule. Thus we have H2O, C02,
HX08, H,C204, C6H1206 (grape sugar), and so forth, without limit.
Attention must be called specifically to the fact that the chemical
unit of combining weight for an element can not be determined by
measuring the molecular weight of the simple substance. We have
no means of knowing how many units are contained in the molecule
until the unit itself shall have been fixed. This is done by comparison
of the molecular weights of several compounds of the element with
the results of analysis of those compounds (p. 200).
Dissociation of Compounds. — Dissociation of the kind under-
gone by sulphur vapor and by iodine vapor (p. 205) is not confined to
simple substances. Phosphorus pentachloride can be converted into
vapor without much difficulty. Its molecular weight, if it underwent
no chemical change during the volatilization, would be 208.25. The
density actually observed at 300° gives by calculation not much more
than half this value. The explanation is found when we examine the
208 INORGANIC CHEMISTRY
nature of the vapor more closely. We find that it is a mixture of
phosphorus trichloride and free chlorine, resulting from a chemical
change according to the equation : PC15 <=> PC18 + 'C12. The low value
of the density tells us at once that dissociation has taken place.
From the value of the density at various temperatures, we may
even calculate the proportion "of the whole material which is disso-
ciated. At 300° it is 97 per cent ; at 250°, 80 per cent ; and at 200°,
48.5 per cent (see Exercise 10 in this chapter).
Chemical Interactions Studied by Observation of Volumes:
Molecular Equations. — According to Avogadro's hypothesis, if we
filled a succession of vessels of equal dimensions with different gases,
and could arrest the motion of the particles and observe their dispo-
sition, we should find that the average distance from particle to par-
ticle would be the same in all cases. This would be true whether our
vessels were filled with single gases, with homogeneous mixtures, or
with gases in layers. Such being the case, if any chemical change is
brought about in the mass which results in a multiplication of the
molecules, it is evident that the volume will have to increase in order
that the spacing may remain the same as before. If any chemical
action results in a diminution of the number of molecules, then a
shrinkage must take place in order that the spacing may be preserved
as before. Thus in a mixture of hydrogen and chlorine, according to
our hypothesis, neighboring molecules of hydrogen and chlorine simply
exchange units, so that H — H + 01— 01 becomes H — C1 + C1 — H.
There being no alteration in the number of particles, no change in
volume occurs. In the case of water, on the other hand,
H-H + 0-0 + H-H becomes H-O-H + H-O-H.
Since the oxygen molecules, which form a third of the whole, dis-
appear into the molecules of hydrogen, the tendency to preserve
spacing results in a diminution of the volume by one-third (p. 126).
This method of looking upon chemical interactions between gases
gives us the nearest sight which we can have of the behavior of the
molecules themselves. We cannot observe the individual molecules,
but in consequence of the spatial arrangement which we suppose them
to observe, the outward change in the volume of a large aggregate of
molecules enables us to draw conclusions at once in regard to the
behavior of the single molecules in detail.
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 209
Take, for example, the case of the burning of solid sulphur in
oxygen gas within, a closed space. We find that the volume of the
gas does not alter during the process. It takes up the sulphur into
itself. The fact that no alteration in volume occurs, leads us to infer
that there are no more gaseous molecules after the action than before
it. A molecule of oxygen contains two units, so that the molecule of
the product must also contain two units of oxygen, with sulphur in
addition. This harmonizes with the formula employed for the sub-
stance which, as we have seen, is S02.
Now, by using in our equations the molecular formulae for all
gaseous substances, we gain the advantage of having a record of the
changes in volume as well as of the proportions by weight. Thus, if
we write S + 20 — >• S02, the equation shows the proportions by weight
correctly. But the equation
S + 0, -> SO
records, in addition, the fact that the number of molecules of oxygen
and of sulphur dioxide is the same. These being the only gases,
there is therefore no change in volume. Again, if we write H + Cl— >
HC1, the proportions by weight are correct. But it appears as if two
gaseous particles had given one, and that therefore the volume should
have been reduced to one-half, which is not the case. The molecular
equation,
shows the same relative weights of material, together with the addi-
tional fact that there are equal numbers of molecules of the two
ingredients, and as many molecules in the product as there are in both
ingredients together. In such an equation each formula of a gaseous
substance represents one volume. The most instructive chemical
equations are thus the molecular ones, and they will be used exclu-
sively in the sequel (see Atomic hypothesis).
When we know the molecular formulae of the single substances
concerned in an action, the equation can be made and the change in
volume determined without actual measurement. For example : What
volume change will be observed when a mixture of carbon monoxide
and oxygen has exploded, and the temperature has once more reached
that of the room ? The molecular formulae are CO, 02, and C02. The
equation representing the weights is CO + 0 — » C02. The molecule of
210 INORGANIC CHEMISTRY
oxygen, however, being 02, we cannot employ less than this quantity
in a molecular equation, so that the equation becomes :
02-»2C02.
Three molecules, therefore, give two, throughout the whole mass,
and therefore three volumes will become two if the pressure and tem-
perature are the same at the beginning and end of the action.
If we remember that all volatile compounds of carbon and hydro-
gen burn to form water and carbon dioxide, the molecular equation
for any such combustion may easily be made, and the volumes of all
the materials ascertained. The volumes of solids and liquids are
always insignificant. Hence, if water is a product, only its volume
as steam is given by the equation (see Exercises 2 and 5-9 in this
chapter).
The molecular formula of acetylene is C2H2. When this gas is mixed with
oxygen in sufficient amount, a violently explosive mixture is obtained. The chem-
ical action results in the formation of carbon dioxide and water. The equation
C2H2 + 5O — » 2C02 + H2O represents the change so far as the weights of material
are concerned. But we must use nothing but complete molecules if we wish to
learn the volume relations. To secure the oxygen in multiples of two combining
weights, we must double the whole equation, when it appears in the form :
2C2H2 + 502 -> 4CO2 + 2H2O.
We then see that seven molecules become six, so that the change in volume, above
100°, will be in the same proportion. Below 100° the water condenses, and the
carbon dioxide is the only gas produced.
Other Methods of Determining Molecular Weights. — The
method we have described is necessarily restricted in its application to volatile
substances which are not decomposed when converted into vapor. The similarity
between the dissolved and the gaseous states (cf. p. 150), however, opens up a much
wider field for the use of what is essentially the same method. Subject to certain
exceptions, equal volumes of solutions having the same temperature and osmotic
pressure contain equal numbers of molecules. Hence all the reasoning we have
applied to gases may be used mutatis mutandis in connection with solutions (see
Chap. xvii).
Properties Possessed by Atomic Weights, but not by Equiv-
alents. — It is manifest that equivalents are much simpler in nature
and much more easily ascertained than atomic weights. It will be
expected, therefore, that we shall be able to show that the units
Ka = 23, Cu = 63.6, Al = 27.1, C = 12, etc., give a better view of the
relations of the elements than do the equivalents 23, 31.8, 9.03, and 3
respectively. Now, these unit (atomic) weights are found to have rive
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS
211
important properties, at least,* which would disappear from view
entirely if equivalents were employed. This gain may be set against
one general law appertaining to equivalents which would be sacrificed
if we were to use atomic weights exclusively. Hence our distinct
preference for the atomic weights, as well as our retention of the
equivalents for certain purposes.
The first great property of atomic weights is that they are in close
correspondence with the volume relations in chemical change, which the
equivalents are not. This matter has formed the main subject of the
present chapter. The atomic weight of an element can have but one
value and is definitely determinable. The equivalent may have more
than one value, because an element, if it gives several series of com-
pounds (p. 105), will have as many different equivalents.
The other three properties of atomic weights concern : Dulong and
Petit's law, the periodic system, and the conception of molecular
structure. The property of equivalents which is not possessed by the
unit weights is described by Faraday's law (q.v.).
Dulong and Petit's Law. — It was first pointed out (1818) by
Dulong and Petit, of the Ecole Polytechnique in Paris, that when the
atomic weights of the elements were multiplied by the specific heats of
the simple substances in the solid condition, the products were approxi-
mately the same in all cases. In other words, the specific heats are
inversely proportional to the magnitudes of the atomic weights. The
table, in which round numbers have been used for the atomic weights,
shows that the product lies usually between 6 and 7, averaging about
6.4:
ELEMENT.
ATOMIC
WT.
SP. Hx.
PRODUCT.
ELEMENT.
ATOMIC
WT.
SP. Hx.
PRODUCT.
Lithium . .
7
.94
6.6
Iron .
56
.112
6.3
Sodium . .
23
.29
6.7
Zinc . .
65.4
.093
6.1
Magnesium
Silicon .
24.4
28.4
.246
.16
6.0
4.5
Bromine
(Solid)
Gold
80
197
.084
.032
6.7
6.3
Phosphorus
(Yellow)
Calcium
31
40
.19
.170
5.9
6.8
Mercury
(Solid)
Uranium
200
238.5
.0335
.0276
6.7
6.6
•* A sixth is mentioned in Chap, xvii, section on the ionic hypothesis.
212 INORGANIC CHEMISTRY
The conspicuous exceptions occur among the elements with low atomic weights
only. The products for four of these, the atomic weights being given in parenthe-
ses, are as follows : Glucinum (9), 3.7, boron (11), 2.8, carbon (12), L7, and silicon
(28.4), 4.5. Further investigation shows, however, that in the case of precisely
these elements the value found for the specific heat varies very markedly with the
temperature at which the specific heat is measured. Their specific heats become
rapidly greater the higher the temperature. Thus at 985° the specific heat of the
diamond (carbon) is 0.45, which when multiplied by the atomic weight gives the
product 5.5. So that even the exceptional elements tend to come into line when
the specific heat is measured at higher temperatures — and Dulong and Petit's state-
ment does not limit our choice to any one temperature.
Another way of expressing this law will give it greater chemical
significance. The specific heats are the amounts of heat required to
raise equal weights of the various elements through one degree. Now
these equal weights contain fewer chemical units in proportion as the
chemical unit weight is greater. Hence this law may be put in the
form : Less heat is required to raise the temperature of equal weights
in proportion as the number of chemical units they contain is smaller.
In fact, equal numbers of units require equal amounts of heat to raise
them through equal intervals of temperature. In other words, atomic
•weights of all elements have equal capacities for heat.
This being true, the equivalents, if used instead of the atomic
weights, must give widely varying products. The quantities of heat
required to raise equivalent weights through one degree are either equal
to, or are fractions of, those required for the atomic weights, according
to the valence of the element, Hence the law applies only to atomic
weights, and not to equivalents.
It will be seen at once that although the law of Dulong and Petit
is purely empirical, it may nevertheless be used for fixing the atomic
weight of an element of which no volatile compounds are known. We
can always measure the equivalent with considerable exactness, and,
when this has been multiplied by the specific heat of the free sub-
stance, we can see at a glance what integral factor will raise the
product to the neighborhood of 6.4. For example, analysis shows us
that in calcium chloride the proportion of chlorine to calcium, using
the known atomic weight of chlorine as one term of the proportion, is
35.5 : 20. If calcium is univalent, 20 is its atomic weight. If it is
bivalent, two units of chlorine are combined with 40 parts of calcium,
and 40 is its atomic weight. If it is trivalent, three units of chlorine
are united with 60 parts of calcium, etc. All we learn in reference to
the atomic weight of calcium from this analysis is that its value
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 213
is 20 or some integral multiple of 20. Nor can we fix the upper limit,
for we are unable to obtain the weight of a known volume of calcium
chloride vapor and so determine the molecular weight. But the specific
heat of calcium being 0.170, we multiply this number by 20, and get
the product 3.4. This is only half large enough, so we assume that 40
is a more probable value for the atomic weight of calcium. The prod-
uct is then 6.8, which agrees fairly well with the average for other
elements. We decide, therefore, that the symbol Ca shall represent
forty parts by weight. The formula of calcium chloride is therefore
CaCl2, and calcium is bivalent.
It will be seen that this does not supply us with a method of ascertaining
chemical unit weights independently of any chemical experiment. We cannot
measure the specific heat and use the quotient from division of this number into
6.4, for we do not know in advance that the product for the element will have
exactly this value. It may be below 6, or it may be as high as 7. In the case of
calcium, for example, 6.4 -f- 0.17 = 37.65. Now 37.65 is 5 per cent below the real
value of the chemical unit, and even the roughest measurement of a chemical com-
bining weight need never be more than 1 per cent in error. Hence the atomic
weight must be founded upon the determination of the equivalent, which can
be measured with accuracy. The rule discussed in this section can be used only
to ascertain what multiple of the equivalent shall be accepted as the atomic weight
after the equivalent itself has been measured with care. In other words, this is a
method of adjusting the result of chemical experimentation, and cannot supersede
it altogether.
Periodic System : Molecular Structure. — The periodic system
(<?.«.) is an arrangement of the elements in the order of magnitude of
their atomic weights. The elements when placed in this order are
found to be grouped in a highly remarkable manner according to chemi-
cal likeness. An arrangement in the order of the equivalents leads to
no significant grouping at all.
The conception of molecular structure finds its chief application in
the chemistry of the compounds of carbon, and a description of its most
remarkable services to chemistry would involve discussions far too
elaborate to be included in this volume. The graphic formulae (q.v.)
which represent this structure find little application in inorganic
chemistry. A simple illustration only can be given.
If the symbol C stood for three parts (the equivalent) of carbon,
methane would receive the formula (7H, the proportion of its constitu-
ents by weight being 3 of carbon and 1 of hydrogen. But, when we
mix this gas with chlorine and expose the mixture to sunlight, no less
than four different compounds are produced (cf. p. 176). WithC = 3,
214 INORGANIC CHEMISTRY
their formulae would be QHgCl, <72HC1, C^HClg, and CC1, the cai-bon
being univalent. The relations of these substances are much simplified
when we change to molecular formulae and substitute C = 12 for Ct = 12,
making the carbon quadrivalent. We then perceive that we are dis-
placing successively the four hydrogen units in one molecule, and that
the substances are CH^ CH8C1, CH2C12, CHC13, and CC14. The whole
prodigious growth of the chemistry of the compounds of carbon, which
has taken place during the last half century, has been the result of the
employment of this seemingly slight improvement by Kekule and Cou-
per, independently of one another, in 1858.
Thus, quite aside from the molecular hypothesis and Avogadro's
addition to it, we have now found ample independent justification for
the multiplication of the equivalents by integers and for the concep-
tion of valence which results from this. Indeed, if Avogadro's hypoth-
esis had not suggested this multiplication, it would have come about
eventually even in the absence of any molecular hypothesis. Dulong
and Petit's law would have been sufficient alone to direct attention to
the atomic weights, and their other properties could not long have
escaped discovery.
The existence of the law of Dulong and Petit and the periodic law, together
with the services of structural formulae to organic chemistry, all demonstrate that
atomic weights are of vastly greater significance in the science than are equivalent
weights. And there are other immense ranges of facts aside from those covered
by these conceptions, which are all dependent upon the atomic weights. That
almost the whole systematization that has been secured in chemistry should thus
center in this one point, furnishes the strongest circumstantial evidence that
Avogadro's hypothesis represents very closely some fundamental property of all
gases. This independent inductive evidence in favor of Avogadro's principle is
especially worth noting because the deduction of the principle from the data of the
kinetic-molecular hypothesis is not absolutely rigid. It involves certain assump-
tions which, while they are plausible enough, are still assumptions.
Application to the Case of a New Element. — By way of review-
ing the principles explained in the chapter, let us apply them to the
imaginary case of a newly discovered element. The bromide of the
element is found to be easy of preparation and to be volatile. It con-
tains 30 per cent of the element, and its vapor density referred to air
is 11.84.
To find the equivalent of the element, that is, the amount combined
with 80 parts (the equivalent) of bromine, we have the proportion
70 : 30 :: 80 : x, from which x = 34.3. The atomic weight must be
this, or some small multiple of it.
MOLECULAR WEIGHTS AND ATOMIC WEIGHTS 215
The G.M.V. of air weighs 28.955 g. (p. 192). Hence the same
volume of the vapor of this bromide, which is 11.84 times as heavy as
air, will weigh 28.955 x 11.84, or 342.84 g. This is therefore the
molar weight of the compound.
Now 30 per cent of this is the new element. 342.84 x 30 -*- 100 =
102.85. So long as no other volatile compound is known, we adopt
this as the atomic weight. The rest of the molar weight (240 parts) is
bromine. Thus the formula of the compound is ElBrs, and from this
we see that the element is trivalent.
In case no volatile compound of the element can be formed, some of
the free simple substance is made, say by electrolysis, and its specific
heat is determined. Application of Dulong and Petit's law then gives
the multiple of the equivalent that constitutes the atomic weight.
Exercises. — 1. The weight of 1 1. of a gas at 0° and 760 mm. is
5.236 g. What is the density referred to air and to hydrogen, and what
is the molecular weight (pp. 192, 199) ?
2. The molecular weight of cyanogen is 52.08. It contains 46.08
per cent carbon and 53.92 per cent nitrogen. What is its density
referred to air, and what the weight of 1 1. at 0° and 760 mm. ? What
is the formula of the substance (p. 204) ? Exploded with oxygen it
forms carbon dioxide and free nitrogen. What will be the relative
volumes of the materials before and after the interaction (p. 210) ?
3. The density (referred to air) of sulphur trioxide vapor, as meas-
ured at a certain temperature and pressure, was 2.85. What is the
observed molecular weight ? What is the actual G.M.V. in this
case, and why does it differ from the average value (p. 199) ?
4. If the atomic weight of oxygen were taken as 100, what would
be the atomic weight of hydrogen ?
5. Knowing the molecular formulae of gaseous substances, it is
unnecessary to burden our minds with other data in regard to their
relative weights. Is hydrogen chloride (HC1) heavier or lighter than
carbon dioxide (C02) ? A G.M.V. of the former weighs H + Cl =
36.45 g., and of the latter C + 20 = 44 g. (p. 204). Hence the former
is a little lighter. Remembering that the G.M.V. of air weighs
28.955 g., we can compare the weight of any gas with that of air in
the same way.
What are the relative weights of acetylene and sulphur dioxide as
compared with air ?
6. Our equations hitherto have attempted to show only the numbers
216 INORGANIC CHEMISTRY
of units in various actions. In future nothing but molecular formulae
of free elements must be used (p. 209). For example (p. 59),
2HgO — » 2Hg + O2. Write in molecular form all the equations in-
volving gases which are found in the preceding chapters.
7. Work out the molecular formulae of all the substances in the
table of analyzed molecular weights (p. 200).
8. If the problem is to know what volume of a gas is produced
from, or combines with, a certain weight of some substance, the use of
the G.M.V. leads us most directly to the answer. Thus, what volume
of oxygen is obtainable from 60 g. of potassium chlorate ? First, we
write the molecular equation, that is, the equation showing nothing but
complete molecules of all gases concerned:
2KC108 -* 2KC1 + 302,
2x122.5 2x74.5 3x32,
and place below each formula the weight of material it represents.
Now 32 g. of oxygen occupy 1 G.M.V. or 22.4 liters. Hence 245 g.
of potassium chlorate give 3 x 22.4 liters, and therefore 60 g. give
60 x 3 x 22.4 -r- 245 liters of oxygen at 0° and 760 mm. If a differ-
ent temperature and pressure is to be considered, either the G.M.V.
or the answer must be converted by rule to the given conditions.
What volume of oxygen at 10° and 750 mm. is obtainable by heat-
ing 50 g. of barium peroxide ?
What volume of oxygen at 20° and 760 mm. is required to convert
16 g. of iron into rust (Fe208) ?
9. Write out the molecular equations for the union of alcohol vapor
(C2H60) and oxygen (p. 210) ; for the interactions of methane and
chlorine (pp. 176, 214) ; and for the burning of phosphorus (vapor)
in oxygen. Deduce the volume relations of the initial substances
and of the products at various temperatures in each case.
10. At 1700° the average molecular weight of phosphorus is 91
(p. 206). What percentage of molecules of P4 has been dissociated
into P2 ?
11. Show that, if an element has more than one equivalent weight,
the atomic weight must be some multiple of each of the equivalents
by a whole number.
12. Prove, without the use of anything hypothetical, that 16 is
preferable to 8 for the atomic weight of oxygen, because the smaller
number involves a fractional value for the atomic weight of hydrogen.
CHAPTER XIII
THE ATOMIC HYPOTHESIS
To determine the nature of chemical phenomena requires, as we
have found, very elaborate experimentation. And this has to be
followed by still more elaborate reasoning before a systematic state-
ment of the precise nature of the change can be made. Yet, when
all this has been done, we are still unable to form a clear conception
of what manner of procedure the change follows, for the details are
entirely inaccessible to observation. We should like to know pre-
cisely how chemical union is consummated, and how chemical exchange
is carried out. We should like to account for the fact that the auto-
matically adjusted proportions of the materials used in every chemical
change are entirely uninfluenced by temperature and other conditions.
Above all, we should like to know, if possible, what state of affairs
determines the employment of an individual unit weight by each
element in all its combinations. None of these questions can be
answered, however, because nothing can be seen which suggests any
answer.
As usual in cases of this kind we construct an imaginary mechanism,
a formulative hypothesis (p. 141) to account for the facts. In doing
so, we are not under the illusion that we are discovering the actual
machinery. We realize that we are simply making a sort of diagram
which will assist our thought about the thing itself. Now a slight
addition to the molecular hypothesis readily furnishes precisely what
we need.
Atomic Hypothesis. — According to the molecular hypothesis, all
matter is made up of small discrete particles, each of which has the
same composition as has the body as a whole. It is difficult, there-
fore, to avoid the conception that the different materials in each com-
pound molecule are more or less distinct entities also. Hence we make
this the basis of a new hypothesis, and attribute to these constituent
parts of molecules the properties of the chemical unit weights which
are closely related to them. These parts must move from one state of
217
218 INORGANIC CHEMISTRY
combination to another without alteration in their mass. Since they
may be restored to their free condition and recombined as often as we
choose, without impairment of their individuality, each kind must be
composed of a distinct variety of matter. These two are almost the
only qualities which the facts thus far presented justify us in attri-
buting to them. We must, therefore, carefully avoid, for the present,
the introduction of unnecessary complications by inventing for them
any other properties gratuitously.
These parts of molecules which we thus suppose to be permanent,
coherent masses, are named atoms. This word signifies objects which
are not disintegrated (Gk. aro/ios, uncut, i.e., not yet cut). The relative
weights of these imaginary masses are numerically the same as the
chemical unit weights (p. 50), and hence, in terms of this hypothesis,
we call the latter atomic weights.
It should be noted that although atoms, like molecules, are fictions,
the atomic weights, since they are measured by experimental methods,
are real. They originally received this name when Dalton, an English
schoolmaster of Manchester, succeeded in unraveling the complications
of the chemical composition of substances by the help of this very
hypothesis, and realized for the first time (1805) that the possession
by all elements of individual chemical unit weights lay at the basis of
the whole system.
We may sum up all that the facts require us to assume about atoms
by saying : Atoms are the units of -which molecules are aggregates.
Those of like kind have equal masses, and differ from those of other
kinds both in mass and the kind of material of which they are made.
The fundamentally different kinds of materials are the chemical
elements. These statements include the whole atomic hypothesis.
It is to be noted, particularly, that there are no facts in chemistry which prove
that atoms are incapable of disintegration. All we know is that in ordinary chemi-
cal changes they are transferred as wholes. It would only occasion the risk of
conflict with facts still to be discovered if we elaborated our fiction any further
than is absolutely necessary. Indeed, the phenomena of radio-activity have already
compelled us to suppose that there are particles much smaller than atoms (see below).
Without this assumption the new facts cannot be accounted for in harmony with
the molecular hypothesis.
The idea that matter is composed of small particles is a very ancient one.
Not even Dalton, however, although he used this conception continually as a
means of thought about chemical and physical phenomena, made any distinction
between atoms and molecules. This distinction was introduced later by Avogadro
(1811). Yet, without this refinement, continual thought of the behavior of matter
as consisting in transactions between small particles led Dalton to see that it was
THE ATOMIC HYPOTHESIS 219
probable that individual unit weights for each element must exist. The discovery
that they did exist soon followed. The numbers which Dalton actually gave out,
aside from the considerable experimental inaccuracies attached to them, were often
equivalents and not modern atomic weights. It was after the publication of
Dalton's ideas that Gay-Lussac discovered the law of combining volumes (1808),
and until this law was discovered there was no criterion for fixing the value of the
atomic weights. Gay-Lussac, at the end of his paper, pointed out that his dis-
covery formed an important confirmation of Dalton's views. Strange to say,
Dalton himself refused to accept Gay-Lussac's law, and so rejected the very means
by which his own principle of chemical unit weights came eventually to be acknowl-
edged as one of the foundation stones of the science. On the other hand, Dalton's
fellow countrymen and contemporaries accepted the principle of unit weights,
but rejected the atomic hypothesis by the help of which Dalton had reached them !
Thus Sir Humphry Davy called them " proportions " instead of atomic weights,
and Wollaston preferred the word " equivalents."
The Atomic Hypothesis and Chemical Change. — The change
from chemical units to atoms is so slight that the application of this
hypothesis to the description of chemical phenomena is very readily
made. The description gains in concreteness, however, from the
change. Thus, we consider iron from every source to be made of
minute portions of iron matter which are all alike in weight, and pre-
sumably also in their other properties. Similarly, all specimens of
sulphur are made of minute particles of sulphur having exactly the
same weight. The weight of a sulphur atom, however, is different
from that of an iron atom. When visible portions of the two sub-
stances unite, we conceive the operation to consist in the union of
each atom, of iron with one atom of sulphur to produce a molecule of
ferrous sulphide. The consummation of this union between multi-
tudes of pairs of the respective kinds of atoms in every second of time
results in a chemical transformation whose progress is perceptible by
the senses.
In more complex chemical changes, a further correspondence be-
tween the hypothesis and the facts comes to our notice. For example,
when sulphuric acid acts upon sodium chloride to produce hydrogen
chloride :
NaCl + H2S04 -» HC1 + NaHS04,
one part by weight of hydrogen takes the place of 23 parts by weight
of sodium, and combines with 35.45 parts of chlorine to form the
hydrogen chloride. This is the way we state the change when we
refer to measurable quantities. According to this hypothesis, the
operation consists in the repetition, millions of times over within a
220 INORGANIC CHEMISTRY
small amount of material, of the substitution of one atom of hydrogen
for one atom of sodium to form a molecule of hydrogen chloride.
The special fact which we notice is that the atom of hydrogen suffices
exactly to occupy the place of the atom of sodium. If it were too
large, then a portion of each atom of hydrogen would remain unused,
and so some free hydrogen would be produced along with the other
products. If it were too small, then some atoms of hydrogen would
be consumed in making up the material of the others to the right
amount, and hence some chlorine might fail to receive any hydrogen
in combination and escape in a free condition. Neither of these things
is observed to take place, however. The remarkable fact about this,
and all other double decompositions of the same sort, is that the little
masses of the various elements exchange places without any altera-
tions to fit the new compound being required. Hence our assumption
that the atoms are permanent, coherent wholes. This, of course, is
simply stating in terms of the atomic hypothesis the facts which
underlie the conception of unit weights. A chemical phenomenon
as we observe it, then, is imagined to consist in some systematic libera-
tion, combination, or exchange of atoms, according to a definite scheme,
and repeated many millions of times (with every molecule) in a body or
mixture of bodies.
The Atomic Hypothesis and the Quantitative ^Laws. — The
idea of atoms simply crystallizes somewhat more definitely the con-
ception of chemical unit weights. Hence, it follows of necessity that
the quantitative laws of chemical combination, out of which the latter
arose, will be found to be entirely in harmony with the atomic hypothe-
sis. The definite composition of each compound, for example, corre-
sponds to the hypothesis that each substance is made up of a specific
kind of molecules, all of which, in turn, contain the same kind and
number of atoms.
The conception of valence (p. 101)' suggests, in terms of this
hypothesis, that some atoms unite with but one other atom, habitually
(NaCl). Some, however, unite with two of the first kind (ZnCl2), or
with one other of their own kind (ZnO), still others with three atoms
of the first kind (A1C18), and so forth. In other words, it involves
the assumption that each kind of atom has a limited capacity for
holding other atoms in combination. Thus, taking the most crudely
mechanical view of the matter, we might elaborate the hypothesis by
suggesting that there is a limit to the number of points at which
THE ATOMIC HYPOTHESIS 221
atoms may be attached to one another. When one atom of chlorine
is attached to one of sodium, the combining capacity of each is ex-
hausted. When one atom of hydrogen is attached to one atom
of oxygen, one combining capacity still remains (H — 0 — ), and
can be satisfied by one more atom of hydrogen or of some other
element.
The fact that the vapor density of the elements is often greater
than that of the compounds in which they are contained, a fact which
led us (p. 194) to the conception that there are two or more unit
weights in one molecule, becomes more concrete in view of this hy-
pothesis. It is as easy to understand that there may be two or more
atoms of the same kind, as two or more of different kinds, united in
one molecule.
Thus the general scheme according to which chemical changes
occur, as well as all the quantitative relations in chemical phenomena,
are admirably described when we think of those phenomena as con-
stituted by transferences of atoms possessing definite weights.
Equations and the Atomic Hypothesis. — Since equations are
simply records of the kinds and quantities of matter taking part in
chemical changes, they require no addition to the atomic hypothesis.
The symbols, which originally represented unit weights of the various
elements, may stand equally well for atoms. Hence, in the current
language of chemistry, potassium chlorate (KC108) is composed of
one atom each of potassium and of chlorine and three atoms of oxy-
gen in each molecule. The word " atom " stands for unit weight, and
the word " molecule " for molecular weight.
Concretely, when these terms are used, the chemical equation
represents one minute specimen of the change which is taking place
throughout the whole mass. It shows enough molecules and atoms to
furnish a complete example of what, repeated millions of times, will
constitute a chemical change in a visible amount of material. There
are a few details which this point of view suggests. For example,
the equation HgO — •» Hg + 0 shows the weights, but does not include
a complete sample of the materials in the point of view of this
hypothesis. The smallest fragment of free oxygen which we can
obtain is made of two atoms. Hence two molecules of mercuric oxide
are required to furnish them. Thus the minimum number of mole-
cules which would suffice for carrying out this change on a molecular
scale is 2HgO — > 2Hg -j- 02. Similarly, since the smallest discrete
222 INORGANIC CHEMISTRY
portions of free hydrogen and chlorine contain two atoms each, we
must write the equation H., + CL, — > 2HC1 (cf. p. 209).
May other Properties, Besides Definite Mass and Specific
Material, be attributed to an Atom ? — When this question is raised,
the ground becomes less certain. A few examples will show the dif-
ficulties encountered and how far useful suggestions have been made.
Since mercury and iodine in a gaseous condition are composed of single atoms
(p. 205), all atoms must be perfectly elastic just as molecules were assumed to be
(p. 129). But this assumption is unnecessary in chemistry, for it throws no light
upon chemical behavior.
The internal constitution of atoms remained for a century rather
indefinite. As fictions, whose properties depended entirely upon
those of the chemical - units which they represented, they were not
regarded as ever undergoing any change in the amount of material
they contained. Recently, however, the peculiar emanations, which
proceed from certain minerals containing uranium and thorium, have
been discovered to be composed of minute particles (see Eadium).
It has been found possible, even, to estimate the size which these
particles must possess as compared with atoms, in order that they may
exhibit the properties which are peculiar to them. Different methods
give somewhat different values. But it appears that some of these
particles, which are called corpuscles (sometimes electrons) to dis-
tinguish them from atoms, are of the same size from whatever element
they are formed, and that their mass is about one-thousandth of that of
a hydrogen atom. The relation is of this order, at all events. Their
behavior also compels us to suppose that, even in the atom from which
they issue, they are scattered at wide distances from one another. Thus,
Professor J. J. Thomson suggests that, if the hydrogen atom were as
large as a church, the thousand or so electrons which it contains
might be compared to that number of full stops (.) scattered through-
out the whole edifice. We have thus been compelled by the facts to
add to the molecular hypothesis, first an atomic hypothesis, and then
a corpuscular hypothesis. Our atoms turn out to be clusters of
particles after all (p. 196). But the atom (chemical unit) of ordinary
chemical change is not thereby affected, however many corpuscles
may be thrust into it by other considerations.
In one direction, namely, that of accounting for the properties of
compounds, no attempt has been made to adapt the atomic hypothesis
so as to explain the facts. Thus, two hydrogen atoms in a molecule
THE ATOMIC HYPOTHESIS 223
give a gas almost insoluble in water ; two chlorine atoms, a gas which
is moderately soluble ; but one of each gives hydrogen chloride, which
dissolves in water in extraordinary quantities. So also, colorless
substances give, by chemical union, strongly colored ones, and odorless
substances, by chemical union, strongly odorous ones. Putting pieces
of iron and sulphur side by side causes absolutely no change in the
properties of either. And yet this hypothesis compels us to assume
that if the particles are made fine enough, and placed close enough to
one another, the individual properties of the constituents will entirely
disappear. Hitherto we have failed to think of any qualities which
might be attributed to the atoms in order to account for facts of this
class. Why should oxygen (02) and ozone (03) be so different in
behavior although the atomic theory hints at nothing but a substitu-
tion of three atoms for two ? What atomic properties shall account
for the difference between red and yellow phosphorus ?
Usually, in explaining these matters , we forsake atoms and employ
the conception of energy. WTe say that ozone has more energy than
oxygen, and yellow phosphorus than red phosphorus. Here the con-
ception of atoms as a means of accounting for properties is subordi-
nate to that of energy. A tendency is plainly perceptible, indeed, not
merely to supplement the atomic theory by the notion of energy, but
to allow energy to supplant it entirely, and that not only in this but
in other connections.
There are certain special peculiarities in the relation between com-
position and properties, however, to the explanation of which the
atomic theoiy has been siiccessfully adapted. Thus, certain com-
pounds precisely similar in composition, such as barium peroxide
(Ba02) and lead dioxide (Pb02), show entirely different chemical be-
havior. The former of these when treated with an acid gives hydrogen
peroxide (q.v.}, while the latter gives only water and oxygen. If it is
merely a question of composition, the substances should behave alike.
So, also, we often have two or more compounds identical in molecular
weight, and in the elements and numbers of units which they contain,
which are nevertheless totally different in physical and chemical prop-
erties. To account for this difference we have found it convenient
to suppose that, although the atoms contained in the two or more
kinds of molecules are of the same numbers and kinds, they are in
different geometrical arrangement towards one another (see Urea).
Sometimes one of these substances can be made directly from the other,
and this gives us that variety of chemical change which was named
224 INORGANIC CHEMISTRY
" internal rearrangement " (p. 15). In discussing facts like the above
in terms of the atomic hypothesis we speak of the structure of the
molecule and the constitution of the substance, and, to represent our
conclusions, we use graphic formulae (see Constitutions of oxygen
acids of chlorine, hydrogen peroxide, sulphuric acid, formic acid and
urea).
It is becoming increasingly difficult to maintain the distinction which formerly
seemed very clear between physical and chemical, phenomena. The atomic hy-
pothesis serves the purpose in a rough way for the simpler cases. If the atoms
are undisturbed, and therefore the chemical properties remain unchanged, the
substance has undergone a physical change only. If, however, the atoms are
altered in their relations to one another, either by rearrangement, or in some more
drastic manner, the change is a chemical one. Now, however, so many shades of
difference among the examples that come under the head of internal rearrange-
ment have been found, that the atomic hypothesis is applied with constantly
increasing difficulty (cf. p. 16). Whether it will be possible to apply this
hypothesis consistently in this region remains yet to be seen.
Summing this up, we find that all the suppositions the chemist
makes are : That an atom has a specific mass and consists of a specific
kind of material, and, sometimes, that the atoms in a molecule are
arranged with reference to one another in space in some definite way.
After making this formulative hypothesis, however, he knows no more
than he did before about the real mechanism of chemical change.
Is the Atomic Hypothesis a Fact. — The language of chemists has
become so saturated with the phraseology of the atomic and molecular hypotheses,
that we speak in terms of atoms and molecules as if they were objects of immedi-
ate observation. It must be reiterated, therefore, that this language is figurative,
and must not be taken literally. The atomic hypothesis provides a convenient form
of speech, which successfully describes many of the facts in a metaphorical manner.
But the handy way in which the atomic hypothesis lends itself to the representa-
tion of the characteristic features of a chemical change falls far short of constituting
a proof that atoms have any real existence. The restatements which we have
made of some of the characteristics of chemical change in this chapter depend upon
the continued reign of the atomic hypothesis. The definitions of the same things
given previously were based upon experiment and are of permanent value, so that
no question can be allowed to arise as to which of the two definitions in each case
is the more important one.
The existence of the unit weights (atomic weights) will probably appear to come
nearer to furnishing verification of the hypothesis that matter is composed of
atoms than any other fact. It may seem unthinkable that each element should use
the same proportions by weight in entering into many different combinations if it is
not actually done up in packets whose weight is represented by this number. Such
an inference, however, is not really justified. An illustration will make this clear.
If we were unable to see wheat close at hand, so as to distinguish its structure, and
THE ATOMIC HYPOTHESIS 225
our information was confined to observation from a distance and the reading of
market quotations, we should be led to infer from the latter that it was a substance
which was always done up in bushels. Yet we know that this inference would be
entirely incorrect. The substance really exists in the form of bushels at the
moment of measurement only. So there might be imagined properties, at present
unknown to us, which directed the quantitative selection of material for chemical
change and rejected the excess, without the existence of any permanent segregra-
tion into pieces of unalterable dimensions. The only bushels and ounces which
we have are in the measuring apparatus and not in the material measured ; so the
only atomic weights may be in the properties controlling chemical combination,
and not in the matter combining.
Exercises. — 1. In previous chapters our definitions have been ex-
perimental. In imitation of the definitions of the law of definite pro-
portions and of valence (p. 220), give theoretical definitions of the
following, in terms of the atomic hypothesis : Physical and chemical
phenomenon, multiple proportions, chemical unit weight, molecular
weight, element, compound, symbol, formula, equation.
2. Criticize the definitions : The atomic weight of an element is the
smallest portion of that element which takes part in chemical change.
An atom is the smallest particle that can be conceived.
3. Define all the varieties of chemical change (p. 187) in terms of
the atomic hypothesis.
CHAPTER XIV
THE HALOGEN FAMILY
THE elements to which we have so far devoted most attention have
been oxygen, hydrogen, and chlorine. If we recall the chemical prop-
erties and relations of these elements we shall recognize the fact that
they all possess very distinct individualities.
The Chemical Relations of Elements. — Hydrogen is the sub-
stance (p. 109) which unites readily with oxygen and chlorine, less
readily with other non-metals, and scarcely at all with metals. Oxy-
gen and chlorine resemble one another somewhat in the greatness of
their chemical activity and the variety of free elements with which
they are capable of uniting, but differ markedly in what we have
called their chemical relations (p. 177). The resulting compounds
belong, in fact, to quite different classes — oxygen forms oxides, chlo-
rine forms chlorides — and elements are considered similar only when
they resemble one another in chemical relations, and produce, by com-
bination with the same element, compounds having similar chemical
properties. Thus the common oxide of hydrogen, water, is a neutral
substance, and is chemically rather indifferent. The chloride of hydro-
gen in aqueous solution is a strong acid and is chemically very active.*
If all the other chemical elements differed from one another as much as
do these three, the study of the chemical elements would be tedious and
tiresome, since we should be denied the satisfaction of tracing resem-'
blances, and the elements would be incapable of classification. In
reality, however, we find that they are not incapable of being grouped
together in sets. They are classified according to the kind of sub-
stances with which they combine and the chemical nature of the prod-
ucts. In some families the resemblance is close, in others less close.
* The difference between oxides and chlorides is seen in their behavior, and will
show itself as we proceed. We have already learned, however, that oxides often
unite with water to form acids or bases (p. 119). Chlorides do not unite with
water to form new substances with marked characteristics (c/. p. 121). They
belong to the large class of compounds designated salts (q.v.).
226
THE HALOGEN FAMILY 227
The present group is of the former class, and will serve, therefore, as
a convenient beginning in the work of tracing relations between the
elements and in classifying the facts of descriptive chemistry.
The Chemical Relations of the Halogens. — The bromide (NaBr),
iodide (Nal), and, to a less extent, the fluoride (N"aF), of sodium,
resemble sodium chloride (NaCl) in appearance and behavior. From
the fact that chlorine, bromine, iodine, and fluorine are thus all
able to produce substances like salt without the presence of any other
constituents, they are known as the halogens (Gk. aA?, salt ; ycwav, to pro-
duce), and their compounds are named the halides. The halogens, as
the above formulae show, are univalent. They all form compounds
with hydrogen, and these compounds closely resemble hydrogen chlo-
ride (q-v.)- For example, they are colorless, they are gases (with the
exception of hydrogen fluoride which is a very volatile liquid), they
are very soluble in water, and their solutions are acids. Other rela-
tions will be given in a summary at the end of the chapter.
BROMINE.
Occurrence. — The compounds of chlorine, bromine, and iodine usu-
ally occur together in nature, while the compounds of fluorine are not
found in the same sources. Bromine occurs chiefly in the form of the
bromides of sodium and magnesium, in the upper layers of the natural
beds of rock salt.
Preparation. — In the chemical point of view there are three dis-
tinct ways in which bromine is made. The first of these is closely
related to the common method of preparing chlorine (p. 172). As
hydrobromic acid, unlike hydrochloric acid, is not formed extensively
in connection with any chemical industry, potassium bromide, the
most accessible compound of bromine, is treated with concentrated
sulphuric acid, and the product is oxidized with powdered manganese
dioxide in one operation. The whole action may be represented in a
single equation (see next section), as follows :
2KBr + 3H2S04 + Mn02 -» MnS04 + 2KHSO4 4- 2H20 + Br2.
Bromine being a volatile liquid, while the two sulphates are involatile,
its vapor passes off along with a little water when the above mixture
is heated. It is condensed in a worm-tube surrounded by cold water.
228 INOKGANIC CHEMISTRY
The second method of preparing bromine depends on the fact that
chlorine is a more active element and displaces bromine from combi-
nation. When, therefore, chlorine is passed into a solution of potas-
sium or sodium bromide, potassium or sodium chloride is formed and
the bromine liberated :
2NaBr + CL,-* 2NaCl + Br2. '
When the liquid is warmed, the bromine passes off along with a part
of the water, and may be condensed as before.
Aqueous solutions of soluble bromides may be decomposed by
means of a current of electricity. The bromine is set free at the
positive electrode.
The whole of the bromine used in commerce is at present manu-
factured in the first two of these ways. Two-thirds of the supply is
obtained from Stassfurt, where, after the extraction of the potas-
sium chloride from the impure carnallite (KC1, MgCl2, 6H20), the
mother-liquid is found to contain the more soluble sodium and mag-
nesium bromides in considerable quantities. The warm mother-liquor
trickles down over round stones in a tower. The chlorine is intro-
duced from below and dissolves in the liquid. The bromine is thus
liberated and passes off as vapor. A part of our supply of bromine is
obtained from the brines of Ohio, West Virginia, and Kentucky. Here
the liquid, after most of the common salt has been removed by crystal-
lization, is assayed to ascertain the quantity of bromine which it con-
tains, and is treated with the calculated amount of sulphuric acid
necessary for the action. Manganese dioxide is then added in small
quantities at a time. In Michigan the brines are treated with elec-
trolytic chlorine. The quantity produced in America in 1904, was
897,100 pounds, and was valued at $269,130.
A Plan for Making Very Complete Equations. — When an equa-
tion involves more than two initial substances or products, as does the
one given above for the first method of preparing bromine, it cannot
readily be worked out by the method formerly recommended (p. 110).
After the formulae of all the substances, on both sides, have been set
down, it is difficult to hit upon the proper numerical factors required
to balance the equation. In such cases a good plan is to select two of
the initial substances and make a partial equation showing part of the
action and including at least one actual product. Any unused units
(not constituting a product) are then set down also and treated as a
THE HALOGEN FAMILY 229
balance. Thus the first two of the substances, in the above equation,
will furnish the second of the products :
KBr + H2S04 -» KHS04 (+ HBr). (1)
Then we may represent the formation of the first of the products from
the materials which evidently must have been used up in forming it,
and set down the balance as before :
Mn02 + H2S04 -> MnS04 + H20 (+ 0). (2)
We then perceive that the last of the products might come from the
oxidation of the first balance by the second :
(2HBr) + (0) -» H20 4- Br2. (3)
The third partial equation shows that 2HBr will be needed for the
amount of 0 obtainable from Mn02, so we take enough of the material
in (1) to get the required 2HBr :
2KBr + 2H2S04 -> 2KHS04 (+ 2HBr). (1)
When we now add the real substances used and produced, as they
occur in these partial equations, and leave out the balances, which
have been adjusted so as to cancel one another, we obtain the final
equation already given for the action. It must be observed that this
subdivision of the action into parts is a purely arithmetical device,
used solely to simplify the arithmetical process of writing the equa-
tions, and is not intended to imply that the chemical change itself
follows these or, indeed, any stages. It happens that the three par-
tial equations we have used in this illustration all represent inter-
actions which can take place separately. But the arithmetical value
of the device does not depend upon this. The partial equations made
for purposes like the present one are often purely fictitious. It is
still true, however, that we are aided in the selection of partial actions
at each step by following some plausible theory as to stages for the
action which, if there were any, would be chemically conceivable.
Physical Properties. — Bromine is a dark-red liquid (sp. gr. 3.18).
It boils at 59°, forming a deep-red vapor, and even at ordinary tem-
peratures gives a high vapor pressure (150 mm. at 18°) and evaporates
quickly. When cooled it forms red, needle-shaped crystals, which melt
at 7.3°. A saturated aqueous solution (bromine-water) at ordinary
temperatures contains about three parts of bromine in one hundred of
230 INORGANIC CHEMISTRY
water. The element is much more soluble in carbon bisulphide,
alcohol, and other organic solvents. Its vapor density up to 750° is 160
(oxygen = 32).
Bromine (Gk. /3ptoyu.os, a stench) has a most pungent odor. It has
a very irritating effect on the mucous membrane of the nostrils and
throat. If spilled upon the hands it has a most destructive action
upon the tissues.
When free bromine is added to starch emulsion no special change
in tint is observable unless a very large amount of the element is used
(see Iodine).
Chemical Properties. — The molecules of bromine are much less
stable than those of hydrogen, oxygen, or nitrogen. The atomic
weight of bromine is 79.96, so that in the form of vapor it has two
atoms in a molecule, and is represented by the formula Br2. At
1050° the vapor density is 150.5, and dissociation into Br has begun.
Like chlorine, it forms an unstable hydrate with water.
Bromine unites directly with hydrogen. The mixture of the gases
is not explosive, and the union is much slower than in the case of
chlorine. In presence of finely divided platinum the speed of the
action may be considerably increased.
Bromine forms compounds directly, both with non-metals, like phos-
phorus and arsenic, and with most of the metals. Towards unsatu-
rated substances and organic compounds it behaves like chlorine (q.v.}.
In all cases the interaction is less violent than when chlorine is used,
and the element, is displaced from combination with hydrogen and
with the metals by free chlorine.
Potassium bromide is employed in making photographic plates and
in medicine. Bromine is required in large quantities in the manufac-
ture of intermediate products used in the preparation of organic
dyes.
HYDROGEN BROMIDE.
Preparation. — It might be expected that the most convenient way
of producing this compound would be similar to that used in prepar-
ing hydrogen chloride, namely, by the action of concentrated sulphuric
acid upon some common bromide such as potassium bromide (KBr +
H2S04^±HBr + KHS04). We find indeed that at first a colorless
gas is given off, which fumes strongly in the air just like hydrogen
chloride, and is the required substance. Almost immediately, how-
THE HALOGEN FAMILY 231
ever, this gas acquires a yellow and then a brown tinge, and we dis-
cover that free bromine is being produced at the same time. If we
examine the gas still further, we recognize also the presence of sulphur
dioxide. It is impossible, therefore, to produce hydrogen bromide
free from those two impurities by this action.
The origin of the bromine and sulphur dioxide which complicate
this chemical change may readily be traced. Hydrogen bromide is
less stable than hydrogen chloride, and its hydrogen can more easily
be removed by the action of substances containing oxygen. In this
case the sulphuric acid acts as the oxidizing agent, yielding oxygen,
sulphur dioxide, and water (H2S04 — > 0 + S02 + H20). Thus the
two extra gaseous products are seen to be formed by a change pro-
ceeding parallel with the main action :
2HBr + H2S04 -» 2H20 + SO2 + Br2.
The simultaneous occurrence, in this fashion, of two more or less
independent actions is not uncommon. The speeds of such actions
may be differently affected by temperature. Thus, here, the second
action seems to become more extensive as the temperature rises (see
Chap. xvi). Since all acids decompose all salts more or less, by use of
an acid which does not give up its oxygen so readily, such as phos-
phoric acid, pure hydrogen bromide may be obtained (KBr -f- H3P04 — >
HBr -f KH2P04). The small solubility of the Salt in concentrated
phosphoric acid retards the interaction (p. 179) and makes the evo-
lution of the gas very slow, however.
Pure hydrogen bromide is prepared by hydrolysis (p. 181) of phos-
phorus tribromide. When bromine and phosphorus are mixed, a vio-
lent union of the two elements takes place, producing phosphorus
tribromide (PBr3). This substance, which is a colorless liquid, is in
turn broken up with great ease by water, producing phosphorous acid,
which is not volatile, and hydrogen bromide :
, Br HOH
P - Br + HOH -» P(OH)3 + 3HBr.
x Br HOH
In practice, those two actions are carried on simultaneously. To dimin-
ish the vigor of the interaction,, red phosphorus is taken instead of
yellow, and is mixed with two or three times its weight of sand in a
flask (Fig. 70). A small quantity of water is added. Excess of
water must be avoided, as the hydrogen bromide produced is extremely
232
INORGANIC CHEMISTRY
soluble, and would therefore be retained in the flask instead of being
disengaged as gas. The bromine is placed in the dropping funnel,
and admitted, a little at a
time, to the mixture. The
gas produced is passed
through a (J-tube containing
glass beads mixed with red
phosphorus. The latter com-
bines with any bromine
which may have escaped
chemical change and have
been carried along with the
gas. The second U-tube,
containing water, may be
attached when a solution of
the gas is required.
FlG- 70- Physical Properties.
— Hydrogen bromide is a
colorless gas with a sharp odor. It is two and a half times as heavy as
air. It is easily reduced to the liquid condition. It is exceedingly sol-
uble in water, and in contact with moist air condenses the water vapor
to clouds of liquid particles. When distilled, the solution in water
behaves like that of hydrogen chloride (p. 182). It loses mainly either
water or hydrogen bromide, according as it is dilute or exceedingly
concentrated, until an acid of constant boiling-point (126° at 760 mm.
pressure), containing 48 per cent of hydrogen bromide, passes over.
Hydrogen bromide, whether in the gaseous condition or in the
liquefied form, is a nonconductor of electricity.
\
Chemical Properties. — The chemical properties of hydrogen
bromide are similar to those of hydrogen chloride (p. 183). It is some-
what less stable, and dissociation into its constituents begins to be
noticeable at 800°. When free from water, it is not an acid (see
below), and is not very active chemically, although it behaves towards
some metals much like hydrogen chloride. When the solution in
water is strongly cooled, crystals of a definite hydrate (HBr, 2H20),
corresponding to that of hydrogen chloride, are obtained. When the
gas is mixed with chlorine, hydrogen chloride and free bromine are
instantly produced, and much heat is evolved by the change, 2HBr-f-
THE HALOGEN FAMILY 233
C12 — > 2HC1 -f Br2 . The heat produced by the union of hydrogen and
bromine vapor is 12,100 calories. This is much less than the amount
produced by the union of chemically equivalent quantities of hydro-
gen and chlorine (22,000 calories). When chlorine displaces bromine
from hydrogen bromide, the heat evolved is found to be the difference
between these two numbers. Using the rule of constant heat summa-
tion (p. 78), we write equation (2) so that HBr is on the same side
with Cl (with which it interacts), and the products of the equation
required (HC1 and Br) are both on the right:
H + C1 ->.HC1 + 22,000 cal. (1)
HBr -> Br + H - 12,100 cal. (2)
Adding, HBr + Cl -> HC1 + Br + 9900 cal.
The 12,100 calories are produced by the union of gaseous bromine
with hydrogen, and the final result is, therefore, that for the produc-
tion of gaseous bromine. If the heat- of formation of liquid bromine
is required, the latent heat of vaporization of bromine (7296 calories),
which will be evolved when the element condenses, must be added.
Chemical Properties of Hydrobromic Acid. — The solution of
the hydrogen bromide in water is an active acid (cf. p. 111). It con-
ducts electricity extremely well. In contact with metals, oxides of
metals, and hydroxides of metals, it behaves exactly like hydrochloric
acid (p. 185). In the first case, hydrogen is set free and the bromide
of the metal produced. In the other two cases, water and the bro-
mides of the metals are produced. Oxidizing agents set bromine free
from hydrobromic acid, and the only difference as compared with
hydrochloric acid is that less powerful oxidizing agents can produce
this result. Chlorine dissolved in water displays bromine from hydro-
bromic acid and from soluble bromides with ease.
IODINE.
Occurrence. — Iodine, like bromine, occurs in sea-water, although
it is present in much smaller quantities. Fortunately, certain species
of sea-weed seem to have the power to remove it from water, and use
it as a constituent in complex organic compounds which they contain.
In the case of certain species, the ash of the sea-weed is said to con-
tain as much as two per cent, or even more. The other chief source
of iodine is in Chili saltpeter (NaN03), in which it is present in the
234 INORGANIC CHEMISTRY
form of small proportions of sodium iodate (NaI03) and sodium
iodide. Of recent years the quantity obtained commercially from this
source has greatly increased, while that from sea-weed has diminished.
It is worth noting that while the total amount of iodine found in the human
body is small, the proportion contained in the thyroid gland is considerable. A
complex organic substance, known as iodothyrin, has been extracted from sheep's
thyroids. It is administered with marked success in certain diseases, such as
cretinism, which are associated with very small development of this gland.
Preparation. — In factories where the iodine is extracted from
sea-weed, the latter is first burned, either roughly in hollows in the
ground, or more carefully in specially constructed ovens. A great
deal of the iodine seems to be lost by volatilization in this process,
but the ash which remains, and which is known in Scotland as kelp
and in Normandy as varec, still contains from 0.5 to 1.5 per cent of
sodium iodide. The ash is treated with water, and the solution is evap-
orated so as to permit the deposition of the sodium chloride and sodium
sulphate which it contains. The sodium iodide, being very soluble,
remains in the mother-liquor. This is then treated with manganese
dioxide and sulphuric acid. The quantity of manganese dioxide is
carefully measured so as to be just sufficient to set free the iodine
contained in the liquid without proceeding farther to the liberation of
the chlorine which it contains in much larger amounts. When the
mixture is heated, the iodine passes off in the form of vapor, and is
condensed in a suitable receiver. The action (cf. pp. 227, 229) is :
MnO2 + 3H2SO4^MnS04 -f- 2NaHS04 + 2H20 + I2.
In France the treatment is similar, excepting that chlorine is used
to liberate the iodine in the last stage (2NaI + C12— > 2NaCl + I2).
The quantity is adjusted so that excess may not be employed. This
is to avoid the formation of a compound of chlorine and iodine, which
easily arises by direct union of the elements. The iodine, being insol-
uble, forms a dense precipitate, and, when the liquid is pressed out, it
remains behind in the form of a paste. Electricity could also be used
for the decomposition of this mother-liquor. The iodine is set free at
the positive electrode, while the metals pass to the other. This
process does not, however, seem as yet to have found actual applica-
tion in commerce.
In all cases the iodine is subjected to purification before being sold.
This is carried out by distilling it with a little powdered potassium
THE HALOGEN FAMILY 235
iodide. It condenses in the solid form directly, in glittering black
plates (sublimed iodine). The distillation of a solid body, when a
condensation takes place directly to the solid form, is spoken of as
sublimation.
Physical Properties. — Iodine (Gk. toetS^s, like a violet) is a
solid substance (sp. gr. 5), exhibiting large crystalline plates of rhombic
form. It melts at 114°, and boils at 184°. The vapor has at first
a reddish- violet tint, and on being more strongly heated becomes deep
blue.
Iodine is much less soluble in water than are the other halogens,
and the solution has a scarcely perceptible brown tint. At ordinary
temperatures one part of iodine dissolves in about 5000 to 6000 parts
of water. It is much more soluble in carbon disulphide (p. 156) and in
chloroform, in which it gives violet solutions. In alcohol it gives a
solution which is brown. The brown color is attributed to the fact
that in alcohol the iodine is in a condition of feeble combination and
not simply in solution. An aqueous solution of potassium iodide,
hydrogen iodide, or any other iodide, has likewise the power to take up
large quantities of iodine. In these cases, however, the formation of
definite compounds (such as, KI + I2 <=> KI8), by a reversible action,
accounts for the amount of iodine which appears to be dissolved.
The behavior of free iodine towards starch forms a distinctive test
for both substances (cf. p. 99). When starch is treated with boiling
water it passes into a state of fine suspension, and the liquid may be
filtered without removal of all the starch. When traces of iodine, con-
tained, for example, in the pale-brown aqueous solution, are added to
this filtered starch emulsion, a deep-blue color is produced. The same
action is used as a test for starch. This blue substance is not a chemi-
cal compound, but a solution of the iodine in the solid starch which is
suspended in the water. That the color becomes so much stronger is in
harmony with the fact that dissolved bodies frequently confer a tint
different from their own upon their solutions.
Chemical Properties. — The molecular weight of iodine, ascer-
tained by weighing the vapor at a temperature above the boiling-point,
is 254.8. The atomic weight being 126.97, the molecule contains two
atoms. This value for the density remains unchanged when the
measurement is made at temperatures up to about 700°. Beyond this
point, however, the vapor diminishes in density more rapidly than
236 INORGANIC CHEMISTRY
Charles' law would lead us to expect, and at 1700° the molecular weight
has fallen to 127. As the vapor is heated, a larger and larger propor-
tion of the molecules is broken up, until the decomposition has become
complete. As in all cases of dissociation, when the vapor is cooled the
atoms recornbine to form molecules. This chemical action is interest-
ing, since it is the most notable case in which we encounter both the
monatomic and the diatomic forms of the same element. The heat
given out when the atoms reunite to form the molecules is very con-
siderable (21 <=» I2 + 28,500 cal.), indicating that the chemical union
of two atoms of identical nature may be as vigorous as that of two
atoms of different chemical substances. The monatomic and diatomic
forms of iodine should be distinct chemical substances, and if the
investigation of the behavior of the former were not hampered by the
very high temperature at which alone it exists, it would doubtless be
found to exhibit different chemical properties.
Iodine forms no hydrate with water. It exhibits little tendency to
unite with hydrogen.. Iodine unites directly with a number of ele-
ments, including some non-metals and the majority of the metals.
When phosphorus is presented in the yellow form, the action takes
place spontaneously without the assistance of heat. Both chlorine
and bromine displace iodine from combination with hydrogen and the
metals (2HI + Br2 — » 2HBr + I2). The action may be brought about
either with the substances in dry form or with their aqueous solutions.
Iodine in water, like chlorine in water, constitutes an oxidizing
agent, although the former is much the feebler of the two. It con-
verts, however, sulphurous acid (q.v.) into sulphuric acid. The
action, so far as the iodine and the water are concerned, is shown by
the partial equation
0).
This action does not take place, however, excepting in presence of a
body capable of assisting in the reduction of the water and of uniting
with the oxygen. In analytical chemistry a solution of iodine in
potassium iodide containing a known proportion of iodine (a standard
solution) is used for estimating the quantity of anoxidizable substance
present in a given specimen. The amount of oxidizable substance
present is measured by' the quantity of the standard iodine solution
which can be decolorized and suffer removal of its iodine. This
method is known as iodimetry.
Iodine and its compounds are much used in the arts and medicine.
THE HALOGEN FAMILY 237
Iodine is applied, in the form of an alcoholic solution ("tincture of
iodine"), for the reduction of some swellings. It is required in mak-
ing iodoform (CHI8), and the iodides of potassium, rubidium, and
sodium, which are used in medicine. Potassium iodide is likewise em-
ployed in the manufacture of photographic plates. Large quantities of
iodine find application in the color industry for making intermediate
products needed in the manufacture of aniline and other organic dyes.
HYDROGEN IODIDE.
Preparation. — The direct union of hydrogen and iodine cannot be
employed in preparing pure hydrogen iodide. The union takes place
slowly, and, since the action is markedly reversible (cf. p. 176), always
remains incomplete : H2 + I2 +± 2HI. In presence of finely divided
platinum the union, so far as it is capable of proceeding, may be
hastened, but the amount of the compound formed is not greater. At
448°, for example, 79 per cent of the constituents unite, and the
remainder is unaffected either by longer heating or by platinum. At
other temperatures the proportion is different, but at none is the
action complete. Hydrogen and iodine unite with little vigor because
the union is not accompanied by a sufficiently marked liberation of
energy (cf. HC1, p. 233). Thus at 18° the action is really endothermal
(— 6100 calories for HI), and at 400°, although the action is exother-
mal, only 535 calories of heat are set free.
The action of concentrated sulphuric acid upon potassium or sodium
iodide is equally inapplicable. In this case, as in that of hydrogen
bromide (p. 231), the hydrogen halide reduces the sulphuric acid, and
much free iodine is formed. Here, on account of the greater instabil-
ity of hydrogen iodide and the greater ease with which it parts with
its hydrogen, the reduction of the sulphuric acid is much more com-
plete, the primary product being apparently hydrogen sulphide. The
actions, which proceed simultaneously (p. 231), are :
KI + H2S04 <=> HI + KHS04 andH2S04 + SHI -> H2S + 4H20 + 4I2.
As soon as the heat produced by the action has raised the temperature
sufficiently, hardly any of the hydrogen iodide escapes oxidation.*
* When much sulphuric acid is used, sulphur dioxide and free sulphur are
formed also. This is in consequence of a secondary action of the hydrogen sulphide
on the sulphuric acid, and of the sulphur dioxide so formed upon the excess of
hydrogen sulphide (see Hydrogen sulphide) :
H2S + H2S04 -+ 2H20 + S + S02, SO2 + 2H2S -> 3S + 2H2O.
238 INORGANIC CHEMISTRY
Finely powdered sodium iodide and concentrated phosphoric acid
(cf. p. 179), when mixed and warmed, give pure hydrogen iodide
(Nal + H3P04<=»HIf + NaHjjPC^). This action was formerly used
in preparing the gas.
The best method is one similar to that described under hydrogen
bromide. Phosphorus and iodine unite directly to form PI8. This is
a yellow solid which is violently decomposed by water and gives phos-
phorous acid and hydrogen iodide :
PI8 + 3H20 -» P(OH)8 + SHI.
If excess of water, which dissolves hydrogen iodide, is avoided, the
latter goes off in a continuous stream in a gaseous condition.
Still another method of making hydrogen iodide is frequently
employed when a solution of the gas in water is required, and not the
gas itself. Powdered iodine is suspended in water, and hydrogen sul-
phide gas (q.v.) is introduced through a tube in a continuous stream.
The iodine dissolves slowly in the water, I2 (solid) <=± I2 (diss'd), and
acts upon the hydrogen sulphide, which likewise dissolves, H2S (gas) ^
H2S (diss'd). Sulphur separates in a fine powder, S (diss'd) ^ S
(solid), and hydrogen iodide is formed in accordance with the
equation :
This action takes place, however, only in presence of water, although
the water does not appear in the equation. The solution is freed from
the deposit of sulphur by filtration, and may be concentrated to 57,
per cent of hydriodic acid by distilling off the water.
The theory of the last method is worthy of attention. It will be
seen that while iodine has little tendency to unite with free hydrogen,
it is here able to decompose a compound containing hydrogen, in order
to secure this element. It is enabled to do this by the fact that the
very large amount of heat given out by the mere solution of hydrogen
iodide in water converts the action, which would otherwise be eiidother-
mal, into an exothermal one. In the absence of water, the reverse of
the above action takes place with ease. In presence of water, however,
the great heat of solution (p. 164) of the hydrogen iodide ( -f- 19,-
200 cal.) more than balances the heat absorbed by the chemical change,
and the action as a whole takes place with evolution of heat (see, also,
Preparation of hydrogen sulphide).
THE HALOGEN FAMILY 239
Physical Properties. — Hydrogen iodide is a colorless gas with
a sharp odor. Its molecular weight is 128, and it is therefore much
heavier than air, the average weight of whose molecules is 28.955
(p. 193). It is a nonconductor of electricity, both in the gaseous
and in the liquefied conditions. It is exceedingly soluble in water, so
that at 0° ten grains of water will absorb ninety grams of the gas, giv-
ing a 90 per cent solution. The behavior of this solution is similar to
those of hydrogen chloride and hydrogen bromide (cf. pp. 182, 232).
The mixture of constant boiling-point distils over at 127° (at 760 mm.),
and contains 57 per cent of hydrogen iodide.
Chemical Properties* — Hydrogen iodide is the least stable of the
hydrogen halides. When heated it begins visibly to decompose into
its constituents at 180°. On account of the ease with which it parts
with the hydrogen which it contains, it can be burned in oxygen gas,
4HI + 02 — > 2H20 -f 2I2. When the gas is mixed with chlorine, a
violent chemical change, accompanied by a flash of light, occurs, the
iodine is set free, and hydrogen chloride is produced, Cl^ + 2HI— >
2HC1 + I2. Bromine vapor will similarly displace the iodine from
hydrogen iodide.
Chemical Properties Of Hydriodic Acid. — In most respects
the aqueous solution behaves exactly like hydrochloric and hydro-
bromic acid. With oxidizing agents, for example, such as manganese
dioxide, it gives free iodine, just as the others give free chlorine and
'bromine respectively. Here, however, the oxidation is so much more
easily carried out, that it is slowly effected by atmospheric oxygen,
so that hydriodic acid left exposed to the air gradually becomes brown
(02 + 4HI — » 2 H90 + 2I2). The free iodine which is produced, re-
mains dissolved in the hydrogen iodide, probably in the form of a
compound HI8. Finally, however, the free iodine as its quantity
becomes greater, and that of the hydrogen iodide smaller, is deposited
in crystalline condition. On account of the ease with which hydriodic
acid parts with its hydrogen, it is frequently used in chemistry as a
reducing agent.
Although the dry gas is not an acid, the solution has all the ordi- *
nary properties of this class of substances (cf. p. 92). The hydrogen
may be displaced by metals like zinc and magnesium. The acid inter-
acts with oxides and hydroxides, forming iodides and water.
240
INORGANIC CHEMISTRY
FLUORINE.
The discussion of this element should logically have preceded that
of chlorine, since it is of all the members of the halogen family the
most active. Chlorine was taken up first, however, because its com-
pounds are more familiar. Fluorine is found in combination in nature.
It occurs chiefly in the mineral fluorite (calcium fluoride, CaF2) and in
cryolite, a double fluoride of aluminium and sodium (A1F3, 3NaF).
Preparation. — When a solution of hydrofluoric acid is heated
with manganese dioxide, oxidation does not occur and free fluorine is
not produced. Until very recently all efforts to isolate the element
failed. It was perfectly understood that the reason of these failures
lay in the greater chemical activity of fluor-
ine, which made it more difficult of separa-
tion from any state of combination than
the other halogens. Its preparation was
finally achieved by Moissan (1886) by the
decomposition of anhydrous hydrogen fluor-
ide, which is liquid below 19°, by means
of electricity. The apparatus (Fig. 71) is
made of copper, which, after receiving a
thin coating of the fluoride, is not further
affected. To reduce the tendency to chem-
ical union, the whole is immersed in a bath
giving a temperature of — 23°. The elec-
trodes are made of an alloy of platinum and
iridium, which is the only substance that
can resist the action of the fluorine when
freshly liberated by the electric current.
Hydrogen fluoride, like other hydrogen
halides, is a nonconductor of electricity, and
a small quantity of potassium fluoride has to be added to enable the
current of electricity to pass. The fluorine is set free at the positive
electrode, and hydrogen appears at the negative. The U-tube is closed
after the introduction of the hydrogen fluoride by means of blocks made
of calcium fluoride, which is naturally unable further to enter into com-
bination with fluorine. For the reception and examination of the
fluorine gas, other copper tubes can be screwed on to the side necks of
the apparatus, and, when necessary, small windows of calcium fluoride
FIG. 71.
THE HALOGEN FAMILY 241
can be provided. It has been found that fluorine dried with extraor-
dinary precautions is without action on glass.
Physical Properties. — Fluorine is a gas whose color is like that
of chlorine, but somewhat paler. Its density has not been measured
with great exactitude, but the value obtained indicates a molecular
weight of 38, showing that there are two atoms in the molecule (the
atomic weight is 19). The gas is the most difficult of the halogens to
liquefy. The liquid boils at - 186°.
Chemical Properties. — Fluorine unites with every element, with
the exception of oxygen, and in many cases does so with such vigor
that the union begins spontaneously without the assistance of external
heat. Dry platinum and gold are the elements least affected. It
explodes with hydrogen at the ordinary temperature, without the
assistance of sunlight. Fluorine displaces oxygen from water instan-
taneously and gives ozone {q.v.}. On the introduction of a drop of
water into a tube of fluorine the vessel is filled with the deep-blue
gas, 3F2 + 3H20 -+ 3H2F2 -f O3.
The chlorine in hydrogen chloride is displaced by fluorine as easily
as chlorine in turn displaces bromine or iodine.
HYDROGEN FLUORIDE.
Preparation. — Pure, dry hydrogen fluoride is best made by heat-
ing potassium hydrogen fluoride, 2KHF2 ^± 2KF + H2F2| . For
ordinary purposes, however, the preparation of an aqueous solution is
the ultimate object. Usually powdered calcium fluoride is treated
with concentrated sulphuric acid, and the mixture distilled in a plat-
inum retort :
CaF, + H2S04<=» CaS04 + H2F2|.
The hydrofluoric acid passes over and is caught in distilled water.
The aqueous solution thus obtained has to be kept in vessels made of lead,
rubber, or paraffin, as glass interacts with the acid with great rapidity
(see below).
Physical Properties. — Hydrogen fluoride is a colorless liquid,
boiling at 19.4°. It mixes freely with water, and, on distillation, an
acid of constant boiling-point (120° at 760 mm.) containing 35 per
cent of hydrogen fluoride is obtained. The vapor density of the
242 INORGANIC CHEMISTRY
hydrogen fluoride between its boiling-point and 30° corresponds to a
molecular weight of 40, and the formula should therefore be H2F2.
Above this temperature the vapor becomes lighter (p. 207), and, when
88° is reached, the molecular weight has fallen to 20, corresponding
to the formula HF.
Association. — Many compounds resemble hydrogen fluoride in
seeming to consist of molecules which are multiples of the simplest
possible. This is spoken of as indicating a tendency to association.*
Thus, sulphuric acid and nitric acid in the liquid condition are com-
posed of more complex aggregates than H2S04 and H]ST08. Even water
is largely (H20)a or even (H20)3, although the vapor is H20. In such
cases dissociation into the simpler molecules takes place gradually as
the temperature is raised. Sometimes this doubling up of molecules
is called polymerization.
In cases of association the observed molar weights, at low tem-
perature, are multiples of the smallest possible molar weight required
for the simplest formula (in above instance, HF). But many sub-
stances naturally possess formulae which are multiples of the simplest
without showing, as the temperature is raised, any tendency to pro-
gressive dissociation into the corresponding simplest molecules. Thus,
acetylene (p. 204) is C2H2 at all temperatures, and acetic acid (q-v.),
although it is associated to C4H804 at its boiling-point, never becomes
simpler than C2H402 at any temperature.
Chemical Properties of Hydrofluoric Acid. — Metals like zinc
and magnesium interact with hydrofluoric acid with evolution of
hydrogen. The action is less violent than with other halogen acids.
The acid interacts with oxides and hydroxides, forming fluorides.
The chief difference in this respect which it exhibits, when compared
with the other halogen acids, is one which we should expect from its
formula, H2F2. We may displace either one or both the hydrogen
atoms in the molecule with a metal. Thus, one of the commonest salts
of hydrofluoric acid is potassium hydrogen fluoride KHF2 mentioned
above, KOH + H2F2 -> KHF2 + H2O. In this respect the acid re-
sembles sulphuric acid and other acids containing more than one
replaceable hydrogen unit. Salts in which a portion of the acid
hydrogen still remains undisplaced are spoken of as acid salts.
* The elementary substances show similar behavior (pp. 205, 206).
THE HALOGEN FAMILY 243
The most remarkable property of hydrofluoric acid depends on
the great tendency which fluorine has to unite with silicon, forming
the gaseous silicon tetrafluoride. Glass, which is commonly made
by fusing together sodium carbonate, calcium oxide, and sand (silicon
dioxide), is a mixture of silicates of calcium and sodium, and is
rapidly decomposed by hydrofluoric acid. The nature of the change
is shown by the two following equations :
CaSi08 +3H2F2-» SiF4 +- CaF2 + 3H20,
Na2Si03 +3H2F2-+ SiF4 + 2NaF + 3H20.
All other silicates are decomposed according to the same plan. The
silicon tetrafluoride passes off. The fluorides of calcium and sodium
are solid and crumble away or dissolve. Thus the glass is completely
disintegrated. The vapor of hydrofluoric acid, generated in the way
described above from calcium fluoride in a lead dish, is used for etch-
ing glass. The surface of the glass is covered with paraffin to protect
it from the action of the vapor, and with a sharp instrument portions
of this paraffin are removed where the etching effect is desired. The
vapor gives a rough surface where it encounters the glass. The
aqueous solution, which may also be employed, makes smooth depres-
sions on the surface. The acid is also used in the analysis of minerals
containing silicates, which frequently are not attacked by other acids.
THE HALOGENS AS A FAMILY.
It may be useful here to bring together some of the facts in regard
to the halogens and their compounds by way of showing more clearly
how far they resemble one another, and in what ways they differ.
The most noticeable fact is that, if we arrange them in order in respect
to any one property chemical or physical, the other properties will
be found to place them in the same order. Thus, if we consider
(1) the physical properties, we find that the color deepens as we pass
from fluorine through chlorine and bromine to iodine. The specific
gravities of the elements increase in the same order. The volatility
of the elements decreases in the same way — fluorine being the hard-
est to liquefy, while iodine is a solid and boils at a fairly high tem-
perature. (2) In their chemical behavior, when, for example, they
unite with the metals and hydrogen, the vigor of the action is greatest
with fluorine and diminishes progressively until we reach iodine.
We shall see later that the affinity for oxygen, on the other hand,
increases as we pass from fluorine to iodine.
244 INORGANIC CHEMISTRY
(3) The relations of these elements in combination show that they
are all univalent in respect to union with hydrogen and metals. In
their oxygen compounds (q.v.~), however, they frequently exhibit a
higher valence. The compounds which they form with any one ele-
ment are usually very similar to one another. All the hydrogen com-
pounds, for example, become acids when dissolved in water. The most
noticeable lack of harmony in this groiip is observed when we consider
the solubilities of the corresponding compounds. Thus, the potassium
salts are all soluble in water. Silver chloride, bromide, and iodide
are almost insoluble, the amount dissolved decreasing in that order.
Silver fluoride, however, is quite soluble. Calcium chloride, bromide,
and iodide are all very soluble, while calcium fluoride is almost com-
pletely insoluble.
It will be noted that the order in which the elements are thus
placed by consideration of most of their properties is the order of
increasing atomic weights (see Periodic system).
COMPOUNDS OF THE HALOGENS WITH EACH OTHER.
We have incidentally mentioned the fact that iodine unites with
chlorine to form a definite compound. In reality there are two such
compounds. The most familiar is a red crystalline substance having
the composition IC1. Another compound, IClg, is made by the use
of excess of chlorine. Iodine unites with bromine to form the com-
pound IBr, while a compound with fluorine, to which the composition
IF5 has been assigned, is supposed to exist. None of these compounds
are particularly stable, and some of them decompose very easily. It
is frequently stated that elements which resemble one another chemi-
cally show little tendency to chemical union. Yet in the case of
IBr, for example, the tendency to decompose into the elements
(2IBr — > I2 + Br2) must be interpreted as meaning that the iodine
and bromine prefer to unite with themselves to form the molecules
I2 and Br2 rather than with one another. In view of this the above-
remark loses some of its point, for an element certainly resembles itself
more than it does any other, and the compounds C12, H2, etc., are
amongst the most stable that we know.
Exercises. — 1. What impurities is commercial iodine likely to
contain ? In what way does heating with potassium iodide (p. 234)
free it from these ?
THE HALOGEN FAMILY 245
2. Classify all the chemical actions in this chapter according as
they belong to one or other of the ten kinds (p. 187).
3. Explain the condensation of atmospheric moisture by the hydro-
gen halides (cf. p. 162).
4. Tabulate, more fully and specifically than is done in the section
on "The Halogens as a Family," (a) the physical properties, (b) the
chemical properties, (c) the chemical relations, of the members of
this group.
5. Construct the equation on p. 234 by the use of partial equations
as in the example on p. 229.
6. Using the method given on p. 229, construct a single equation
for the formation of iodine, water, and hydrogen sulphide directly from
potassium iodide and sulphuric acid.
CHAPTER XV
CHEMICAL EQUILIBRIUM
IN spite of its formidable title, this chapter will introduce nothing
novel. Its purpose is to collect together and organize more definitely
a number of scattered facts and ideas which have already come up in
various connections. On this account, however, it will be all the more
necessary for the reader to refresh his remembrance of these facts
and ideas by re-reading the pages to which reference may be made.
Reversible Actions. — In discussing the union of hydrogen and
iodine (p. 237), it was stated that the progress of the action ceases
while yet a large amount of both the substances necessary for its main-
tenance still remains available. Now the materials left over are pre-
sumably no less capable of uniting than the parts which have already
united. The solution of this mystery lies in the fact (p. 239) that
decomposition of the compound can begin at 180°, and therefore takes
place actively at 448°. Hence the product of the union must begin
to dissociate, in part at least, as soon as any of it is formed. Thus
two changes, one of which undoes the work of the other, must go on
simultaneously. In consequence of this, neither can reach comple-
tion. As we should expect, experiment shows that it makes no differ-
ence whether we start with the elements or with the compound : the
proportions of the materials found in the tube, after it has been heated
for a sufficient length of time, are in both cases the same. A general
statement may be founded on facts like this, to the effect that a chem-
ical action must remain more or less incomplete when the reverse
action also takes place under the same conditions (cf. p. 35). Two
arrows pointing in opposite directions are used in equations repre-
senting reversible changes : *
H2 + ],<=» 2HI or 2HI <=> H2 + I2.
It will be observed that representing reversible actions by equations involves
a departure from the original meaning of an equation. Thus at 448°, 80 per cent
* The reader must avoid the idea that a reversible action is one which goes to
completion, and then runs back to a certain extent. This conception would be
contrary to the fact, and opposed to the principles of energetics, as well as inex-
plicable by the kinetic hypothesis.
246
CHEMICAL EQUILIBRIUM 247
of the substances are in the form HI and 20 per cent in the uncombined state:
(80%)2HI<=±H2 + I2(20%).
In other words, the amounts of matter on the two sides are not equal. Each side,
taken separately, shows correctly the proportions used in the interaction for which
it stands, however. Hence the equation in a reversible action professes to show
quantitatively the change which would occur if each of the two opposed actions it
includes were to be allowed separately to proceed to completion.
The following are some other examples of actions of the same
kind : (1) The interactions of sulphuric acid and sodium chloride
(p. 179), and (2) of chlorine and water (p. 176), which were fully
discussed at the time; (3) that of equivalent amounts of iron and
water, or of magnetic oxide of iron and hydrogen (p. 110), which does
not proceed to completion in either direction when either set of mate-
rials is sealed up in a tube and heated ; (4) the behavior of barium
peroxide (p. 63), of hydrates (p. 120), of iodine vapor (p. 205), of
water vapor (p. 119), of sulphur vapor (p. 205), of phosphorus vapor
(p. 205), and of mercuric oxide (pp. 62, 77).
When the action is one which is reversible, but, under the circum-
stances being discussed, proceeds far towards completion in one direc-
tion, the arrow will be modified to indicate this fact:
C12 + H20 *=5 HC1 + HC10 (p. 176).
When this relative completeness is due to precipitation or volatiliza-
tion, the fact will be indicated by vertical arrows :
NaCl + H2S04±=>NaHS04 + HC1 f (p. 178).
NaCl j+ H2S04<=>NaHS04 + HC1 (p. 179).
All chemical actions do not belong to this class. Many proceed
uninterruptedly to exhaustion of one, or all, of the ingredients. For
example, equivalent amounts of magnesium and oxygen combine com-
pletely (2Mg + 02 — > 2MgO). Here, however, the product is not
decomposed even at the white heat produced by the vigor of the union.
Indeed, magnesium oxide cannot be decomposed, and the action re-
versed, at any temperature we can command. The other complete
actions are so because they are likewise irreversible.
Chemists now look upon reversibility as being the normal property of all chem-
ical changes. They consider the lack of obvious incompleteness as being due to
the very small degree to which the reverse action can occur under the given con-
ditions, rather than to the entire absence of any reverse action.
Kinetic Explanation. — Restating thefce facts in. terms of the
kinetic hypothesis will enable us to reason more clearly about this
248 INORGANIC CHEMISTRY
variety of chemical change. Suppose we start with the materials rep-
resented on one side of such an equation. The molecules of these
materials will encounter one another frequently in the course of their
movements. In a certain proportion of these collisions the chemical
change will take place. In the earliest stages there will be few of the
new kind of molecules, but, as the action goes on, these will increase
in quantity. There will be two consequences of this. In the first
place the parent materials will diminish in amount, the collisions
between their molecules will become fewer, and the speed of the for-
ward action will therefore become less and less. In the second place
the increase in the number of molecules of the products will result in
more frequent collisions between them, in more frequent occurrence of
the chemical change which they can undergo, and thus in an increase
in the speed of the reverse action. The forward action decreases in
speed progressively ; the reverse action increases in speed. Finally
the two speeds must become equal, and at that point perceptible
change in the condition of the whole must cease.
The most immediate inference from this mode of viewing the
matter is, that the apparent halt in the progress of the action does
not indicate any cessation of either chemical change. Both changes
must go on in consequence of the continued encounter of the proper
molecules. But since they proceed with equal speed they produce no
alteration in the mass as a whole. In fact, the final state is one of
equilibrium, and not of repose. Hence, chemical changes which are
reversible lead to that condition of seemingly suspended action which
we speak of as chemical equilibrium. The changes themselves are
called reversible, or, since they represent a state of balance between
opposing tendencies, balanced actions.
The detailed discussion of the relations of liquid and vapor (pp. 117,
135), of ice and water (p. 115), and of saturated solution and undis-
solved solid (pp. 152, 160), liquid or gas (p. 154), has already famil-
iarized us with this term and its significance. By the use of the
kinetic hypothesis we can, in fact, apply the very same sets of ideas
to the discussion of any kind of reversible phenomena.
When this is done for reversible chemical actions we perceive that
two circumstances, in particular, are likely to have a noteworthy effect
in the promotion or retardation of a given chemical change. These
are the homogeneity, or the reverse, of the mixture, and the molecular
concentration of each component.
CHEMICAL EQUILIBRIUM 249
The Influence of Homogeneous Mixture. — Evidently the inter-
action of substances will be greatly favored if they are capable of
remaining intimately mixed. When this is the case, as in gaseous
or liquid mixtures, every molecule has an equal opportunity freely to
encounter every other molecule. On the other hand, when one sub-
stance is a liquid and the other a gas not especially soluble in the
liquid, or when one is a liquid and the other a solid suspended, but
not dissolved, in the liquid, or when they are solid and gas respec-
tively, the chances for mutual encounters between the two different
kinds of molecules will be very notably restricted. Thus the progress
of an action between inhomogeueously mixed bodies, otherwise capable
of interaction, must be greatly affected by purely physical circum-
stances. Some of the remarkable consequences of this were discussed
fully in connection with the preparation of hydrogen chloride (p. 179).
The Influence of Molecular Concentration. — Even when the
mixture is homogeneous, a second factor will affect the action. The
frequency of the encounters amongst a given set of molecules, result-
ing in a definite chemical change, will evidently depend entirely upon
the degree to which they are concentrated in each other's neighborhood.
Larger amounts of one of the materials, for example, will not result
in more rapid chemical action in the sense which this material favors,
if the larger amount of material is also scattered through a larger
space. Chemical changes of this kind, therefore, are not accelerated
by increasing the mere quantity of any ingredient, but only by increas-
ing the concentration of its molecules. Hence, in practice, we find
that the distribution of the molecules in a system which has reached a
state of equilibrium is at once affected if we alter the molecular con-
centration of any of its components. Thus, if, in the action of hydro-
gen upon iodine, we introduce into the same space, an extra amount
of hydrogen, this facilitates the formation of hydrogen iodide by in-
creasing the possibilities of encounter between hydrogen and iodine,
while at the same time it does not affect (cf. p. 88) the number of
encounters in a given time between hydrogen iodide molecules which
result in the reverse transformation. The proportion of hydrogen
iodide formed, therefore, from a given amount of iodine will be
greater, although the potential maximum has not been altered since
the quantity of one ingredient only has been increased. The intro-
duction of an excess of iodine would have had precisely the same
effect.
250 INORGANIC CHEMISTRY
It is easy to illustrate this experimentally. If ferric chloride
and ammonium thiocyanate are mixed in aqueous solution, a liquid
containing the soluble, blood-red ferric thiocyanate is produced ;
'FeCl3 + 3NH4CNS^Fe(CNS)3 + 3KH4Cl. The action is a rever-
sible one. Now, if the two just-named salts are mixed in very dilute
solution in the proportions required by the equation, say by adding
20 c.c. of a deci-normal solution (p. 149) of each to several liters of
water, a pale-reddish solution is obtained. When this is divided
into four parts, and one is kept for reference, the addition of a little
of a concentrated solution of ferric chloride to one jar, and of ammo-
nium thiocyanate to another, will be found to deepen the color by
producing more of the ferric thiocyanate. On the other hand, mix-
ing a few drops of concentrated ammonium chloride solution with the
fourth portion will be found to remove the color almost entirely on
account of its influence in accelerating the backward change.
We may state (see below) as a general principle that, in reversible
actions between substances in homogeneous mixture, the amount of a
chemical change taking place in a given time -will be dependent upon
the molecular concentration of each ingredient. It will also naturally
depend upon the intrinsic affinity (cf. p. 111). It will likewise vary
with the temperature, since the affinity is affected by change in tem-
perature.
The phrase " active mass " is sometimes employed instead of the words " molec-
ular concentration." It is somewhat misleading, however, for, as we have seen,
it is not on the mass of a substance, but on the quantity of it in a given volume,
that the speed of the action depends.
Law Connecting Molecular Concentration and Speed of Re-
action. — These principles must now be stated in somewhat more pre-
cise terms. We confine our attention to homogeneous mixtures, in
the first place, and to non-reversible actions.
The concentration of the molecules is usually expressed, for each
substance, in terms of the number (whole or fractional) of moles
(gram-molecular weights, p. 197) of the substance contained in a liter
of the whole mixture. There is the same number of molecules in a
mole of every substance, namely, the number of molecules in 32 g. of
oxygen (cf. p. 199). Hence the number of moles per liter defines the
concentration of the substance in terms of this number of molecules
in a liter as the unit of concentration.
Thus, in a solution containing 25.4 g. of free iodine (-fa of a formula
CHEMICAL EQUILIBRIUM 251
weight, I2) per liter, the solution is 0.1 molar (p. 149), and the molec-
ular concentration of the iodine is 0.1. When the substance is a gas,
the concentration of the molecules is proportional to the partial press-
ure of the gas. Now, one mole of a gas occupies 22.4 liters at one
atmosphere pressure (and 0°). Hence, when one mole of a gas is con-
tained in 1 liter, and its molecular concentration is therefore 1, it exer-
cises 22.4 atmospheres partial pressure. When the partial pressure of
one gas in a mixture is two atmospheres, its molecular concentration is
-2- or 0.09.
22-4
In an action of the form A + B — » C + D, let cl and c2 represent
the number of moles per liter of the materials A and B, respectively,
at any stage of the interaction. Then the speed (S) of the forward
action, expressed in moles of A and B transformed per unit of time
(say, per hour), is found to be defined by the relation cx x c2 x F = S.
Of course ^ and c2 (and therefore, also, S) diminish steadily, for the
materials A and B are being progressively used up. S, therefore, at
any moment is the amount which would be transformed if the concen-
trations present at that moment were artificially maintained during the
whole hour, ^expresses the intrinsic activity (affinity) propelling the
action, which is independent of concentration. If unit concentrations
are taken, ct = c2 = 1, and therefore F = S. Thus F, the activity, is
always represented numerically by the speed in moles per unit of time
when the concentration of each ingredient is unity. For rapid actions
a shorter unit of time than the hour may be used.
In practice the value of F for any action can easily be determined,
because cv c2, and S can be measured. As a check on the results,
different concentrations will be used at the same temperature, and the
amounts transformed in measured times will be observed in each
case. We can then calculate from the data of each set the speed
(moles transformed per hour or minute) which would be shown by
constantly maintained unit concentrations of the materials. The
answers are the values of F based upon different concentrations of the
same substances, and they agree closely with one another. The values
of F for several different chemical actions, however, may differ widely,
and are measures of the relative activity of each.
If the action is more complex, the same principles apply. Thus, if an expres-
sion for the speed of the action 2 A + 3B + C — * D + E is required, we must con-
sider this equivalent to A + A + B + B + B+C—+D + E. Here the speed
forwards (S) = c, x ct x c2 x c2 x c2 x c3 x F, or c^c^F = S.
It need hardly be added that, clearly, when the concentration of any ingredient
252 INORGANIC CHEMISTRY
becomes zero (say by all of it entering into combination), or when the affinity is
zero, the speed S must become zero, that is to say, no action takes place.
The relation of the theoretical speed with constant concentration,
used in the above formulae, to that which is observed with diminishing
concentration is rather complex. For an action of the form A — » B + C,
where the change in only one molecule constitutes the action, if ct is
the initial molecular concentration of A, and x is the fraction of this
which is transformed in the time t,
When two molecules have to interact, the formula is still more com-
plex. If the substances are present in equivalent proportions, their
molecular concentrations in this special case are alike, and may each
be represented by cr The relation is then :
C C ~
• f
-r- t.
The mathematical derivation of these relations will be found in any
work on physical chemistry.
The law of molecular concentration is, therefore : In every chemi-
cal experiment, the speed of the action at any moment is proportional
to the first, or some higher power of the molar concentration, for the
time being, of each interacting substance, and to the affinity at work.
The following illustration (see also Sulphurous acid) will make all
this clearer. When arsine AsHs (q.v.} is heated at 310°, it decomposes
gradually into hydrogen and arsenic :
AsH -> As
The action is not appreciably reversible. The arsenic assumes the
solid form. The gas is inclosed in a tube which is kept in a bath at
310°, and a manometer shows changes in pressure. Since, as the action
proceeds, 1£ molecules of hydrogen take the place of each molecule of
arsine, the total pressure slowly increases. Every increase of 1 mm. in
the pressure is the result, therefore, of an addition of 3 mm. partial
pressure of hydrogen and a reduction of 2 mm. in the partial pressure
* Ordinarily we should write the equation 2AsH3 — > 2As + 3H2. But this
form would make the speed proportional to c^, and calculation then gives us in-
constant values for S. The reason for this irregularity is given in works on
physical chemistry.
CHEMICAL EQUILIBRIUM
253
of the arsine. The molecular concentrations are proportional to the
pressures, and change, therefore, in the same ratios. The observations
(first two columns), together with the data deduced from the first two
by calculation, were as follows :
MOLECULAR CONCENTRATIONS.
Tl H
Total.
AsH3 Transfmd.
0
784.84
0.02159 (c,)
3
878.50
0.02416
6.00514
0.0906
4
904.05
• . .
.
8
987.19
We must first ascertain the molecular concentration of the arsine
corresponding to the observed pressure at the beginning. We remember
that at 22.4 atmospheres, or 22.4 x 760 mm. and 0°, the concentration
of a gas has the value 1 (p. 251). The actual initial pressure 784.84
784 84 x 273
mm. at 310° would become, at 0°, ' - , or 367.5 mm. The
-
(oil) -j- £i oj
K
'
or 0.02159 moles
molecular concentration is here, therefore, -^^- —
LIU A X • t)U
per liter. After 3 hours, some hydrogen has been formed. The
pressure has increased to 878.50 mm. Reducing as before, this repre-
sents a molecular concentration of all ingredients of 0.0241 6 moles per
liter. The increase is 0.00257. This, as was demonstrated above, corre-
sponds to a loss of 2 x 0.00257, or 0.00514 moles per liter of arsine.
Now, employing the formula given above, we find the speed per hour :
c 0.02159
-3 =
This result means that, if the concentration of the arsine were to be
maintained at the initial value by continual renewal of the waste,
then 0.0906 (9.06 per cent) of the initial amount would be decomposed
in an hour. Using the pressures at 4 and at 8 hours, the reader will
obtain by calculation, practically the same value for S. Other experi-
ments with still different concentrations, provided the temperature
was the same (p. 250), would likewise give the same result. Hence,
when we take unit concentration (1 mole per liter), c = 1, and the ex-
pression c^F = S becomes F = 0.0906. Thus the affinity or activity
254 INORGANIC CHEMISTRY
of the action may be measured with any concentration, and expressed
in moles transformed per hour with unit concentration. Similar
measurements with other actions then enable comparisons of their
relative activities to be made (see Exercise 9, end of this chapter).
The Condition for Chemical Equilibrium. — Let us now take a
reversible action,
A*±B+ C,
ci C2 ca>
in which cv c2, cs are the molecular concentrations after the condition
of equilibrium has been established. The above mathematical form of
statement for the influence of concentration and affinity on speed of
transformation, being true for all concentrations, will be true for this
set. The speed of the forward and reverse actions will be,
St = c^ and S2 = c2csF2,
respectively, where Fv is the intrinsic tendency of A to decompose,
and Fz the tendency of B and C to combine. As the concentrations
have so adjusted themselves that an equal amount of material is being
transformed each way, we have
C.F. = c2csF: or 5 = ^ •
11 2 » 2 Tf n
** J «|
77*
-=£-, being the ratio of two constants, is constant ( = &). This is the
-*2
ratio of the affinities driving the opposed actions, and is known as the
affinity constant of the reversible chemical change. If, for example,
its value is £, then the speeds of the two actions, if each were to pro-
ceed unimpeded (say in separate vessels) with constantly maintained
unit concentrations of the materials, would be in the ratio F1 : F2
or 1 : 4. From this it will be seen that measurement of the concentra-
tions present in a system which has reached equilibrium gives us the
data for calculating the value of this ratio. In other words, it gives us
the means of ascertaining the relative magnitudes of the intrinsic
affinities of the opposed actions.
We may apply this to the data given (p. 237) for hydrogen iodide. The equa-
tion is
2HI<=±H2 + I2,
"l "2 ^3'
and c,, c2, and c3 are the molecular concentrations after the system has come to rest.
The speed of the forward action (S,) is cl2F1 and that of the reverse action (S2) is
CjC3F,. As concentrations are now such that the speeds are equal, we have
CHEMICAL EQUILIBRIUM 255
T? re
V — 2 3 —
Now, at 448°, with equivalent quantities of the two elements, nearly 0.8 (more
exactly, 79 per cent) of the weight of each is in the form HI, and 0.2 in the mix-
ture H2 + Ij. Thus in every 100 molecules, 80 are HI(c,), 10 are H2(c2), and 10
are I2(c3). Thus k = O.I2 -=- 0.82 = fa. That is to say, the union of hydrogen and
iodine would take place with 64 times as great a speed as the dissociation of hy-
drogen iodide if each action could proceed without reversal and under identical
conditions. Or, in terms of the kinetic theory, the collisions of the H2 and I2
molecules result many times more often in chemical change than do collisions
of HI molecules.
The case of hydrogen iodide is comparatively simple because the volume is not
altered by the progress of the action (see below). The expansion when phos-
phorus pentachloride (p. 208) dissociates compels us to take account of the
volume. The equation is :
PC16^± PC13 + C12, aud^L = k = C-^-
c, c2 ca F2 ct
If one gram molecule of the substance is taken, and x is the proportion dissociated,
/•g 1 _ /j«
and v the volume occupied by the whole, then c2 = c3 = - and ct = -- Thus
x2
k = — - — • Now at 250° (and 760 mm.), for example, 0.8 of the whole weight
{L — XI *0
of material is dissociated : x = 0.8, 1 — x = 0.2. Hence k — 0.82 -r- 0.20 = 3.2 -H v.
To obtain the value of v we note that a gram molecule at 760 mm. and 0° occupies
22.4 liters. At 260° it occupies 22.4 x (250 + 273) -=- 273 1. But this mass of gas
contains 0.8 more molecules because of dissociation, and its volume is, therefore,
1.8 x 22.4 (250 + 273) H- 273 = v = 77.2 1. Thus k = 3.2-=- 77.2 = ± • Otherwise
310
stated, 25 Ft = F2. That is to say, the union of the trichloride and chlorine would
proceed twenty-five times as fast as the dissociation, if each of the three sub-
stances was present in unit concentration, and each action could proceed inde-
pendently without reversal.
The Effect of Changes of Volume on Chemical Equilibrium.
— Our applications of the theory of equilibrium will be chiefly to dis-
solved bodies, and hence the effect on the equilibrium point of changes
in volume (by dilution or the reverse) will require frequent considera-
tion. Now dilution, for example, diminishes opportunities for en-
counters between the substances on both sides of the equation. In the
first of the above illustrations, increasing the volume decreases the
rate at which the chlorine and the trichloride can combine. Since,
however, the speed of the dissociation depends on the state of the
PC15 molecules only, and is unaffected by their nearness to or remote-
ness from one another, the forward action will not be weakened at all.
256 INORGANIC CHEMISTRY
Hence, dilution increases the degree of dissociation. In general,
change in volume will affect the equilibrium point whenever there are
more molecules on one side of the equation than on the other.
In mathematical terms, when we change the volume to - times its
former value (n whole or fractional), the concentration changes n
times. The equation for equilibrium then becomes, momentarily,
nc2 x nca -H n^ ^ k, or nc2ca -i- c1^ k. To restore the value of the
expression to equality with k, change must occur in the concentrations
C2, cs, and cx. When n is < 1, that is, when the volume increases,
some PC16 must pass into the form PC13 and C12 until c2'cs' H- c/= k,
as before. That is, dilution increases the degree of dissociation.
In the case of hydrogen iodide, and in all others where the nurnter of molecules
taking part in the direct and reverse actions is the same, change in the volume of
the system has no effect on the position of the equilibrium point. Thus dilution
diminishes the chance of encounter between two HI molecules to the same extent
that it interferes with encounters between H2 and Lj molecules. Conversely,
increase in all concentrations, by diminution of volume, favors both actions equally.
Hence, at 448°, 79 per cent of HI will always be present at last, whatever the
volume occupied by a given amount of the materials. In mathematical terms, if we
diminish the volume n times (n whole or fractional), we increase the concentration
of each constituent n times. The values become ncu nc2, and nc3 respectively,
, nzc,c3 c2c,
and k = -5-3 = _L«
Heterogeneous Equilibrium. — A modification of the above
conceptions is necessary when the mixture is not homogeneous. If,
for example, one of the constituents is present as a solid or a gas, in
greater amount than can be dissolved by the liquid in which alone the
chemical change takes place, then, according to the definition of
saturated solution (p. 158), the concentration of the dissolved material
will be constant at a given temperature as long as physical equilibrium
between the solid and the solution is maintained. This is a case
especially likely to occur when slightly soluble (so-called " insoluble")
bodies (cf. p. 146) are concerned.
The same reasoning applies also to very slightly volatile solids.
The concentration of the vapor of a solid body present in excess (meas-
ured by its vapor pressure) will be constant so long as the temperature
is fixed, and interaction with a superincumbent gas will take place
chiefly through the vapor.
In both these cases the concentrations of the active parts of the
CHEMICAL EQUILIBRIUM 257
slightly soluble and slightly volatile bodies, respectively, are not sub-
ject to variation — they are constant. Thus, with the action,
2Ba02 «=± 2BaO + 02,
Cl C2 CS>
the concentrations of the vapor of Ba02(c1) and of BaO(c2) are constant,
and that of oxygen (c8) alone is subject to alteration. We have, there-
fore,
^17
or c. = — A;,
in which, since cl} c2, and k are constant, cs must be constant also. But
the pressure of a gas is proportional to its molecular concentration,
according to Avogadro's hypothesis. Therefore, in this action, the
pressure of the oxygen (the dissociation pressure) should be constant
irrespective of the extent to which the dissociation has progressed.
Observation shows that this is the case. When barium dioxide is
maintained at a definite temperature, the pressure of the oxygen rises
to a fixed value and remains constant. Any pressure of oxygen above
the fixed value forces the gas into combination until no barium mon-
oxide is left ; any pressure below this value permits the dioxide to
decompose until none remains. This explains the Brin method of
making oxygen (p. 63), in which a temperature is chosen at which
the dissociation pressure is approximately equal to that of the partial
pressure of oxygen in the air (i atmosphere). All hydrates, as we
have seen (pp. 121-123), behave in a precisely similar way, and fur-
nish numberless confirmations of this application of the law of molecu-
lar concentration.
Applications in Chemistry : Displacement of Equilibria by
Changes Affecting Concentration. — Of especial interest to the
chemist are the conditions under which the equilibrium point may be
displaced and more nearly complete realization of one of the two
opposed changes may be brought about.
We have just seen (p. 249) that one way in which a reversible action
may be forced nearer to completion in one direction or the other is the
introduction of an excess of one of the ingredients contributing to the
action. This method of displacing the equilibrium point, however,
cannot be very effective unless it is possible to introduce an exceedingly
large excess of the selected ingredient in a high degree of molecular
concentration, since this operation does not in any way affect or, in
258 INORGANIC CHEMISTRY
particular, restrain the reverse action which is continually undoing the
work of the forward one. A much more effective means of furthering
such actions is found in removing one of the components of the system.
Any agency which could remove the free iodine as fast as it was
formed in the decomposition of hydrogen iodide, for example, would
entirely stop the reproduction of the compound and so would enable
the dissociation (2HI <=± H2 + I2) to run to completion.
This might be realized * by causing one end of a sealed tube charged
with hydrogen and iodine, after the contents had settled down to a
condition of equilibrium, to project from the bath in which the whole
had been kept at 448° (Fig. 72, which is simply diagrammatic). By
cooling this end, a large part of the
21 per cent of free iodine would
quickly be condensed in it to the solid
form, while the hydrogen would re-
main gaseous. Only the trace of
vapor which cold iodine gives would
Flo- 72- then be available to interact with the
hydrogen and reproduce hydrogen
iodide. Meanwhile the decomposition of the latter would go on, and
thus, eventually, almost all the iodine would be found free in one end
of the tube, and the hydrogen, all free likewise, would occupy the rest.
By this purely mechanical adjustment the chemical change would in
this way be carried from 21 per cent completion to almost absolute
completion.
If, on the other hand, arrangements were made to have powdered
marble, in a sealed bulb of thin glass, inclosed in the tube, we might
imagine the very opposite effect of the above to be produced. The
breaking of the bulb of marble, when equilibrium had been reached,
would provide means for the removal of all the hydrogen iodide ,f while
the hydrogen and iodine would still be gaseous. Thus, the com-
pound having been removed, there would be no reverse action to
compensate for the union of the elements. The whole material would,
therefore, soon have passed through the form HI. Hence, by another
* For another illustration, see under Ammonia.
f The hydrogen iodide would be destroyed by interaction with the marble:
2HI + CaCO3 — > CaI2 + C02 + H20.
The calcium iodide is a solid. The two gases, carbon dioxide and water vapor,
do not interact with hydrogen or with iodine, and would not, therefore, inter-
fere with the formation of fresh hydrogen iodide.
CHEMICAL EQUILIBRIUM 259
mechanical arrangement, an action -which ordinarily could progress to
only 79 per cent would be turned into a complete one.
In e very-day chemical work, since our object is usually to prepare
some one substance, chemists either avoid chemical changes which are
notably reversible, or adjust the conditions so that the reverse of the
action which they desire is prevented. In consequence of this, when
carrying out the directions for making familiar preparations, the fact
that such actions are reversible at all very readily escapes our notice.
Arranging the conditions so that the separation of a solid body by pre-
cipitation, or the liberation of a gas, takes place, are the two com-
monest ways of rendering a reversible action complete. Excellent
examples of both of these are furnished by the chemical change used
in producing hydrogen chloride by the interaction of salt and sul-
phuric acid (p. 179), the discussion of which should once more be
studied attentively.
The escape of one member of a system engaged in chemical interaction,
because it is gaseous or solid, and in either case immiscible with the rest of the
members of the system, is the commonest cause of the obstruction of one direc-
tion of a reversible action and the triumph of the other. This, as we have seen,
is the combined result of the natural behavior of a system in chemical equilibrium,
and of the physical properties, particularly the solubility, of the members of the
system. Two rules, attributed to Berthollet, have been made, however, to describe
these special cases of a broader principle. Unfortunately, it is difficult so to word
them that they shall be entirely unambiguous and entirely correct.
The "rule of precipitation," for example, might read: When certain classes of
materials are brought together in solution, if an exchange of radicals would produce
an insoluble body, this exchange will occur. But then the fact is that, in such
cases, the exchange always occurs to some extent whether any product is insoluble
or not. The insolubility is responsible only for the greater completeness of the
exchange. Crude statements to the effect that "when an insoluble body can
be formed, it will be formed," when close scrutiny shows them to possess any
definite meaning whatever, are grossly misleading. They suggest that insolubility
is a sort of especially desirable career on which the elements are ambitious of
entering.
All forms of these so-called laws are objectionable, because they necessarily
suggest that the positive direction of the action is assisted by the immiscibility of
the product, and this is the precise converse of the fact. The immiscibility does
nothing at all towards assisting the formation of the insoluble substance itself, but
does whatever it can towards preventing the destruction of that substance, once it
is formed, by hampering the negative action.
Affinity vs. Solubility. — The question of the relation of affinity
to the apparently much greater efficiency of one of the directions of
some reversible actions, may now be put in a much clearer light (pp.
260 INORGANIC CHEMISTRY
181, 110, and this Chap.). The whole of the possibilities of progress
for any action are expressed by a function (p. 251) of the form c^c2F= S.
If any one of the variables, say one of the concentrations (cj, is neg-
ligible, the product must be small, irrespective of the values of the
other factors. Thus the feebleness of a chemical action only shows
that the product of all the variables is minute, and not that the
affinity factor per se is of small magnitude.
Displacement of Equilibria by Changes in Temperature :
Van't Hoff's Law and Le Chatelier's Law. — Since the affinities
driving the two opposed actions are differently affected by change in
temperature, the value of their ratio alters, and the equilibrium point,
which depends on this ratio, suffers displacement. It is found that the
direction of the displacement depends on which of the two actions
absorbs heat, and which gives it out. In a system which is in equilib-
rium, of the two opposed interactions, that one which is endothermal
is promoted, while that which is exothermal is resisted, by raising the
temperature, and vice versa. This is van't Hoff's law of mobile equilib-
rium. Thus the dissociation of phosphorus pentachloride (p. 208) is
endothermal (p. 27), and hence becomes greater at higher tempera-
tures. All thermal dissociations are of this kind. Again, the inter-
action of steam and iron (p. 110) is exothermal, and so the higher the
temperature, the more conspicuous the opposite action becomes. In
view of this generalization, we can understand the formation of endo-
thermal substances of an unstable character at high temperatures.
Thus, ozone (q.v.) is produced by electrical discharges, and cyanogen
(g.v.) is formed in the blast-furnace.
The principle applies also to physical equilibria. Thus, as the
temperature rises, a compound which gives out heat in dissolving is
less soluble in a solution already almost saturated with the compound ;
while one which absorbs heat in dissolving is more soluble in such a
solution. For example, anhydrous sodium sulphate gives out heat
in dissolving (p. 165), and its solubility diminishes (p. 158) with ris-
ing temperature, while with hydrated sodium sulphate just the con-
verse is observed. So, also, the vaporization of a liquid absorbs heat,
and hence the concentration of vapor it can maintain increases with
rise in temperature.
The above law is a particular case of a still more general one frequently
called Le Chatelier's law: If some stress (for example, by change of tem-
perature, pressure, or concentration) is brought to bear on a system in equi-
CHEMICAL EQUILIBRIUM 261
librium, by which the equilibrium is displaced, the equilibrium is dis-
placed in that direction which tends to undo the effect of the stress.
Thus, pressure causes ice to melt because the water which is formed occupies a
smaller volume, and the change therefore tends to relieve the pressure. So, also,
the production (say, by a chemical interaction) of a large amount of a body in
solution, by which its concentration is increased beyond saturation, leads to crys-
tallization (or precipitation) of the excess.
Exercises. — 1. What is the molecular concentration of the oxygen
in the air (pp. 155, 250), of the nitrogen in the air, of the aqueous vapor
above water at 10° and at 20° (p. 116), of a solution containing one
formula- weight of sodium chloride in 10 liters, of a solution containing
65 g. of hydrogen iodide in 250 c.c. ?
2. What are the partial pressures of the three components of
phosphorus pentachloride vapor at 250° and 760 mm. (p. 207) ? What
are their molecular concentrations ?
3. Using the model on p. 254, study the dissociation of KI8
(p. 235), of iodine vapor (p. 236), and of hydrogen iodide (p. 237),
and the formation of ferric thiocyanate (p. 250). Show in each
case the effect on the system of increase in volume without change in
the amount of material (p. 255).
4. Using the model on p. 257, study the dissociation of mercuric
oxide (p. 12), assuming the compound to be involatile, and the inter-
action of iron and steam (p. 110). Why can magnetic oxide of iron
be -reduced completely by a stream of hydrogen (p. 110), and iron
oxidized completely by a current of steam (p. 99) ?
5. What actions in Chap, xiv are complete for the same reason
that the action of sulphuric acid on salt (pp. 178-180) is so ?
6. Why is the formation of the following substances complete :
silver chloride (p. 186), and hydrogen chloride and water by union
of the elements ?
7. What inference should you draw from the fact that the solu-
bilities of potassium nitrate, sodium chloride, and Glauber's salt
(p. 157) increase with rise in temperature (p. 260), and from the fact
that those of calcium citrate (p. 156) and triethylamine decrease with
rise in temperature ?
8. Is the heat of solution of lead nitrate (p. 157) positive or nega-
tive ?
9. Carry out the calculation of S for 4 and 8 hours (p. 253).
262 INOEGANIC CHEMISTRY
SUMMARY OF PRINCIPLES.
The summary of some of the chief principles of the science (p. 188) may now
receive several important additions. For the sake of completeness, reference to
the periodic system is made in No. 21, to isomers in No. 22, and to the phase rule
in No. 23, although these subjects have not yet been taken up. As before, all
hypothetical matters have been excluded as far as possible.
15. That weight of each substance which in the gaseous condition occupies the
same volume as 32 grams of oxygen, temperature and pressure being alike for
both (namely, 22.4 liters at 0° and 760 mm.), is taken as the chemical unit of
weight for the substance, and is known as its molar weight (p. 199).
16. That weight of each element which is the greatest common measure of the
quantities of the element found in the molar weights of its compounds is taken
as the chemical unit of weight for the element, and is known as its atomic weight.
This weight has the property described in 6 (p. 188).
The composition of each substance is expressed in terms of the atomic weights
as units, and the sum of the atomic weights is multiplied by an integer, when
necessary, so as to equal the molar weight (p. 204).
17. The number of equivalent weights of hydrogen which combine with, or are
displaced by the atomic weight of an element is called the valence of the element
(p. 103).
18. The speed of every interaction is a function of the first, or some higher
power of the molar concentration of each interacting substance (p. 252).
19. Substances undergoing, at a fixed temperature, an interaction which is
reversible, reach a condition of equilibrium. The final proportions of the mate-
rials are such that the speeds (see 18) of the opposed actions are equal (p. 254).
20. Van't Hoff's law and Le Chatelier's law (p. 260).
21. Each element has its own set of chemical relations (pp. 177, 226) : e.g., it
can exist in combination with certain other elements ; it has a certain valence,
and may have more than one valence ; it confers certain properties on its com-
pounds as a class ; it resembles certain other elements in several of these respects
(e.g., the halogens), and differs from others, in a way more or less definitely
described by its place in the periodic system (q.v.).
In complex cases, the inter-relations of the elementary units in a compound,
and the relations of the compound to other compounds (see No. 22), are represented
graphically by formulae based upon an hypothesis of molecular structure (p. 224).
22. Identical combinations of matter may constitute more than one compound
substance (isomers, see Urea). These may have equal molar weights (optical
and structural isomers), or they may have different molar weights (pp. 204,
242).
23. In a system in equilibrium the number of components plus two equals the
number of phases plus the number of degrees of freedom (Phase rule, q.v.).
CHAPTER XVI
OXIDES AND OXYGEN ACIDS OF THE HALOGENS
THE chief subjects of practical importance touched upon in this
chapter are connected with bleaching powder (CaCl(OCl)), and potas-
sium chlorate (KC103) and perchlorate (KC104). Hence our attention
will be largely directed to the modes of making these substances and to
their relations to one another. Incidentally, we shall encounter many
actions of a complex and, to us, more or less novel kind.
Compounds of Chlorine Containing Oxygen. — The following
are the names and formulae of the substances :
HC10 Hypochlorous acid, C12O Hypochlorous anhydride,
(HC1O2) Chlorous acid,
C1O2 Chlorine dioxide,
HC103 Chloric acid,
HC1O4 Perchloric acid, C12O7 Perchloric anhydride.
There are also compounds of metals with the negative radicals of
these acids. Of this nature are the three substances mentioned in the
first paragraph. Chlorous acid is itself unknown, but potassium chlorite
(KC102) and some other derivatives have been made.
The two anhydrides (p. 71 ), when brought into contact with water,
combine with it to form the acids opposite which they stand in the
table. Chlorine dioxide (<?.?>.), however, is not related to any one acid
in this way.
All these compounds differ from most that we have hitherto dis-
cussed inasmuch as not one of them can be made by direct union of
the simple substances.
Nomenclature of Acids and Salts- — When several compounds
closely related in composition, like the above acids, are known, a sys-
tematic method of naming them is used. The terminations -ous and
•4c indicate smaller and larger proportions of oxygen respectively.
For compounds below or above those two in their degree of oxidation,
the prefixes hypo- and per- are employed,
263
264 INORGANIC CHEMISTRY
When the radicals (p. 93) contained in the acids are combined
with metals, the compounds are spoken of as salts of the respective
acids. Thus, KC103 is described as the potassium salt of chloric acid.
The specific names for these salts are distinguished by terminations
corresponding to those of the acids :
KC1O Potassium hypochlorite, KC103 Potassium chlorate,
KC102 Potassium chlorite, KC104 Potassium perchlorate.
The termination -ite corresponds to -ous, -ate to -4c. This principle is
applied systematically, so that the salts of sulphuric and sulphurous
acids, for example, are called sulphates and sulphites respectively.
Compounds containing no oxygen receive the termination -ide.
Thus, KC1 is potassium chloride, FeS is ferrous sulphide.
Salts and Double Decomposition. — We have just been using
the word salt in a general sense. It is the class name for a set of sub-
stances which includes common salt or sodium chloride (NaCl), potas-
sium nitrate (KN"08), sodium sulphate (NagSC^), silver chloride (AgCl),
potassium chlorate (KC103), etc. The majority of the substances used
in elementary chemistry belong to this class. They receive the name
because in certain important chemical respects they behave like com-
mon salt. For example, when sodium chloride is treated with acids,
such as sulphuric acid (p. 178) or phosphoric acid (p. 179), hydrogen
chloride is liberated. Actions of this kind consist in an exchange of
radicals and are all reversible. An action of the same type, although
outwardly different, is that of sodium chloride and silver nitrate in
aqueous solution (p. 99). Here we have two salts interacting instead
of an acid and a salt, and we get a precipitate instead of a gas. But
the interchange of radicals is exactly similar.
Now salts in general behave in these respects in the same way as
does common salt. They interact with acids or other salts, particu-
larly in solution, in such a way that an exchange of radicals takes
place. In the first case, a salt and an acid, and in the second case
two salts, are produced. These actions are all reversible. Acids differ
thus from salts only in the fact that one of their radicals is hydrogen.
Hence they are frequently called hydrogen chloride, hydrogen sulphate
(H2S04), and so forth. It may be added that bases (p. 119), like potas-
sium hydroxide (KOH), interact reversibly with salts and acids,
exchanging radicals after the same fashion.
All salts are named according to the radicals which they contain.
OXIDES AND OXYGEN ACIDS OF THE HALOGENS 265
Thus, all containing — S04 are sulphates. Conversely, when the name
of a salt is given, the formula can be written down at once. In doing
this, however, regard must be had to the valence of the radicals
(p. 104).
In view of the reversibility of most of the interactions of salts,
acids, and bases, we encounter completed changes chiefly when pre-
cipitation occurs, or when one product is volatile (p. 259). If neither
of the products formed by the exchange of radicals is insoluble, the
reversibility of the action prevents our obtaining anything but a
mixture. Only those double decompositions which involve more or less
insoluble or volatile substances are thus of use for preparing salts. The
action of sodium chloride on silver nitrate (p. 99) is an example. The sil-
ver chloride is almost completely insoluble, while the sodium nitrate pro-
duced by the change remains dissolved. By nitration we obtain the
silver chloride as a powder, while the evaporation of the filtrate gives
us the soluble product. This sort of action can be used, therefore,
either for the preparation of a soluble or an insoluble substance. If
the problem is to make a soluble product, then we must arrange an
action between two substances, each containing one of the two required
radicals, and possessing two other radicals, which, when united, give
an insoluble body. This plan is illustrated frequently in what follows.
Hypochlorites. — Since none of the acids in our list can be made
directly from their elements, we generally have to prepare, first, the
corresponding salt. From the salt, by double decomposition, the acid
is then secured. Hence, in each case, the salts will be discussed
first.
Chlorine interacts slightly with water (p. 176), producing small
quantities of hydrogen chloride and hypochlorous acid (equation (1),
below). The action is very strongly reversible. That is to say, since
the last two substances interact very vigorously to reproduce chlorine
and water, the direct action does not make much progress.
When, however, some substance which can interact with these
products is added to the solution of chlorine, or when chlorine gas is
simply passed into an aqueous solution of such a substance, displace-
ment of the equilibrium point at once occurs (p. 258). Now potas-
sium hydroxide is a suitable substance. It interacts almost completely
in solution with both the products of this action, producing potassium
chloride (2) and potassium hypochlorite (3), according to the last two
of the following equations :
266 INORGANIC CHEMISTRY
Cls + H2O ±5 HC1 + HOC1, (1)
HC1 + KOH ->KC1 + H20, (2)
HOC1 + KOH -» KOC1 + H.,0.* (3)
Thus, omitting the water which appears both among products and
initial substances, and the two acids which are used up as quickly as
they are produced by the first action, we get, by addition of the three
partial equations (cf. p. 228), the final equation :
C12 + 2KOH -> KC1 + KOC1 + H20.
This sort of action does not give a pure hypochlorite, but for some
purposes the presence of the chloride in the solution is not objection-
able.
Bleaching powder, CaCl(OCl), is manufactured on a large scale by
an action exactly like the above. The neutralization of a molecule of
each of the two acids, however, can be accomplished by a single molecule
of slaked lime (calcium hydroxide, Ca(OH)2), since the latter contains
two hydroxyl (OH) groups. The hydroxide can be applied either in
the dry form or mixed with some water as a paste. The separate
actions and final equation are as follows :
Cl, +H20 i=;HCl + HOC1 (1)
/OH + HC1 /Cl
'a ' <a\OCl
CL, +Ca (OH)2 -> CaCl(OCl) + H20
Bleaching powder (see Calcium) is a salt of calcium involving two
different acids (a mixed salt). This condition, again, does not inter-
fere with the application of the substance commercially. A method
of obtaining pure hypochlorites, however, will be found below.
Hypochlorites change into chlorates (q.v.) when heated. They
may also give off oxygen, 2CaCl(OCl) — » 2CaCLj + 02. Although
this decomposition is slow in cold solutions of hypochlorites, or when
they are preserved in the dry form, it may be hastened by means of
catalytic agents. The addition of a little cobalt hydroxide (q.v.)
to bleaching powder solution causes rapid evolution of oxygen. The
hypochlorites are manufactured because hypochlorous acid, which is
used in bleaching, can readily be made from them. The acid itself
will not keep, except when largely diluted, and consequently cannot
be transported conveniently.
* The interaction of potassium hydroxide, or any other base, with any acid to
produce a salt and water, is called neutralization (cf. p. 186),
OXIDES AND OXYGEN ACIDS OF THE HALOGENS 267
Preparation of Hypochlorous Acid: Hypochlorous Anhy-
dride. — 1. The common method of obtaining the acid is by double
decomposition, using some other acid (p. 264). Even from such a
mixture, or mixed salt, as is produced by the action of chlorine on
a base (p. 266), the acid may be obtained in fairly pure condition.
Thus, with nitric acid (an active acid), we have, simultaneously, the
two reversible actions :
( KOC1 + HN08 t=» KN03 + HOC1,
\ KCI + msro3 <± KN03 + HCI.
But hypochlorous acid is a feeble acid, while hydrochloric acid is an
active one, so that in the former action the reversing tendency is very
slight, while in the latter it is vigorous. Hence, by adding nitric acid,
in amount barely sufficient for the liberation of the hypochlorous acid
alone, and doing this in a very dilute solution, the object is attained.
The potassium chloride is hardly affected. By gently warming the
liquid a dilute solution of hypochlorous acid can be distilled off. The
operation may be performed even more successfully by use of a weak
acid, instead of an active one like nitric acid. Boric acid (q-v.') is
suited to the purpose.
2. Advantage may be taken of the feebleness of hypochlorous acid
in another way. When powdered chalk (CaC08) is added to chlorine
water, or chlorine is passed into water holding chalk in suspension, only
the hydrogen chloride has any appreciable action upon the chalk. It
gives, by exchange of radicals, calcium chloride and carbonic acid
(equation (2), below). The latter decomposes immediately with effer-
vescence, yielding water and carbon dioxide gas (equation (3)).
In this particular example of equation-making (cf. p. 229) the
three partial equations, in their simplest forms, may not be added
together. Before this can be done we must always see that the terms
which appear on both 'sides of the equations, awe? are not actual
products, will disappear by cancellation. Here, the 2HC1 in equation
(2) will not cancel with the first product in equation (1), unless the
whole of this equation is multiplied by two. The multiplication has
therefore been effected :
2C11 + 2H20 <^2HC1+2HOC1 (1)
(2)
(3)
2C12 + CaC03 + H20 -> CaCL, + C02 f + 2HOC1.
2HCH
-CaCO^CaCL,-
H2C08 fcj H20 -+
f H2C08
• C02
268 INORGANIC CHEMISTRY
This action gives a solution containing calcium chloride and hypo-
chlorous acid, and, by distillation, as in the first method, a dilute solu-
tion of the acid may be secured.
3. The anhydride of hypochlorous acid (CijO) may be obtained by
passing chlorine gas over mercuric oxide. For this purpose precipi-
tated mercuric oxide (q.v.) must be used, and it should be heated in
advance in order to coarsen the grain of its particles and so diminish
the speed with which it is acted upon. Each of the constituents of the
oxide combines with chlorine :
The mercuric chloride then unites with another formula-weight of the
mercuric oxide to form a compound HgO,HgCl2, which remains in
the tube. The chlorine monoxide is a reddish-yellow gas. It may
be reduced to the liquid form, and boils at + 5°. Both the gaseous
and liquefied forms of it, the former when heated, the latter when
touched by paper or dust, decompose into the constituents with explo-
sion. The gas dissolves in water very easily (200: 1, by vol.). The
yellow solution of hypochlorous acid which results,
has a strong odor of chlorine monoxide. The combination is reversible,
and, when the liquid is warm, a little of the gas escapes.
Properties of Hypochlorous Acid. — 1. Hypochlorous acid
cannot be made, excepting in solution, or kept, excepting in dilute
solution. This is in consequence of its tendency to decompose in three
different ways, one of which has just been mentioned (see 3 and 4
below).
2. As an acid it neutralizes (p. 266) active bases, giving hypo-
chlorites.
3. If the solution is concentrated, much of the hypochlorous acid
changes gradually into chloric acid and hydrogen chloride. This
occurs even in the dark.
3HC10->HC103+2HC1.
4. When the solution is warmed, but more especially when it is
exposed to sunlight, oxygen is evolved rapidly.
OXIDES AND OXYGEN ACIDS OF THE HALOGENS 269
*
This decomposition always takes place whether the acid is present
alone in the water, or along with other substances. Hence, the solu-
tion of chlorine in water, which contains a small amount of hypo-
chlorous acid, on being exposed to the bright sunlight gives off bubbles
of oxygen in rapid succession. This decomposition, since it re-
moves one of the interacting substances in the reverse action, C12 -{-
H20 <— ; HC1 + HOC1, enables the interaction of chlorine and water to
go on to completion. Consequently, the final liquid contains nothing
but hydrochloric acid and water. Leaving out the intermediate steps
again, the action appears, therefore, to be simply a decomposition of
water by chlorine.
2CL, + 2H20^ 4HC1 + 02.
5. In consequence of the ease with which it gives up oxygen, hypo-
chlorous acid is a strong oxidizing agent.
Hypochlorous Acid as an Oxidizing Agent : BleacJiing. — Both
iodine and bromine are oxidized by hypochlorous acid, the former
much more rapidly than the latter, 2HOC1 + I2 -+ 2HOI + CL,.
Further oxidation to HI03 occurs immediately. Although iodine has
less affinity for hydrogen than has chlorine (p. 239), this action shows
that the relation towards oxygen is just the opposite. Here the iodine
goes into combination and the chlorine is displaced.
It is on account of its oxidizing power that hypochlorous acid is
used commercially in bleaching. It is not applied to paints, which
are chiefly mineral substances, but to complex compounds of carbon,
such as constitute the coloring matters of plants and of those artificial
dyes whose manufacture has now become so gigantic an industry. It
should be understood that the great majority of the complex com-
pounds of carbon are colorless. Even a slight chemical change, affect-
ing only one or two of the atoms in a complex molecule, is thus almost
sure to give a colorless or much less strongly colored material. Indigo
(C16H10lSr202), which has a deep-blue color, is an example of a vegetable
dye which is also made artificially. Hypochlorous acid oxidizes it to
isatin, a yellow substance relatively pale in color :
C16H10N202 + 2HOC1 -> 2C8H5N02 + 2HC1.
In ways just as definite as this, hypochlorous acid will change the
composition of other colored substances, although, since we do not
know the formulae of all these substances, we cannot always write
270 INORGANIC CHEMISTRY
«
equations for the actions. Thus, the interaction by which chloro-
phyll, the green coloring matter of plants, is bleached is doubtless
similar to the above, although the formulae of materials concerned are
unknown.
On account of the hypochlorous acid which is already present in
chlorine water, this solution is a very efficient bleaching agent. The
removal of this one of the factors in the reverse action (p. 265) enables
more of the acids to be produced from the chlorine and water until the
whole of the former has been consumed.
As a rule, bleaching is actually carried out by liberating hypochlo-
rous acid from bleaching powder by means of sulphuric acid :
,OCl H. HOC1)
Ca + S04->CaS04+ [t^H^O + Clj.
XC1 H/ HC1 )
Of course, temporarily, most of the hypochlorous acid interacts with
the hydrochloric acid to give chlorine and water, but, as the residual
hypochlorous acid loses its oxygen, the secondary action is again dis-
placed backwards until the chlorine is all used up.
The yarn or cloth is first cleansed from fatty or oily material by
boiling with soap solution. It is then immersed in bleaching powder
solution, and finally in dilute sulphuric acid. Both solutions must be
very weak in order that no interaction may occur with the fabric it-
self. The last two processes may be repeated, if the brownish or
yellowish coloring material has not disappeared after the first treat-
ment.
Hypochlorous acid can be used to bleach linen or cotton, because
the body of these materials, apart from the small amount of coloring
matter, is composed of compounds containing nothing but carbon,
hydrogen, and oxygen. These compounds are very slowly affected by
hypochlorous acid, unless too strong a solution is used, or the exposure
to its influence is too long. That chemical interaction does occur is
shown by the "rotting" of goods which have not been washed
thoroughly after bleaching. Wool, silk, and feathers, on the other
hand, are composed largely of compounds containing nitrogen in addi-
tion to the above three elements. Their constituent material interacts
as easily with hypochlorous acid as do the traces of coloring substances.
Hence, since the fabric itself would be attacked by this agent, different
means of bleaching have to be used for materials of this class.
It should be understood that a cold dilute solution of hypochlorous
OXIDES AND OXYGEN ACIDS OF THE HALOGENS 271
acid may be kept almost indefinitely and will not give up its oxygen
spontaneously. The transfer takes place when, and only when, the
acid comes in contact with some substance capable of uniting with
oxygen.
Thermochemistry of Hypochlorous Acid. — As we have seen
(p. 27), chemical changes which proceed spontaneously to completion
are accompanied by a transformation of chemical energy into some
other form of energy. Hence, a substance, or system of substances,
which undergoes such a change, possesses more chemical energy
and activity before the change than after it. In consequence, if
some given chemical change uses the products of such an action,
and can be brought about by the employment of the original sub-
stance, the employment of the latter will involve a greater liberation
of energy, and will therefore be more likely to secure the consumma-
tion of the change in question.
The decomposition of hypochlorous acid and of chlorine monoxide
are cases where there is a very marked difference between the amount
of chemical energy in the original substances and in the products of
decomposition, hydrogen chloride and free oxygen in the first case, and
free chlorine and oxygen in the second. Hence the changes into these
substances sometimes are of the nature described as explosive. A more
important fact, however, is that, on this account, hypochlorous acid
and chlorine monoxide are more active oxidizing agents than is free
oxygen gas. The energy liberated in the decomposition of the hypo-
chlorous acid has to be added (p. 78) to that which free oxygen
could give, if performing the same oxidation, in order that the total
fall in energy, which measures the tendency of the action to take
place, may be estimated. Hence, substances that are not affected by
free oxygen may be changed instantly by hypochlorous acid. This
explains, for example, the oxidation by hypochlorous acid of many
carbon compounds, including those which are colored, when atmos-
pheric air is without action. Thus, the heat liberated in the oxidation
of indigo to isatin by oxygen gas, if it could be carried out, would be
1800 cal. The much greater heat liberated when hypochlorous acid
is used, we obtain by adding the thermochemical equations :
2HC10 = 2HC1 + 20 + 18,600 cal.
C16H10Na02+ 20 = 2C8H5N02 + 1800 cal.
CWH10N80, + 2HC1O = 2C8HBNO2 + 2HC1 + 20,400 cal.
272 INORGANIC CHEMISTRY
The following thermochemical equations give a rough Idea of the relative
oxidizing powers of the chief oxygen acids of the halogens :
HC1O, Aq = HC1, Aq + 0 + 9300 cal., or + 9,300 1
HC103, Aq = HC1, Aq + 30 + 15,300 cal., or + 5,100
HC1O4, Aq = HC1, Aq + 40 + 700 cal., or + 170
HBrO3,Aq = HBr, Aq + 3O + 15,000 cal., or + 5,000
HIO3, Aq = HI, Aq + 3O - 42,900 cal., or -14,300
cal. for
each atomic
* weight of
oxygen.
HI04, Aq = HI, Aq + 40 - 34,500 cal., or- 8,600 ,
Formerly a different explanation for actions like that of hypochlorous acid,
•when it behaves as an oxidizing agent, was offered. It was suggested that the
oxygen was first liberated from the acid (HOC1— »HC1 + O), and that the single
atoms of the element so produced were more active than molecular oxygen. This
oxygen, which was supposed to interact in the moment of its production, was
called nascent oxygen. But it will be seen that such an explanation is entirely
unnecessary. The activity of the hypochlorous acid on account of its large store of
free energy sufficiently accounts for the facts (see Nascent hydrogen).
Simultaneous, Independent Chemical Changes in the Same
Substance. — As we have seen, hypochlorous acid undergoes three
different changes. Some molecules decompose into water and chlorine
monoxide (p. 268), while others give chloric acid and hydrogen chlo-
ride, and still others hydrogen chloride and oxygen. Since the
same molecule cannot undergo more than one of these different
changes, it follows that the actions are independent of one another.
This is shown by the fact that in sunlight the third predominates,
while in the dark it falls far behind the second. Since the relative
quantities of the products vary, the several simultaneous actions
cannot be put in the same equation. The fundamental property of
an equation is to show the constant proportions by weight between
every pair of substances in it. Hence three separate equations are re-
quired in the present, and in all similar cases where all the proportions
are not constant (cf. p. 231 and see Perchlorates). Successive actions,
like (1) preceded by (2) in the next section (cf. p. 266), however, may
be combined in one equation, since in them all the proportions must
necessarily be constant. These equations are interlocked, for (1)
consumes what (2) produces.
Chlorates. — Like hypochlorous acid itself, the hypochlorites turn
into chlorates. Thus, when chlorine is passed into a warm, concen-
trated solution of potassium hydroxide, and particularly when an
excess of chlorine is used, the hypochlorite changes into chlorate as
fast as it forms :
3KC10 -> KC108 + 2KC1. (1)
OXIDES AND OXYGEN ACIDS OF THE HALOGENS 273
To secure the three molecules of the hypochlorite, the equation form-
erly given (p. 266) must be tripled :
3CL, + 6KOH -» 3KC1 + 3KC10 + 3H20. (2)
When these are added, and the intermediate substance is left out, the
final equation is obtained :
3C1, + 6KOH -> KC103 + 5KC1 + 3H20.
When the solution is cooled, the chlorate crystallizes out.
This action involves converting five-sixths of the valuable potas-
sium hydroxide into the relatively less valuable potassium, chloride.
Hence, in practice, the makers carry out the corresponding action
with calcium hydroxide. They then, add potassium chloride to the
resulting solution, containing calcium chloride and calcium chlorate
(Ca(ClO8)2). The potassium chlorate, formed by double decomposi-
tion, crystallizes when the solution is cooled.
All chlorates are at least moderately soluble in water. Potas-
sium chlorate is used in making fireworks, explosives, and matches.
An intimate mixture with sugar (C^H^Ojj) burns with semi-explosive
violence, the oxygen of the salt combining with the carbon and hydro-
gen to form carbon dioxide and water. Detonating fuses for artillery
are made of a mixture of this salt with antimony trisulphide (q-v.~).
The Separation of Substances by their Solubility. — When
neither of the products of an action approaches absolute insolubility, a separation
may nevertheless be effected more or less perfectly by taking advantage of differ-
ence in solubility. Thus, in the practical method of making potassium chlorate,
the calcium chloride is exceedingly soluble, while the potassium chlorate is only
moderately so. Then, too, the solubility of the latter decreases rapidly as the
temperature is lowered (Fig. 61, p. 167). Hence, it is found that when the
mixture is cooled to — 18° only about 13.5 g. of potassium chlorate remain dis-
solved in each liter, and are lost. At 0° the loss would be greater, for at this
temperature a liter of pure water would hold 33.3 g., and a liter of this solution
would contain more than this on account of the uncompleted reversible action
(c/. p. 264) :
Ca (C1O3)3 + 2KC1 <=± CaCl2 + 2KC1O3.
It will be seen that we reason as if the solubility of each substance was inde-
pendent of the presence of other dissolved bodies (p. 153).
By the use of this principle, and the data in regard to solubility in Fig. 61
(p. 157), a rough idea may be obtained of what may be expected in any given
case. From the diagram the solubilities at any given temperature may be read.
Suppose, for example, the question is in regard to the quantity of potassium
274
INORGANIC CHEMISTRY
chlorate we may expect to obtain from 3 g. of potassium hydroxide dissolved in
7 g. of water (a 30 per cent solution). From the equation :
3C12 + 6KOH — >• KC1O3 + 5KC1 + 3H2O
6x56 122.5 5X74.5
we find that 336 g. of potassium hydroxide give 122.5 g. of chlorate and 372.5 g. of
chloride. Hence, by proportion, 3 g. will give about 1 g. and 3 g. respectively.
The solubility, read from the diagram, is the amount of the salt dissolved by 100
c.c. of water, for example, 56.5 g. of potassium chloride at 100°. Some of the
results are given in the form of a table :
POTASSIUM
CHLOBIDE.
POTASSIUM
CHLOKATE.
Amount formed from 3 g. KOH. . . .
Solubility at ) 100 c c Aq
3.0
56 5
1.0
56 5
100° in ) 7 c c. Aq
4.0
4.0
Solubility at ( 100 c.c. Aq ......
34.7
7.5
20° in j 7 c c. Aq
2.5
0 5
Solubility at ) 100 c c. Aq
28.0
3 3
0° in ( 7 c c. Aq
2.0
0 25
Thus, at 20°, at least 2.5 g. of the 3 g. of potassium chloride will remain dissolved,
while half of the potassium chlorate will crystallize out. If the solubilities are
examined, it will be seen that the potassium chlorate is even more easily obtain-
able in pure condition when calcium chloride takes the place of potassium
chloride.
Chloric Acid. — This acid may be obtained in solution in water,
by adding the calculated amount of diluted sulphuric acid to a solution
of barium chlorate :
Ba(C103)2 + H2S04 <=> BaS04 \ + 2HC103.
The barium sulphate, being insoluble, is removed by filtration (cf.
p. 265).
The solution may be concentrated (to about 40 per cent) by evapora-
tion, but must not be heated above 40°, as the acid decomposes near this
temperature. The resulting thick, colorless liquid has powerful oxidiz-
ing qualities, setting fire to paper (made of cellulose, C6H1005) which
has been dipped into it. It converts iodine into iodic acid, 2HC103 -+-
I2 — » 2HIOS + CLj. When warmed beyond 40° the acid decomposes,
giving chlorine dioxide and perchloric acid (see below).
Making of Equations Once More. — The equation for the last
action, although far from simple, may be made readily by use of a device
OXIDES AND OXYGEN ACIDS OF THE HALOGENS 275
which can always be applied where an oxygen acid gives an oxide. The
formula of the initial substance, chloric acid, may be written thus,
H2O,C1205, so as to show the anhydride (in this case imaginary) to
which it is related. Now the products are C102 and perchloric acid,
and the latter may be written H^ClgOf. Disregarding the elements
of water, we perceive that some G1205 becomes C1207, while the rest of
the Cl^ furnishes the oxygen for this change and itself falls to
2C102. Evidently one molecule undergoing the former change will
require two undergoing the latter in order that it may secure the two
units of oxygen :
2H20,C1205 -+ 2H,0 + 4C102 (+ 20), (1)
H20,C1A (+ 20) -» H20,C1207. (2)
Adding and dividing by 2, we have :
3HC103 -> H20 + 2C102 + HC104.
Chlorine Dioxide : Chlorous Acid. — Chlorine dioxide is a
yellow gas which may be liquefied, and boils at + 10°. The gas and
liquid are violently explosive, the substance being resolved into its
elements with liberation of much heat. It is formed whenever chloric
acid is set free, and hence it is seen when a little powdered potassium
chlorate is touched with a drop of concentrated sulphuric acid. Con-
centrated hydrochloric acid turns yellow from the same cause when
any chlorate is added to it. These actions are used as tests for chlo-
rates, and distinguish them from perchlorates (q.v.}. With water, chlo-
rine dioxide gives a mixture of chlorous and chloric acids, and with
bases a mixture of the chlorite and chlorate.
Perchlorates, Perchloric Acid, and Perchloric Anhydride.
- When heated, chloric acid and chlorates give perchloric acid (p. 274)
and perchlorates respectively. The chlorates also give oxygen at the
same time (p. 64) :
r2KC108->2KCl + 302,
\4KC103 -» 3KC104 + KC1.
These actions, like the three decompositions of hypochlorous acid,
are independent, and proceed simultaneously (p. 272). Their relative
speed, however, varies with the temperature, and the decomposition
into chloride and oxygen may completely outrun the other when a.
276 INORGANIC CHEMISTRY
catalytic agent like manganese dioxide is added (p. 65). When pure
potassium chlorate is heated cautiously, about one-fifth of it has lost
all its oxygen by the time the rest has turned into perchlorate. The
mixture may be separated by grinding with the minimum quantity of
water which will dissolve the chloride it contains. The perchlorate,
having at 15° less than one-twentieth of the solubility of the chloride,
will remain, for the most part, undissolved. The perchlorates are
much more stable (p. 119) than the chlorates, or hypochlorites : they
are all soluble in water, and they are used in making matches and fire-
works.
Pure perchloric acid explodes when heated above 92°. But, like
other liquids, its boiling-point is lower when its vapor is under reduced
pressure (p. 118). At 56 mm. pressure it boils at 39°, a temperature
at which hardly any decomposition is noticeable. Hence the acid may
be made by mixing potassium perchlorate and concentrated sulphuric
acid and distilling the mixture cautiously in a vacuum :
KC104 + H2S04 <=> KHS04 + HC10J .
To secure the requisite low pressure, the ordinary distilling apparatus
(Fig. 16, p. 38) is made completely air-tight, and is connected by a
branch tube with a water-pump.
Perchloric acid is a colorless liquid, which decomposes, and often
explodes spontaneously, when kept. A 70 per cent solution in water
is perfectly stable, however. Although it is an active oxidizing agent,
it is not so active as chloric acid, and does not oxidize hydrogen chlo-
ride in cold aqueous solution. When liberated by concentrated
sulphuric acid it does not at once give the yellow chlorine dioxide
(p. 275).
The anhydride (C12O7) may be prepared by adding phosphoric anhydride to
perchloric acid iu a vessel immersed in a freezing mixture, P2O5 + 2HC1O4 — » 2HPO3
+ C12O7. Phosphoric anhydride is often used in this way for removing the
elements of water from compounds. By gently warming the mixture, the perchlo-
ric anhydride can be distilled off. It is a colorless liquid boiling at 82° (760 mm.),
and exploding when struck or too strongly heated.
Oxygen Acids of Bromine. — No oxides of bromine have been
made, but the acids HBrO (hypobromous acid) and HBr03 (bromic
acid) and their salts are familiar.
By the action of bromine on dilute, cold potassium hydroxide
solution, the bromide and hypobromite are formed :
Br2 + 2KOH -» KBr + KBrO + H30,
OXIDES AND OXYGEN ACIDS OF THE HALOGENS 277
When the solution is heated, the hypobromite turns into bromate and
bromide. The actions are exact parallels of the corresponding ones
for chlorine.
Aqueous bromic acid may be made in the same way as chloric acid
(p. 274), or by the action of chlorine and water on bromine :
5C12 + 6H20 + Br2 -* 2HBr03 +10HC1.
The solution is colorless and has powerful oxidizing properties. Thus,
it converts iodine into iodic acid.
2HBr08 + I2 -* 2HI03 + Br2.
It appears, therefore, that iodine has more affinity for oxygen than has
bromine.
Oxides and Oxygen Acids of Iodine. — The following are the
acids and their corresponding salts :
[HIO Hypoiodous acid], [KIO Potassium hypoiodite] ,
HI08 Iodic acid, KI03 Potassium iodate,
[HI04 Periodic acid], NaIO4 Sodium periodate,
H6I06 Periodic acid, Na2H8I06 Disodium periodate.
The substances in parenthesis have not been isolated. There is one
oxide, I205.
lodates and Iodic Acid. — The potassium and sodium salts of
iodic acid are found in Chili saltpeter. They may be made, in much
the same fashion as are the chlorates and bromates (p. 272), by adding
powdered iodine to a hot solution of potassium or sodium hydroxide.
There is evidence that hypoiodites are formed in cold solutions, but
they change quickly to iodates.
Iodic Acid is formed by passing chlorine through powdered
iodine suspended in water. The action is parallel to that of chlorine
on bromine water. A still better way is to boil iodine with aqueous
nitric acid (q-v.)- The latter gives up oxygen readily, and is here
used solely on this account. Hence it may be omitted from the
equation, only the oxygen, of which it is the source, appearing :
. I, + H,0 + 50 -+ 2HI03.
278 INORGANIC CHEMISTRY
In both these actions the initial substances (including the excess of
nitric acid) and the products, with the exception of the iodic acid
itself, are all volatile. When the solution is concentrated by evapora-
tion, the iodic acid crystallizes. It is a white solid, perfectly stable at
ordinary temperatures, and can be kept indefinitely. At 170° it begins
to give off water vapor (2HI03 <=± H20 -+- ^2^5)' leaying the pentoxide
of iodine. The latter is a white crystalline powder which may be
raised to 300° before it, in turn, breaks up, giving iodine and oxygen.
In aqueous solution iodic acid is an oxidizing agent, but does not
part with its oxygen so readily as do chloric acid and bromic acid. It
oxidizes hydrogen iodide in dilute solution :
HI03 + 5HI -> 3H2O + 3I2,
all the iodine being liberated In this respect it resembles concen-
trated sulphuric acid (p. 237). Dilute sulphuric acid shows no
oxidizing qualities.
Various Acids Derived from One Anhydride. — Some acids are
related to their anhydrides as are hypochlorous acid (p. 268) and
sulphurous acid (p. 71). One molecule of the anhydride combines
with one molecule of water. In other cases, however, the proportion
of water may be less or greater than this. Thus phosphoric anhydride
(P205) takes up three formula- weights of water (p. 71). Now if
periodic acid were of the former type (H2O, I207 = 2HI04), its formula
would be HI04- It does form salts of this type, such as NaI04 and
AgI04. But the free acid is a deliquescent solid of the formula
HBI06(= 5H20, 1207), and the most easily prepared salt belongs to this
type. All types are called periodates, however, because their composi-
tions are all founded upon the same anhydride. The latter has not
itself been made. We usually speak of various acids and salts as
being derived from the same anhydride, the word " derived " being used
in a figurative and not a literal sense.
The difference between two acids HI04 and H5I06 is not at all the
same as between HI03 and HIO4. The latter would represent differ-
ent stages of oxidation, being derived from I2O5 and I207 respectively,
and accordingly would be named iodic acid and periodic acid. The
former differ only by 2H20, and an addition or subtraction of the two
elements of water in equivalent quantities is neither oxidation nor
reduction. Hence they are both periodic acids (see Phosphoric acid).
OXIDES AND OXYGEN ACIDS OF THE HALOGENS 279
Periodates and Periodic Acid. — Sodium periodate (NaI04) is
found in Chili saltpeter. When sodium iodate (NaI03) is dissolved
along with sodium hydroxide in water, and chlorine is passed into the
mixture, the sodium hypochlorite formed from the latter oxidizes the
iodate (NaIO3 -f O — > Nal04). But the somewhat insoluble salt which
crystallizes out is Na2H8I06 :
NaI03 + 0 + NaOH + H2O -* Na2HsI06.
Other salts may be made from this one.
An aqueous solution of periodic acid is obtained, like that of
chloric acid (p. 274), by the action of sulphuric acid on barium period-
ate. A white, very soluble solid (H5I06) remains when the liquid is
evaporated. When this is heated, water and oxygen are both given
off, and iodine pentoxide (I205) alone remains.
Chemical Relations. — The compounds of the halogens with
metals and with hydrogen diminish in stability, with ascending atomic
weight of the halogen, in the order: F(19), 01(35.5), Br (80), I (127).
Each halogen will displace those following it from this kind of combi-
nation. In the case of the oxygen compounds, the order of stability
is just the reverse, those of iodine, for example, being the only ones
which are reasonably stable. The order of displacement in such com-
pounds confirms this conclusion.
Amongst the oxygen acids of any one halogen, those containing
most oxygen are most stable. The salts are in all cases more stable
by far than the corresponding acids.
The halogens when combined with metals and hydrogen are univa-
lent (HI, KC1, etc.). It is clear, however, that, when united with oxy-
gen, their valence is higher. The maximum is. shown in perchloric
anhydride (C1207), where chlorine appears to be heptavalent.
The formulae of the acids might be written so as to retain the
univalence :
H-C1, H-0-C1, H-0-O-C1, H-0-0-0-C1,
H-O-0-O-O-C1.
But compounds in which we are compelled to believe that two oxygen
units are united are usually unstable (see Hydrogen peroxide), and we
should expect the instability would be greater with three and with four
units of oxygen in combination. Here, however, the reverse state
of affairs must be taken account of in our formulae, for HC104 is the
280 INORGANIC CHEMISTRY
most stable of the chlorine set. This reasoning, together with the
heptavalence in CLj07, leads us to assume the valence seven in per-
chloric acid (see Periodic system). The structural formulae (cf. p. 224)
of some of these substances are therefore often written as follows :
O O
II II
H-C1, H-0-C1, H-0-C1=0, Na-0-I=0.
II II
0 O
Exercises. — 1. Assign to its proper class (p. 187) each of the
actions mentioned in this chapter.
2. Knowing that potassium fluosilicate (K9SiF6) is insoluble, how
should you make chloric acid (p. 265) ?
3. Make the equation for the interaction of chlorine with calcium
hydroxide in hot water (p. 266). How should you make zinc chlorate
from zinc hydroxide (Zn(OH)2) ?
4. How should you make pure potassium hypochlorite from hypo-
chlorous acid (p. 268)?
5. On what circumstances would the possibility of making barium
chlorate by action of chlorine on barium hydroxide depend (p. 274) ?
6. Make the equations for : (a) the preparation of potassium bro-
mate ; (#) pure aqueous bromic acid ; (c) the interaction of iodine with
aqueous potassium hydroxide in the cold, and when heated.
7. Using the method given on p. 274, make the equations for the
interactions of chlorine dioxide with water, and with aqueous potas-
sium hydroxide.
CHAPTER XVII
DISSOCIATION IN SOLUTION
THE employment of interacting substances in the form of solutions
is so constant in chemistry, and the reasons for this are so cogent, that
we must now resume the discussion of the subject of solution (cf.
p. 145).
The present chapter will be devoted to giving the proofs that, to
speak in terms of the molecular hypothesis, the molecules of acids,
bases, and salts in aqueous solutions, are actually dissociated into parts
by the solvent. This will be shown by consideration, successively, of
certain peculiarities in the chemical behavior, the osmotic pressures,
the freezing-points, and the boiling-points of the solutions of these
substances. We shall see that these parts coincide in composition
with the radicals, and are called ions. Finally, the principles of
chemical equilibrium will be applied to the relations of the ions to
that proportion of the molecules which has remained undissociated.
Some Characteristic Properties of Acids, Bases, and Salts,
Shown in Aqueous Solution. — Acids all contain hydrogen (p. 93).
In aqueous solution, if soluble, they are sour in taste, they turn blue
litmus red, and their hydrogen is displaced by certain metals (p. 95),
and has the properties of a radical. By the last statement is meant
that it very readily exchanges places with other radicals in reversible
double decompositions (p. 264). Many other bodies, like sugar, kero-
sene, and alcohol, contain hydrogen also, but not one of them shows
all of these properties. Again, all salts are made up of two radicals,
and the reversible double decompositions into which they enter with
acids, bases, and other salts, consist in exchanges of these radicals.
Other substances may include the same combination of atoms, but in
their actions these groupings are often disregarded. Thus, sodium
chloride and silver nitrate exchange radicals completely (p. 13), and,
in dilute solution, hydrogen chloride and sodium hydrogen sulphate
do so partially (p. 180). But sodium chloride and nitroglycerine
C3H5(N03)g do not interact at all. The latter is not a salt, although it
contains the same proportion of nitrogen to oxygen as does any nitrate,
281
282 INORGANIC CHEMISTRY
Furthermore, it is chiefly in aqueous solution that these special
properties of acids, bases, and salts become apparent. Their behavior
is often quite different in the absence of this solvent. If, for example,
we mix together ammonium carbonate and anhydrous cupric nitrate,
and apply heat, a violent interaction begins. An immense cloud of
smoke and gas is thrown out of the tube, and the substance remaining
is either black or reddish, in parts, according to the proportions of the
substances employed. The residue contains cupric oxide, and some-
times red cuprous oxide (Cu20). The gas is tinged red by the pres-
ence of nitrogen peroxide (N02), while a more careful examination
would show that it contained carbon dioxide, nitrogen, nitrous oxide
(N20), water vapor, and perhaps still other products. The contrast,
when the substances are dissolved in water before being brought in
contact with one another, is very great. A pale-green precipitate is
formed at once, and rapidly settles out. On examination, this turns
out to be cupric carbonate, while evaporation of the solution furnishes
us with ammonium nitrate. There are only two main products, and
the essential part of the action in solution is represented by the
equation :
(NH,)2C03 H- Cu(N03)2 -> CuC03 + 2NH4N03.
In the interaction between the dry substances the molecules are com-
pletely disintegrated, and the whole change is very complex. In the
action in water no heating is required, the substances are neatly broken
apart, certain groups of atoms, which we call radicals, are transferred
as wholes from one state of combination to another, and the rearrange-
ment takes place in a machine-like manner. Contrasts like this
between the interactions of anhydrous and dissolved bodies are very
common. Thus, we have had occasion (p. 96) to mention the difference
between the action of metals on concentrated and on dilute sulphuric
acid.
Many compounds, however, do not show any change in behavior
when dissolved in water. Sugar, for example, is, as a rule, more readily
acted upon in the absence of any solvent. Then again, while water
is not the only solvent which has the effect we have just described,
the majority of solvents, if they affect chemical change at all, simply
retard it. Thus the union of iodine and phosphorus in the absence of
a solvent takes place spontaneously with a violent evolution of heat.
When the elements are dissolved in carbon bisulphide before being
mixed the action is much milder, although the product is the same
DISSOCIATION IN SOLUTION 283
(phosphorus tri-iodide). The diminution in the concentration of the
ingredients has decreased the speed of the action in the normal way
(p. 249). That water and some other solvents have a specific influence
tending to increase the activity of certain classes of substances, shows
that a special explanation of the phenomenon must be found.
Summing up these points we see that the peculiarity of acids, bases,
and salts in aqueous solution is that each compound always splits in
the same way. Thus, cupric nitrate always gives changes involving
Cu and ISTOg and never interacts so as to use CuN2 and 08, or Cu02 and
N02, as the basis of exchange. Similarly, acids always offer hydrogen
in exchange, and so nitric acid behaves as if composed of H and N08,
and sulphuric acid as if composed of 2H and SO4, and never as if
made up of HSO and H03, or H2S and 04. The sour taste and the
effect upon litmus seem to be properties of this easily separable hydro-
gen, for they are shown only by acids. The result is that we can
make a list of the units of exchange, such as H, OH, N08, C08, S04,
Cu, K, and Cl, employed by acids, bases, and salts in their interac-
tions. The molecule of each compound of these classes contains at
least two of them. Even when these units contain more than one
atom, their coherence is as noticeable within this class of actions, as
is the permanence of the atomic masses themselves in all actions.
The question raised in our minds is whether solution in water
alters the character of the molecule simply by producing a sort of
plane of cleavage in it which creates a predisposition to a uniform
kind of chemical change, or whether it actually divides the molecules
into separate parts consisting of the above units of exchange, and
leaves subsequent chemical actions to occur by cross-combination of
these fragments. The fact that the dissolved substances can be recov-
ered by evaporation of the liquid does not demonstrate that they have
not been changed temporarily while in solution. The alteration which
the water produces, whatever it be, will naturally be reversed when the
water is removed. Since our question involves nothing but the count-
ing of particles, the number of which would be much greater in the
event that actual subdivision of molecules is the explanation, it can
be answered by a study of the physical properties of solutions.
OSMOTIC PRESSURE.
In the earlier discussion of solution (p. 150) the condition of a
dissolved substance was viewed as akin to that of a gas. We con-
ceived the molecules of the dissolved substance as being distributed
INORGANIC CHEMISTRY
through
another,
because
solution
gases.
The
facts of
the space occupied by the solvent, as being separate from one
and as moving about independently of each other. This was
the phenomena of diffusion and osmotic pressure (p. 151) in
closely resemble those of diffusion and gaseous pressure in
invention of a suitable hypothesis for the explanation of the
osmosis presents some difficulties, but the facts themselves
are undoubted. It will conduce, therefore, to clearness
if we speak first of some things which may be observed
and are true, irrespective of any explanation.
Phenomena Produced by Osmotic Pressure. —
In order that the osmotic pressure (Gk. OKT/AOS, impulsion)
of the molecules of a diffusing body may be perceived,
a partition, which they are unable to traverse, must be
interposed between the solution and a contiguous mass
of the pure solvent (Fig. 57, p. 151). The partition must
be permeable by the solvent, however. Such a partition
is described as semi-permeable.
The general nature of the phenomena may be seen
by using an inverted thistle-tube (Fig. 73), with the
opening covered by a piece of the prepared peritoneal
membrane of oxen, and using sugar as the dissolved
body. This membrane is not strong enough to serve
for measurement. It is also somewhat permeable by
the sugar. But the water traverses it very easily, and
so an exhibition of the general result of a stricter test
is obtained quickly. The bulb of the tube is suspended
in pure water.
The water is able to pass freely through the mem-
brane in either direction, while the sugar is not. As
the result of the interchange of water, the liquid rises slowly but
steadily in the tube. The pure solvent always passes into the solu-
tion. If, further, two solutions of different concentrations of the same
substance are employed, then, invariably, water passes from the more
dilute solution into the more concentrated one through the membrane.
There is apparently a tendency for the water so to distribute itself
that the solutions may eventually become equal in strength. The
water passes from a dilute solution, leaving it more concentrated than
before, into a more concentrated solution, rendering it more dilute.
Fro. 73.
DISSOCIATION IN SOLUTION 285
These phenomena were first studied by Pfeffer (1877), a botanist,
who used certain plant cells for the purpose. The cell content in-
cluded a liquid containing various salts in solution, and a protoplas-
mic layer which was not attached to the cell wall. This protoplasmic
layer behaved like a semi-permeable membrane. When such cells
were immersed in a concentrated solution of any substance, the water
passed from the interior of the cell to the solution, and by means of
a microscope a shrinkage of the protoplasmic layer away from the cell
wall could be observed. Conversely, when such cells were placed in
pure water, or a solution of a very dilute nature, water passed from the
outside into the interior, and the protoplasmic layer was distended so
as to fill the corners completely. The distension of the cells of droop-
ing flowers, when their stems are placed in water, and the consequent
revival, is a familiar illustration of the same sort of thing. All solu-
tions which produced neither the one effect nor the other on a given
set of plant cells, were named is-osmotic. The osmotic pressures of
their contents were the same as the pressure of the cell fluid.
Professor Crum Brown has devised an arrangement which exhibits the action
of a perfectly semi-permeable membrane very strikingly. A concentrated solution
of calcium nitrate is shaken with a small amount of phenol (carbolic acid), so as
to become saturated with the latter, and the mixture is then poured into a tall,
narrow cylinder. The phenol rises and floats upon the surface of the calcium
nitrate. The amount of phenol should not be more than sufficient to saturate the
liquid and give a layer a few millimeters in thickness. Distilled water, also satu-
rated with phenol, is cautiously introduced above all. The water on both sides of
the layer of phenol is soluble in phenol, and consequently, by dissolving in this
and passing out on the other side, can traverse the partition. The calcium nitrate,
however, which is here the dissolved substance, cannot traverse the phenol in
which it is not soluble. The phenol therefore constitutes a perfect semi-permeable
membrane. If the level of the lower side of the phenol is marked on the outside
of the cylinder by means of a strip of paper, it will be found, as the arrangement
is watched from day to day, that the water passes through the phenol into the solu-
tion, and the phenol rises higher and higher, until finally it surmounts all the rest
of the liquid.
The Phenomena a Logical Consequence of Semi-Permea-
bility. — The passage of the water into the solution in which the
greater osmotic pressure exists seems at first paradoxical. We must
remember, however, that the system, consisting of the liquids on each
side of the membrane, can be in equilibrium only when the osmotic
pressure on the two sides is identical. But the equalization of the
osmotic pressures cannot take place by the passage of part of the
solute from one side to the other. The membrane has been taken,
286 INORGANIC CHEMISTRY
purposely, of such a nature that the dissolved substance is unable to
traverse it. The equalization must occur, therefore, in the only other
possible manner, namely, by the passage of the solvent in the other
direction.
An imitation of this behavior may easily be exhibited by the use of gases. A
piece of peritoneal membrane is stretched across the mouth of a thistle-tube and
moistened with water. The tube, which has been bent in U-f°rm to serve as a
manometer, contains a small amount of some colored liquid, whose motions will
exhibit any change in pressure in the interior. When an inverted cylinder of
ammonia gas is placed round the head of the thistle-tube, the ammonia gas dis-
solves in the water on the membrane until this water is saturated, that is, until
the ammonia molecules leaving the water are as numerous as those entering
it. It will be seen, however, that the ammonia solution really has two surfaces,
one of them towards the interior, and the ammonia particles must eventually leave
both surfaces at the same rate at which they are landing upon one of them. The
ammonia gas being at the pressure of the atmosphere, the particles of ammonia
leaving the film will produce a tension of one atmosphere of ammonia over each
surface. Thus ammonia gas will be transferred from the cylinder to the interior
of the thistle-tube until its partial pressure in the latter is equal to that in the
former. The membrane is semi-permeable, since, of the air and ammonia con-
tained in the thistle-tube, only the ammonia can traverse the film. The con-
tents of the thistle-tube therefore correspond to the solution, air being the solute
and ammonia the solvent. The original air in the apparatus was at a pressure of
one atmosphere, but the ammonia, although under no greater pressure, enters
nevertheless. Indeed, it would continue to do so until the pressure inside became
equal to that of the ammonia outside plus the original pressure of the air, a total of
two atmospheres. The case corresponds to that of water entering a solution whose
osmotic pressure is one atmosphere. It enters until the contents of the apparatus
are under a pressure one atmosphere greater than that existing outside.
Measurement of Osmotic Pressure. — It will be seen that the
whole phenomenon rests upon the fact that the membrane used is
permeable by one of the constituents only. The preparation of a
vessel of sufficient strength, and possessing walls with the maximum
permeability by water and the minimum permeability by dissolved
substances, presents great difficulties. A device of Pfeffer's is still
found to be the best. A cylinder of porous porcelain, much like a
Pasteur filter-tube, is treated so that its pores are partially filled with
a gelatinous precipitate of cupric ferrocyanide (q.v.~).
The porous cylinder, after removal under the air-pump of the air which its walls
contain, is placed in a solution of cupric sulphate. Its interior is then filled with a
solution of potassium ferrocyanide. When these two liquids meet by diffusion
inside the wall, they interact, producing a dense precipitate of the substance above
mentioned :
2CuS04 + K4Fe(CN)6— >Cu2Fe(CN)6 + 2K2SO<.
DISSOCIATION IN SOLUTION
287
If such a prepared vessel, after being filled with a one per cent
sugar solution, could be closed by a piston (e.g. Fig. 46) and be placed
in pure water, it would be found necessary to place weights on the
piston to prevent an upward movement, due to access of water to the
interior through the walls. Finally a weight would be found that would
just balance the inward tendency of the water. With more weight
than this, water would be squeezed out through the pores ; with less,
the water would force its way in and the piston would rise. When
this weight has been placed in position, the
water inside and outside, having reached a
condition of equilibrium, must be exerting
equal pressures on each side of the wall of the
vessel. Hence, the excess of pressure inside
must be due to the osmotic pressure of the
dissolved sugar. It cannot be due to the water
itself, for that is able to escape through the
pores. The weight opposing the osmotic press-
ure at 15° in the case of a one per cent sugar
solution is found to be about 0.7 kg. for every
sq. cm. of the exposed surface. Since 1.03 kg.
per sq. cm. equals 760 mm., this would indicate
a pressure of 760 x 0.7 -H 1.03, or 516 mm.
(0.68 atmospheres).
In practice a small bent tube opening into
the cylinder is used as a manometer (Fig. 74).
The other end of the tube is closed, and some
air is confined in this end by mercury. The
diminution in the volume of the air registers
the pressure. The smaller tube, drawn out to
a point, is used for filling the cell with the
solution and is then sealed before the blow-
pipe. The whole apparatus is immersed in
a large bath of water whose temperature can
be maintained constant during the experiment. Concordant readings
are hard to get in consequence of difficulties inherent in the prepara-
tion and use of the apparatus, but the general relations of the results
can be stated in a very simple form.
Ten years after Pfeffer's experimental work, van 't Hoff first formulated the
laws of osmotic pressure. He showed that the general analogy between the gase-
ous state and the state of solution could be developed so as to exhibit a complete
FIG. 74.
288
INORGANIC CHEMISTRY
correspondence between the laws of both. His conclusions (1887) were founded
partly on Pfeffer's results and partly on supplementary experiments, and are given
in the three following paragraphs (see, also, appendix to this chapter).
Osmotic Pressure and Concentration. — A part of one of
Pfeffer's sets of experiments will show the relation in this respect :
PKB CENT OF SUGAR.
OSMOTIC PRESSURE.
PER CENT OF SUGAR.
OSMOTIC PRESSURE.
1
2
553 mm.
1010
4
6
2082 mm.
3075
The osmotic pressures of a series of solutions of the same substance
are proportional to their concentrations. The form of the law is the
same as that of Boyle's law (p. 81). The pressures are in general the
same for the same concentrations whatever solvent is used.
Osmotic Pressure and Temperature. — Pfeffer, using a one per
cent solution of sugar at two different temperatures, found the pressure
at 14.2° to be 510 mm. and at 32° to be 544 mm. Not only does osmotic
pressure change, with alteration in temperature, in the same direction
as does gaseous pressure, but the changes can be expressed by Charles'
law (p. 87). The osmotic pressure increases in proportion to the
absolute temperature. A gas which at 14.2° C. exhibits a pressure of
510 mm., at 32° exercises a pressure of 542 mm.
An Analogue of Avogadro 's Hypothesis. — Still more interest-
ing is the fact that, if we compare the concentrations of different solu-
tions -which at the same temperature exhibit equal osmotic pressures,
we find that they contain equal numbers of molecules of the dissolved
substance in equal volumes. Thus, if we dissolve one mole (342 g.)
of sugar (C12H22On) and one mole (74 g.) of methyl acetate (CH3C2H802)
in equal volumes of water, we have taken equal numbers of molecules
of the two substances, and the osmotic pressiires which the solutions
exhibit are found to be equal.
As will be seen below, certain substances in certain solvents, par-
ticularly in water, exhibit pressures which are greater than this hypoth-
esis would permit.
The closer study of the figures enabled van 't Hoff (1887) to state
the most interesting fact of all : The osmotic pressure exercised by a
substance in solution is identical in value with the gaseous presssure
DISSOCIATION IN SOLUTION 289
which it would exhibit if the same quantity of it were contained as
a gas in the same volume at the same temperature. For example,
44 g. of carbon dioxide in the gaseous condition fills the G.M.V.
(22.41.), and at 0° exercises a pressure of one atmosphere. We find
that when we dissolve the same quantity of the same substance in
22.4 1. of any solvent at the same temperature, it exercises one atmos-
phere of osmotic pressure. Certain substances, however, particularly
when dissolved in water, exhibit greater pressures than this (see
below).
Determination of Molecular Weights. — It is evident that we have
here an experimental method which may be used in measuring the molecular
weight, and is applicable to substances which cannot be converted into vapor. All
that is necessary is to dissolve a weighed amount of the substance in a known
amount of water, or some suitable solvent, and by means of the apparatus
described above to measure the osmotic pressure at some fixed temperature. From
the result, by means of the laws corresponding to those of Boyle and Charles, we
may calculate the concentration of the solution which would have given one atmos-
phere pressure at 0°. That quantity of substance which would give this concen-
tration in 22.4 1. of the solvent is then the molecular weight (cf, p. 199). The time
required for measurements of osmotic pressure and the experimental difficulties
alone prevent the employment of this method in practice.
Osmotic Pressure and Dissociation in Solutions. — What
inference is to be drawn in the cases in which abnormally high osmotic
pressures are observed ? In view of the fact that the pressure is sup-
posed to be produced by the impact of the particles, and depends on
the number of them in the given volume, we must infer that where the
pressure is greater, more particles are present in the given volume
than we had supposed. In other words, dissociation of the original
molecules must have occurred. This phenomenon is observed when-
ever acids, bases, or salts in aqueous solution are under observation.
Thus a solution of sugar, which does not belong to these classes, con-
taining 342 g. in the G.M.V., exhibits the normal osmotic pressure of
one atmosphere at 0°. A solution of one molecular weight of potas-
sium chloride (74.5 g.) in the same volume of water, however, exhibits
an osmotic pressure of about 1.88 atmospheres at 0°. The greater
pressure must be due to the fact that, although the number of mole-
cules of potassium chloride taken is the same as in the case of sugar,
the number of actual particles whose impacts constitute the pressure is
greater, — is, in fact, 88 per cent greater. Now the multiplication of
particles from potassium chloride molecules can occur only by their
290
INORGANIC CHEMISTRY
dissociation into particles of K and Cl by a chemical change rep-
resented by the equation KC1 ^± K + Cl. In this case, seeing that
each original molecule can give but two particles, the excess of press-
ure indicates that 0.88 (88 per cent) of the molecules of potassium
chloride have been broken up. Comparison shows that the degree of
dissociation for equi-molar solutions of different acids, bases, or salts
varies widely. For the same substance, it is always relatively greater
in dilute than in concentrated solutions.
It will be seen that we have thus a purely physical and perfectly
. independent confirmation of the indications
already found in the chemical behavior of
substances of this kind. In practice, on
accoiint of the experimental difficulties,
this method is not used for measuring the
degree of dissociation.
DEPRESSION IN THE FREEZING-POINT
OF A SOLVENT.
Measurement of Freezing- Points.
— The task consists in measuring exactly
the temperature at which a previously
weighed quantity of the solvent freezes,
and then, after dissolving in it a known
weight of some soluble substance, deter-
mining the freezing-point once more. The
absolute values of these two points are not
required, it is simply the difference be-
tween them that has to be known with
exactness (cf. p. 163). By means of a
very delicate thermometer (Fig. 75) having
only six degrees on the whole scale, the
temperature of the freezing liquid may
be read to one one-thousandth of a degree.
A. reservoir at the top enables us to add to, or subtract from, the
mercury contained in the bulb and column, and so the same instrument
may be used with solvents having widely different freezing-points.
When water is being employed as the solvent, the outer jar -must be
filled with a freezing mixture of ice and water containing salt. With
solutions in benzene, ice and water are used alone. To avoid super-
FiG. 75.
DISSOCIATION IN SOLUTION 291
cooling, the solvent or solution must be vigorously stirred after it
has been cooled down to a point just below the freezing-point.
Laws of Freezing-Point Depression. — The depression is
directly proportional to the weight of dissolved substance in a
given amount of the solvent. The depression is inversely propor-
tional to the amount of solvent. Thus, if we double the concen-
tration of the solution, the depression in the freezing-point is doubled.
Further, equal numbers of molecules of different solutes in the same
quantity of solvent give equal depressions. Or, in other words, the
depression is proportional to the concentration of the molec\iles of
the solute. Thus, solutions containing 342 g. of sugar (C12H22OU), or
46 g. of alcohol (C2H60), or 74 g. of methyl acetate (CH8C2H3O2), in 1000
g. of water, show a depression below the freezing-point of water of
1.89° in each case. This depression produced by a mole of the solute
in 1 1. of water is called the molecular depression constant and has a
different value for each solvent. For solutions of the same molecular
concentration in benzene the depression is 4.9°, in phenol (carbolic
acid) 7.5°. Combining these facts in one expression :
The observed depression \ _ Wt. of Solute 1000
in an aqueous solution ) Mol. Wt. of Solute Wt. of Solvent.
For other solvents, the corresponding value of the depression constant
must be substituted for 1.89°.
These principles may be expressed mathematically in a form which is con-
venient for use. If A represent the depression in any actual experiment, 5 the de-
pression produced by one molecular weight in 1000 grams of solvent, W the
weight of the substance, M its molecular weight, and g the weight of the solvent in
grains, then :
W x 1000
A = 8 X — TT-
M x g
In the case of water, as we have seen, 5 is 1.89°. For each solvent the value of 8
must be determined by means of a substance of known molecular weight.
These laws describe the facts most exactly when the solutions are
dilute. They hold only when there is no chemical interaction between
solute and solvent. Even so, however, acids, bases, and salts dissolved
in water present many apparent exceptions and must be discussed
separately.
292 INORGANIC CHEMISTRY
Determination of Molecular Weights. — When the depression
constant of a solvent has once been ascertained by means of a sub-
stance of known molecular weight, this method may be used for deter-
mining the molecular weight of other substances which are soluble in
the same liquid. All the other factors can be observed and substituted
in the formula. This method is especially useful when the substance
. cannot be converted into vapor without undergoing decomposition (see
Hydrogen peroxide).
Freezing- Points and Dissociation in Solution. — The sub-
stances which present the most conspicuous exceptions to the above
rules are acids, bases, and salts in aqueous solution. With most of
these, the depression produced is greater than we should expect from
the concentration of the solution. Thus, in an actual experiment, two
equi-molar solutions were compared. One contained one mole (74 g.)
of methyl acetate, and the other one mole (58.5 g.) of sodium chloride,
each dissolved in 2000 g. (2 liters) of water. The freezing-points ob-
served, on the arbitrary scale of the thermometer, were :
Pare water 3.580° Pure water 3.580°
Solution of methyl acetate . 2.610° Solution of salt . . . 1.902°
Depression 0.970° Depression 1.678°
0.970°
Excess depression by salt 0.708J
The solution of methyl acetate, as it contained only 0.5 moles of the
solute per liter of water, showed, as it should do, about half the aver-
age molecular depression (1.89°, p. 291). This is typical of the class of
substances showing normal behavior. Sugar, alcohol, and hundreds
of other substances, in solutions of the same molar concentration,
would have given the same value.
The freezing-point of the salt solution, however, was much lower.
If this solution had contained the same concentration of dissolved
particles as the other solution, its depression would have been 0.970°
likewise. The number of particles must therefore have been greater
than we should have expected from the number of molecules taken.
In other words, a portion of the molecules of the salt must have been
broken up, and the excess depression, 0.708°, must have been due to
the extra particles produced by dissociation. Now sodium chloride
molecules cannot give more than two particles each, and the depression
is proportional to the number of particles. It follows, therefore, that
l$%, or 0.732 (73.2 per cent) of the molecules were dissociated.
DISSOCIATION IN SOLUTION 293
This result is typical also. Acids, bases, and salts of which one
mole is dissolved in two liters of water, are found to give irregular
values, all more or less in excess of 0.970°. Those which contain but
two radicals, like sodium chloride (NaCl) and potassium nitrate
(KNO,), give values between 0.970° and 2 x 0.970°. Substances like
calcium chloride (GaCL,) and sodium sulphate (Na2S04) give depres-
sions approaching three times the normal value: their molecules
contain three radicals. The excess depression depends, therefore,
upon the number of particles which each molecule can furnish, and
upon the proportion of all the molecules which is dissociated into
these fragments.
In the case of an acid, base, or salt, the depression is not strictly
proportional to the concentration. Thus, one mole of salt in four liters
of water does not give half the depression of the two-liter solution
(0.839°) but somewhat more (about 0.844°). The same method of
calculation indicates, therefore, a greater degree of dissociation
(about 79 per cent) in the more dilute solution (see Ionic equilib-
rium, below).
Acids, bases, and salts, so far as they are soluble in materials like
toluene, benzene, chloroform, and carbon bisulphide, exhibit simply
normal depressions in these solvents. It appears, therefore, that dis-
sociation does not take place in many solvents. In common experi-
ence it is encountered only in solutions in water, and, perhaps,
alcohol.
Soiling-Points and Dissociation in Solution^ — If space per-
mitted, a series of statements might be made in regard to the boil-
ing-poiuts of solutions (cf. p. 162) which would be closely parallel to
those about freezing-points. The boiling-point, as we have seen, is
elevated, however, by the introduction of a foreign body. Thus, when
water is the solvent, one mole of a solute in 1000 g. of the solvent
normally raises the boiling-point 0.52° (that is, from 100° to 100.52°).
But acids, bases, and salts form an exception to this rule, as before,
and the excess elevation which they give is a measure of the degree of
dissociation.
Comparison of the Results of the Three Methods. — When we
measure the osmotic pressure, the freezing-point depression, and the
elevation in the boiling-point of the same solution, and calculate the
degree of dissociation from the result of each measurement, we find
294 INORGANIC CHEMISTRY
that the values obtained are usually identical, within the limits of
error to which the methods are liable. Indeed, the theory of this
subject, developed by van't Hoff, enables us to connect the osmotic
pressure by a mathematical relation with the other two phenomena, and
to calculate any one of the three from any other.
The connection between the three sets of phenomena cannot be explained
here. It is treated in all works on Physical Chemistry. It may be pointed out,
however, that, in one essential respect, experiments in osmotic pressure, and in
the freezing and boiling of solutions, are all alike. The perception of osmotic
pressure involves a partition which the solvent alone can pass, and the osmotic
pressure for a given solution is the one required to force the solvent out. In
freezing a solution, pure ice is separated, and so a similar extrusion of a part of the
pure solvent is effected. In a boiling solution, for which the above rules hold, the
vapor is composed of the pure solvent, and the solute remains behind. The rela-
tion between the three operations lies in the fact that in each case the same thing,
namely, the separation of a part of the solvent, is done. Each method effects this
in a different way. But the expressions representing the work done, in terms of
the factors which define the work in each case, can be equated in pairs and the
required relation established. Thus the molecular depression of the freezing-point,
or the molecular elevation in the boiling-point, as we have defined them, is equal
to 0.002 T2 -i- g, where T is the absolute temperature of the freezing- or boiling-
point, and q is the heat of fusion or vaporization, as the case may be. Water, for
example, freezes at 273° abs., and its heat of fusion is 79 cal. per gram, from which
the calculated molecular depression, 0.002x2732 ~- 79, or 1.88°, is obtained. Simi-
larly, using the boiling-point, 373° abs., and the heat of vaporization, 537 cal. per
gram, we calculate the molecular elevation of the boiling-point to be 0.518°.
It ought to be added that abnormally small osmotic pressures, freezing-point de-
pressions, and boiling-point elevations, are also frequently observed. This occurs,
however, almost wholly in non-aqueous solvents, such as benzene. It is shown
particularly by substances containing oxygen, and is even noticed in the case of
acids, bases, and salts. By parity of reasoning we infer that in these cases associa-
tion (cf. p. 242) of the molecules has occurred, and that the physical unit of the
solute in these solvents is larger than the ordinary molecule.
THE APPLICATION OF THESE CONCLUSIONS IN CHEMISTRY.
The Constitution of Solutions of Acids, Bases, and Salts. — The
composition of solutions which are normal or abnormal, in respect to
osmotic pressure, freezing-point, and boiling-point, may be shown thus :
SOLUTES.
DISSOLVED IN WA-
TER, ALCOHOL, ETC.
DISSOLVED IN
TOLUENE, CHLOBO-
FOBM, ETC.
Acids
bases salts
Abnormal
Normal
Other
substances
Normal
Normal
DISSOCIATION IN SOLUTION 295
It appears that water and soine other solvents have the power of
breaking up the molecules of acids, bases, and salts and of holding the
fragments apart from one another and hindering their reunion. In
consequence of this, our view of the nature of an aqueous solution of
hydrogen chloride (HC1), or common salt (NaCl), or sodium hydroxide
(NaOH), or any of the substances of the classes which these represent,
may now be stated in definite terms. Such a solution contains, besides
undivided molecules of the solute, at least two other kinds of material,
H, Na,* Cl, OH, etc., which result from the breaking up of the molecules.
We shall see that these subdivisions of the original molecules have
distinct physical and chemical properties of their own. The descrip-
tions of the " properties " of the solutions, as they used to be given in
chemistry, were really a confused statement of the properties of the
different constituents of a mixture of molecules and their fragments.
Thus the indications of dissociation found in the chemical behavior
of acids, bases, and salts (p. 282) are fully confirmed by a study of
the physical properties of their solutions.!
The suggestion that the multiplication of particles takes place by interaction
of the salt with part of the water, NaCl + H2O ^± NaOH + HC1, resulting in the
production of two molecules of dissolved matter from one. is open to several fatal
objections. In the case of a highly dissociated salt, according to this explanation,
the mixing of the acid and base in dilute solution should result in no particular
change and give rise, therefore, to no development of heat. But the heat of neutral-
ization is very great in such cases. This is an example of a stochastic hypothesis
(p. 142), be it noted, and its verity or falsity can be put to the test at once. Its
inapplicability is further seen in the fact that it cannot explain the dissociation of
acids and bases themselves.
The free radicals, of whose existence we have thus become convinced,
constitute a new set of materials. Thus the hydrogen radical of acids,
although a form of uncombiiied hydrogen, differs totally from the gas
which is composed of the same material. The latter has no sour taste
or effect upon litmus. It is very slightly soluble in water, while the
hydrogen radical exists as a separate substance only in solution. Again,
substances with the composition of the radicals N03 and S04 are not
known at all except in solutions. The chief peculiarity of these sub-
stances is that a solution cannot be made which contains less than two
* The objection that separate atoms of sodium could not remain free in water,
will be disposed of later.
t Recent observations, showing that in some cases rapid double decompositions
of the normal kind take place in solutions which exhibit no physical evidence of the
existence of dissociation, demonstrates that it would have been unsafe to infer dis-
sociation from chemical evidence alone.
296 INORGANIC CHEMISTRY
kinds of them side by side. A niche, therefore, must be created in our
molecular hypothesis to receive these new substances.
Nomenclature : The Ionic Hypothesis. — Our chemical mole-
cules are the units of material in the gaseous condition (p. 198). Evi-
dently smaller units, which may nevertheless contain more than one
atom, must be assumed to exist in solution. These units, for a reason
that will appear later, are called ions, and their composition corresponds
to that of the radicals. The dissociation of molecules into ions is
named ionization. The substances of the three classes which alone
are ionized may be designated ionogens. Since ions are discrete parti-
cles, they are, in all physical respects, molecules. Thus we speak of
the molecular concentration of ionic hydrogen, just as we do of that of
dissolved or of gaseous hydrogen.
The solution of an ionized substance is called an electrolyte (q.v.), and often
this term is applied also to acids, bases, and salts themselves, because, when dis-
solved, they produce electrolytes. This is rather a confusing metonymy, however,
because these bodies by themselves are not conductors. This use of the term also
introduces obscurity because it connects the ionization with electrolysis and always
conveys the impression that the latter produces the former. The electrolytic prop-
erty of ions is only one amongst many special properties of electrolytes, and the
majority of these properties are chemical and have nothing to do with electrolysis.
Hence we have preferred the more general word " ionogen."
The radicals and their chemical behavior are real, and all the pecul-
iarities of aqueous solutions of acids, bases, and salts are experimental
facts. Ions, however, like corpuscles, atoms, and molecules, are part
of our great system of formulative hypotheses and are added to it in
order to maintain its self-consistency. We apply that part of the
hypothetical system known as Avogadro's hypothesis to solutions, and
finding portions of molecules which do not exist in the gaseous con-
dition and which have special properties of their own, a new class of
unit masses has to be established. Molecules are units which are not
commonly disintegrated by vaporization (p. 198) ; ions, those which
are not commonly disintegrated in double decomposition in solution;
atoms, those which are not commonly disintegrated in any chemical
action. But there are exceptions in each of the three cases. The ionic
hypothesis was first suggested by Arrhenius (1887) immediately after
the publication of van 't Hoff's correlation of the facts about osmotic
pressure (p. 288).
It is worth noting that the quantities expressed by the formulae Al, Ca, and
K, when existing as ions, produce equal osmotic pressures, and have equal effects
DISSOCIATION IN SOLUTION 297
upon the freezing- and boiling-points. This is a further justification for our choice
of chemical unit quantities of the elements (atomic weights), for the atomic weights
have these properties in common, and equivalents, of course, do not (cf. p. 210).
Ionic Equilibrium. — Since the ions are chemically different from
their parent molecules, their formation represents a variety of chemi-
cal change. The change does not involve any chemical interaction with
the water, of the nature of hydrolysis, for cases in which this takes plaqe
are expressly excluded from consideration. It is simply a dissocia-
tion, i.e. reversible decomposition of the dissolved substance.
From the fact that the proportion of molecules ionized is shown to
become greater as more and more of the solvent' is added, and that
removal of the solvent diminishes the proportion of ions to molecules,
and finally leaves us the substance entirely restored to the molecular
condition, we know that this is a reversible action and therefore a true
dissociation. The molecules and their ions adjust themselves like the
constituents in any case of chemical equilibrium. In the cases above
mentioned we should have the following actions taking place : *
HC1 <=> H* + Cl' NaCl «=* Na* + CF NaOH <=>. Na* + OH'.
These equilibria are all of precisely the same nature as that of
phosphorus pentachloride vapor (p. 255), and the discussion of the
latter should be reexamined and applied by the reader. The sole
difference is that here change in volume is effected, not by compression
or by release of pressure, but by removing or adding water. The
adjustment to a condition of equilibrium, however, seems to be instan-
taneous where ions are concerned, while in other chemical actions it
always takes a perceptible, and often a considerable interval of time.
Using Cv C2, and Cs for the molecular concentrations (numbers of
moles per liter) of the molecules, and the two ions, respectively, we have
an equilibrium constant (cf. p. 254), in this case called the ionization
constant :
C, X Ca
When we dissolve a single substance which gives only two ions, the
molecular concentrations of the ions are necessarily equal. Hence,
C 2
in such a case, — f- = K. When some other ionogen with a common
ci
ion is present, however, the values of (?„ and Ca will be different.
* The symbols Na', Cl', etc., are used to indicate that these are ions, and not
identical with atoms, Na, Cl, etc. The negative radicals and hydroxyl are distin-
guished thus, NO,', OH', and the others by a dot, H*, K'.
298 INORGANIC CHEMISTRY
Considering the form of the above mathematical expression, it will
be seen that when the degree of ionization is great, C2 and Cs are larger
than Cv and the value of K, the ionization constant, will be great.
On the other hand, in the case of feebly dissociated substances, the
value of K will be small. Furthermore, if by the addition of more
water we diminish all the concentrations, this will momentarily affect
the numerator more than the denominator (cf. p. 298). In order, there-
fore, that the value of the whole expression may remain constant, the
concentrations of the ions, represented by C2 and Cs, must become
greater, and can only do so at the expense of the concentration of the
undissociated molecules, represented by Cr Our formula, therefore,
represents successfully the fact that dilution, which diminishes the con-
centration of all the substances, produces a greater degree of ionization.
A more general form of treatment will be required later. If a be
the number of moles, say of acetic acid, originally taken, v the
volume of the solution in liters, and x the number of moles ionized,
then the molar concentrations at equilibrium will be
=
v v
Therefore,
/x\2 a — x x2
Or
v ' (a — x} v
This is known as Ostwald's dilution formula.
Exercises. — 1. A one per cent sugar solution gives an osmotic
pressure of 516 mm. at 15°. What is the molecular weight of sugar ?
Assume that the sp. gr. of the solution is 1.
2. What gaseous pressure would be exerted by a gas of the same
molecular concentration as a one per cent solution of sugar at 15°
(p. 288) ? Compare the answer with the osmotic pressure of the
solution.
3. What depression in the f.-p. of water will be produced by
dissolving 10 g. of bromine in 1 kg. of this solvent ?
4. WThat depressions in the f.-p. of benzene and of phenol would
be produced by 10 g. of bromine to 1 kg. of the solvent, if no chemi-
cal action took place ?
5. What is the molecular depression-constant of a solvent in
which 5 g. of iodine in 500 g. of the solvent lowers the f.-p. 0.7°?
6. What is the degree of dissociation of zinc sulphate if 5 g. of it
DISSOCIATION IN SOLUTION
299
dissolved in 125 g. of water produce a lowering of 0.603° in the f.-p. ?
What is the molecular concentration of each of the three substances
present in this solution ?
7. What will be the approximate b.-p. of a solution of common
salt, saturated at 100° (p. 157) ? Assume that the solute is 80 per
cent dissociated.
Appendix: Recent Measurements of Osmotic Pressure. —
Measurements of osmotic pressure by Morse and Frazer, published since the fore-
going was written, are the only ones yet made which deal with pressures of more
than three or four atmospheres.
The actual measurement is made according to the general method described on
p. 287, that is, by means of a porous porcelain cup or cylinder which contains
the semi-permeable membrane within its walls, and is attached to a manometer
containing an inclosed volume of air over mercury.
The membrane, however, is deposited electrolytically, instead of by diffusion.
The method of procedure is briefly as follows : The porous cylinder, after the
removal of the air from its walls by "electrical endosmose," is surrounded by a
copper electrode, and both cup and electrode are immersed in a 0.1 N solution of
copper sulphate. The other electrode, the anode, is placed within the cup, which
is filled with a 0.1 N solution of potassium ferrocyanide. A current of electricity
with an electromotive force of 110 volts is then passed through the solutions and
the porous wall of the cylinder, from the copper to the platinum electrode, and the
membrane of copper ferrocyanide is deposited either upon the interior surface of
the cup, or within its walls. By proceeding in this way, only from one to three
hours, instead of several days, are required for the formation of a suitable membrane.
The following table contains the results of some measurements made with
solutions of cane-sugar in water. The proportionality between concentration and
osmotic pressure is evident.
MOLES SUGAR IN
1000 GRAMS WATER.
OSMOTIC PRESSURE
(ATMOSPHERES).
MOLES SUGAR IN
1000 GRAMS WATER.
OSMOTIC PRESSURE
(ATMOSPHERES).
0.2
0.4
4.83
9.72
0.5
1.0
12.15
24.46
From data already obtained by Morse and Frazer, it seems very probable that
the statement of one of van 't Hoff's laws (p. 288) will have to be modified, for it is
found that "cane-sugar, dissolved in water, exerts an osmotic pressure equal to
that which it would exert if it were gasified at the same temperature and the volume
of the gas were reduced to that of the solvent in the pure state" (Amer. Chem.
Jour., July, 1905).
CHAPTER XVIII
OZONE AND HYDROGEN PEROXIDE
A FRESH, penetrating odor, resembling that of very dilute chlorine,
was noticed by van Marum (1785) as being perceptible near an elec-
trical machine in operation. Schonbein (1840) showed that the odor
was that of a distinct substance, which he named ozone (Gk. o£«v,
to smell), and he discovered a number of ways of obtaining it. It is
very questionable whether there is any ozone in the air, excepting
temporarily in the immediate neighborhood of a natural or artificial
discharge of electricity.
Preparation of Ozone. — The most satisfactory way of prepar-
ing ozone (08) is to allow electric waves to pass through oxygen.
The apparatus (Fig. 76) consists of two co-axial glass tubes, between
which the oxygen flows. The waves are generated by connecting an
FIG. 76.
outer layer of tinfoil on the outer tube, and an inner layer of tinfoil
in the inner tube with the poles of an induction coil. With dry, cold
oxygen, about 7.5 per cent of the gas is easily turned into ozone.
Under the best conditions this proportion cannot be much exceeded.
Ozone is found in the oxygen generated by electrolysis of dilute
sulphuric acid (p. 95). Some of it is produced when sulphuric acid
acts upon oxides which, with this reagent, liberate oxygen, e.g.,
2Ba02 + 2H2S04 -» 2BaS04 -f 2H,0 -f 0,. It arises during the slow
oxidation of phosphorus by the air, resulting, probably, from the de-
composition of unstable, highly oxidized bodies which are formed
during the action. Oxygen containing as much as 15 per cent of it
is produced by the interaction of fluorine and water (p. 241).
300
OZONE AND HYDROGEN PEROXIDE 301
Physical Properties of Ozone. — Ozone is a gas of blue color.
It boils at — 119°, so that when a mixture of oxygen and ozone is led
through a U-tube immersed in liquid oxygen (— 182.5°), the ozone is
liquefied. The opaque, deep-blue fluid contains only about 14 per
cent of oxygen, and this may be removed by evaporation.
Ozone is much more soluble in water than is oxygen. Its solubility
shows that, at 12°, 100 volumes of water would dissolve 50 volumes of
the gas at one atmosphere pressure. Its solubility, when mixed with
oxygen, is in proportion to its partial pressure (p. 155).
Chemical Properties of Ozone. — Ozone is relatively stable only
when mixed with much oxygen. Hence its density and molar weight
cannot be ascertained save by indirect means. The weight of a liter
of the mixture at 0° and 760 mm. having been measured, the ozone
may be removed by absorption in turpentine and the proportion of it
present in the gaseous mixture be thus ascertained. For example, if the
weight of 1 1. was 1.468 g. and 50 c.c. were absorbed by turpentine,
there were 950 c.c. of oxygen. The weight of this oxygen is 1000 :
950 : : 1.429 : x, from which x = 1.361 g. The rest of the weight,
1.468-1.361 or 0.107 g., was that of 50 c.c. of ozone. The weight of
1 1. of ozone at 0° and 760 mm. is therefore 2.140 g. The molecular
weight (weight of 22.4 1.) is thus 47.9 g., or nearly 48 g. The for-
mula of ozone is therefore 03.
When ozonized oxygen is heated, the ozone is decomposed at about
250-300°. The action for its formation :
302 <=± 203
is therefore reversible. That this equation, showing that three mole-
cules of oxygen give two molecules of ozone, is correct, may be demon-
strated by measuring the diminution in volume which accompanies the
action. If a shrinkage of 5 c.c. is observed in forming the ozone, it is
found that 10 c.c. more are then absorbed by turpentine. Thus the ozone
occupied 10 c.c., and the total oxygen from which it was made was
therefore 15 c.c. Hence three volumes of oxygen give two of ozone.
The formation of ozone absorbs much energy from the electric waves,
or, in other methods of making it, from the concomitant chemical
changes :
0 + 02 = 03 - 32,400 cal.
Ozone is a much more active oxidizing agent than oxygen. Mercury
and silver, which are not affected by the latter, are converted into
302 INORGANIC CHEMISTRY
oxides by the former. Silver gives the peroxide, Ag20.,. Paper dipped
in starch emulsion containing a little potassium iodide is used as a
test for ozone :
03 + 2KI + H20 -» 02 + 2KOH + I2.
The iodine gives a deep-blue color to the starch (cf. p. 235). This
test, however, will not distinguish ozone from chlorine or hydrogen
peroxide, and may, therefore, be used only in the absence of these sub-
stances. The last substance is always present in the air, and, since air
usually shows the above action, is probably responsible for the belief
that air contains ozone. The action on silver has never been obtained
with air. Ozone also removes the color from organic dyes, such as
indigo, by oxidizing them (cf. p. 269). Its activity as an oxidizing
agent, like the similar activity of hypochlorous acid, is due to the fact
that it contains much more energy than oxygen. In all its actions the
energy set free is greater by this excess than that liberated when oxy-
gen is used.
Oxygen and ozone are different substances (p. 35), that is, have
different properties. The difference in density, interpreted in terms
of the molecular hypothesis, gives us the statement of the nature of
the difference which is embodied in the formulae 02 and 08. The differ-
ence in activity, interpreted in terms of the conception of energy, gives
us the other method of stating the nature of the difference. The re-
cent preference for the second method is well illustrated by this case.
The first method uses a mere physical property, the second a fact
which is intimately connected with the whole chemical behavior of the
substance, a matter of much greater interest to the chemist.
Ozone may be distinguished from chlorine, nitrogen peroxide, and other oxi-
dizing agents, with the exception of hydrogen peroxide, by using pink litmus paper
instead of plain paper to carry the potassium iodide solution in the above test.
The potassium hydroxide set free by ozone turns the paper blue. Chlorine, for
example, gives an entirely different action : C12 + 2KI — > 2KC1 + I2.
Ozone is used commercially in bleaching oils and in purifying
starch. It is employed also for sterilizing drinking water in Lille and
other cities.
HYDROGEN PEROXIDE.
Hydrogen peroxide (H202) is found in minute amounts in rain and
snow. It is formed in small quantities, in a way not at present under-
stood, when moist metals rust.
OZONE AND HYDROGEN PEROXIDE • 303
Preparation of Hydrogen Peroxide. — When sodium peroxide
(<?.v.) is added, a little at a time, to a dilute acid, hydrogen peroxide is
set free :
Na202 + 2HC1 fc» 2NaCl + H20
2.
It may be separated from the salt (and a large part of the water)
by repeatedly shaking the mixture with ether (cf: p. 155). The rela-
tive solubility in water and ether is 1 : 0.0596, however, so that much
ether is needed. The ethereal layer, which rises to the top, when
evaporated, leaves a strong aqueous solution of the compound behind.
When hydrated barium peroxide (Ba02, 8H20) is shaken with cold,
dilute sulphuric acid a similar action takes place :
Ba02 + H2S04 fc? BaS04 j + H20
The excess of sulphuric acid may be removed by adding barium
hydroxide solution cautiously until no further precipitation of barium
sulphate occurs : Ba(OH)2 + H2S04 fc» BaSOJ + 2H20. Hydrochloric
acid or phosphoric acid may be used instead of sulphuric acid. The
second is largely employed in the commercial manufacture of hydrogen
peroxide. In each case, great care has to be taken to precipitate the
other products and all impurities from the solution. When hydro-
chloric acid is used, for example, the barium chloride produced by the
action is removed by adding silver sulphate :
BaCl2 + Ag2S04i=> BaSOJ + 2AgClJ.
An aqueous solution is also obtained by passing carbon dioxide
through barium peroxide suspended in water :
Ba02 + C02 + H20 <=> BaC03 J +H202.
Pure hydrogen peroxide is isolated from any of these solutions by
distillation under reduced pressure (p. 276). It is much less volatile than
water, but decomposes into water and oxygen violently at 100°. Hence
the lower pressure is required to make possible its volatilization at a
temperature below this point. At 68 mm. pressure, the water begins to
pass off first (at about 45°). The last portion of the liquid boils at 84-
85° and is almost all hydrogen peroxide.
By evaporating the commercial (3 per cent) solution at 70°, a liquid
containing 45 per cent of hydrogen peroxide may be made without
much loss of the material by volatilization.
304 INORGANIC CHEMISTRY
The Interaction of Barium Peroxide and Sulphuric Acid. —
It is -worth noting that, although common barium peroxide is not less
soluble in water than is the hydrated form, it dissolves much more
slowly. The fact that it is made by heating barium oxide in oxygen
and is composed of compact particles is accountable for this.
Every action upon a little-soluble, or slowly dissolving body, like
the barium peroxide in the above actions, is rather complex. It is
only the dissolved part of the substance that interacts. There is thus
a physical equilibrium between the undissolved and the dissolved
bodies, Ba02 (solid) +± BaO2 (diss'd), the displacement of which fur-
nishes the material for the chemical action. The latter has therefore to
follow the pace set by the former. When barium sulphate is precipi-
tated, another physical equilibrium follows the chemical change :
BaSO4 (diss'd) «=± BaSO4 (solid). When relatively insoluble bodies are
used or produced, there is thus a chain of equilibria each depending
on the others :
BaO2 (solid) <-; Ba02 (diss'd) + H2SO4 ±^ H2O2 + BaSO4 (diss'd) <^ BaS04 (solid).
If the barium sulphate ceased to be precipitated, its interaction in
solution with the hydrogen peroxide would drive the central action
backwards, and barium peroxide would be precipitated instead. The
success of the process thus depends on the fact that barium sulphate
is even less soluble than barium peroxide.
When carbon dioxide is used (see above), a similar chain of equili-
bria exists, and in that case it is the barium carbonate that is the less
soluble substance.
Other Modes of Formation. — Hydrogen peroxide is formed by
the direct union of hydrogen and oxygen. When a hydrogen flame is
allowed to play upon ice, appreciable amounts of the peroxide are
saved from being decomposed, as they ordinarily would be by the
heat of the action, and are found in the water.
It may be obtained by the action of acids upon the peroxides of
calcium, strontium, zinc, and copper.
Traces of hydrogen peroxide are formed when zinc, copper, lead, and other
metals are shaken with air and dilute sulphuric acid. It is produced when oxygen
is passed, in the neighborhood of the negative electrode, through the liquid in an
electrolytic cell containing dilute sulphuric acid. The gas is reduced by the hydro-
gen being liberated on the platinum plate.
OZONE AND HYDROGEN PEROXIDE 305
Physical Properties. — Hydrogen peroxide is a syrupy liquid of
sp. gr. 1.5. It blisters the skin, and, when diluted, has a disagreeable
metallic taste. It has been frozen (m.-p. — 2°).
Chemical Properties. — Hydrogen peroxide is very unstable, and
decomposes slowly even at — 20°. The dilute aqueous solution, when
free from impurities, keeps fairly well. The presence of a trace of
free acid increases its stability. Free alkalies and most salts assist
the decomposition ; hence the necessity for purifying the commercial
solution. Addition of powdered metals, of manganese dioxide, and
of charcoal causes effervescence even in dilute solutions, and oxygen
escapes :
2H202 -» 2H20 + 02.
The more concentrated solutions (38 per cent) remain quiescent in a
dish of polished platinum even at 60°, but the making of a slight
scratch on the bottom, beneath the surface of the liquid, causes pro-
fuse liberation of oxygen along the sharp edge thus produced. The
action of the catalytic agents is therefore probably mechanical.
Since the substance cannot be vaporized, even at low pressure,
without some decomposition, its molar weight has been determined by
the freezing-point method (p. 291). The freezing-point of a 3.3 per
cent solution in water was 2.03° below that of the water itself. Hence,
in 1000 g. of water, 3.3 g. would have given a depression of
2.03 x 96.7 -4- 1000, or 0.196°. Therefore a depression of 1.89° would
have been caused by 3.3 x 1.89 -5- 0.196, or 31.8 g., which is the re-
quired molar weight. Now the formula HO corresponds to a molar
weight of 17 and H202 to one of 34. It is evident, therefore, that the
latter is the correct formula.
Hydrogen peroxide, in solution in water, is a feeble acid. The
normal molar weight and very small electrical conductivity (see Chap,
xix) show that only a very small proportion of it can be ionized. As an
acid it enters into double decomposition readily. Thus, when it is
added to solutions of barium and strontium hydroxides, the hydrated
peroxides appear as crystalline precipitates :
Sr(OH)2 + H202 <=> 2H20 + Sr02.
The precipitation involves another equilibrium : Sr02 + 8H20 <=± Sr02,
8H20 (solid). The action of hydrogen peroxide upon chromic acid
(H2Cr04) is probably of a similar nature. The composition of the
product, which gives a beautiful blue solution, is not definitely known,
306 INORGANIC CHEMISTRY
as it decomposes almost immediately. Its formation, by adding a
drop of potassium dichromate to an acidulated solution of the peroxide,
is used as a test for the latter. The acid interacts with the dichro-
mate, giving the necessary chromic acid :
H2S04 + K2Cr207 + H20 +± 2H2Cr04 + K2SO4.
The blue substance has the property, unusual in inorganic compounds,
of dissolving much more readily in ether than in water. It is also
much less unstable when removed from the foreign materials in the
aqueous solution. Hence the test is rendered more delicate by ex-
tracting the solution with a small amount of ether. In the ethereal
layer the color of the compound is more permanent, as well as more
distinctly visible on account of the greater concentration.
Hydrogen peroxide is a much more active oxidizing agent than free
oxygen. It liberates iodine from hydrogen iodide, an action which,
in presence of starch emulsion (cf. p. 235), is used as a test for its
presence :
2HI + H202 -> 2H20 + I2.
It converts sulphides into sulphates. The white lead (q.v.) used in
paintings is changed by the hydrogen sulphide in the air of cities to
black lead sulphide, PbC03 + H2S -> PbS + H20 + C0a. This may
be oxidized to white lead sulphate by means of hydrogen peroxide :
PbS + 4H202 -> PbS04 + 4H2O,
and in this way the original tints of the picture may be practically re-
stored. Organic coloring matters are changed into colorless substances
by an action similar to that of hypochlorous acid (cf. p. 269). Hence
hydrogen peroxide is used for bleaching silk, feathers, hair, and ivory,
which would be destroyed by the more violent agent. The products
of its decomposition, being water and oxygen only, are harmless, and,
on this account, it is used as a bactericide in surgery.
Hydrogen peroxide exercises the functions of a reducing agent in
special cases, also. Thus, silver oxide is reduced by it to silver :
Ag20 + H202 «> 2Ag + H20 + 02.
A solution of potassium permanganate, in which the permanganic acid
has been set free by an acid, KMn04 + H2S04 <=± HMnO4 + KHSO4,
is rapidly reduced. The permanganic acid, with excess of sulphuric
acid, tends to undergo the first of the following changes, provided a
OZONE AND HYDROGEN PEROXIDE 307
substance is present which can take possession of the oxygen that
would remain as a balance :
50) (1)
_ (2)
2HMn04 + 2H2S04 + 5H202 -> 2MnS04 + 8^,0 + 502
2HMn04 + 2H2SO, -
(50) + 5H202 -
-> 2MnS04 + 3H,0 (+ 5(
-» 5H20 + 502
The first partial equation has been doubled to secure the even num-
ber of units of hydrogen required for the formula of water.
In all reductions by hydrogen peroxide, each molecule of the latter removes but
one atomic weight of oxygen. Whether it behaves thus because its two hydro-
gen units combine with this oxygen and all its own oxygen escapes, or because
it furnishes water and one oxygen unit of the pair required to form the mole-
cule of free oxygen (the substance reduced furnishing the other), has not been
determined.
The above action is used in quantitative analysis for estimating
the quantity of hydrogen peroxide in a given liquid after the liquid
has been acidified. The amount of a standard (p. 236) solution of the
permanganate which is required to decompose all the peroxide is
measured by means of a burette (q.v.*). The permanganate is deep
reddish-purple in color, while the products are colorless. Hence, after
the peroxide is exhausted, the next drop of the permanganate confers
a distinct, permanent, pink tinge upon the liquid. The addition of
the permanganate solution is stopped so soon as this condition is
reached and the volume of it that has been used is read off.
Thermochemistry of Hydrogen Peroxide* — The formation of
hydrogen peroxide from the free elements is accompanied by evolution
of heat :
H2 + 02 = H202 Aq + 45,300 cal.
Hence the substance is formed by direct union (p. 304). But its de-
composition into water and oxygen gives out a further supply of
heat:
H2 + O, = H20 + 0 + 23,100 cal.
The sum of these two stages, of course, yields the same result
(cf. p. 78) as the direct formation of water (68,400 cal.).
When hydrogen peroxide is used, instead of free oxygen, for oxidiz-
ing purposes, each such action liberates 23,100 calories of heat more
in the former case than it would in the latter. Hence the activity of
the substance as an oxidizer (cf. p. 271),
308 INORGANIC CHEMISTRY
Peroxides : Chemical Constitution and Molecular Structure.
-To represent the chemical behavior of hydrogen peroxide, two
different graphic formulae (cf. p. 279) have been proposed :
H-0 H
I and XO = 0.
H-O n/
In the latter of the two formulae, one of the oxygen units holds four
equivalents of other materials instead of two. Oxygen being or-
dinarily bivalent, the two extra valences may plausibly be supposed
to involve a feebler state of combination, and therefore to portray the
tendency of the compound easily to give up one unit of oxygen.
The linking of the oxygen units, common to both formulae, expresses
the fact that we do not obtain hydrogen peroxide from substances con-
taining less than two units of this element in each mole. Thus, to be
consistent, we write Na^, although we have no means of determin-
ing the molecular weight 6f this particular peroxide. The former of
the above formulae is more generally used.
All oxides containing two units of oxygen do not yield hydrogen
peroxide, however. Thus, lead dioxide (Pb0.2) gives, by interaction
with dilute acids, water and oxygen only. Carrying out our system,
therefore, we make barium and hydrogen peroxides alike, and we
assign different constitutions (p. 224) to barium and lead dioxides and
different structures to their molecules :
0 HC1 Cl H-O
Ba' | + -»Ba' + |
XO HC1 XC1 H-O
.0 Cl
Pbf +2HCl-*Pb/ -f H,0+0.
In confirmation of this we find that lead can form an unstable tetra-
chloride and other compounds in which, as in PbIV02, it is quadriva-
lent. But barium gives no other compounds in which there is even
the semblance of quadrivalence. So our structural formula ingeni-
ously leaves it bivalent even in Ba02. We assign, therefore, to all
substances which give hydrogen peroxide, and are therefore true
peroxides, the " peroxide structure," consisting of linked oxygen units,
while for those which give no hydrogen peroxide, we write formulae in
which the oxygen units are independent of one another.
OZONE AND HYDROGEN PEROXIDE 309
Exercises. — 1. What volume of ozone will be taken up by 100 c.c.
of water at 12° from a stream of oxygen containing 7.5 per cent of
ozone (p. 155) ?
2. Formulate the action of carbon dioxide on barium dioxide
(p. 303) after the manner of that of sulphuric acid on the same sub-
stance (p. 304). The dissolving gas gives an additional equilibrium :
CO2 (gas) + H20 <=> H2CO8 (diss'd).
3. At what temperature will a ten per cent solution of hydrogen
peroxide freeze (p. 305) ?
4. Write the thermochemical equations for oxidation of indigo
by ozone (pp. 271, 302) and by hydrogen peroxide.
CHAPTER XIX
ELECTROLYSIS
Introductory. — Experiment shows that most solutions which
exhibit chemical transformations by interchange of groups (p. 283), and
all which show evidence of dissociation by measurements of osmotic
pressure, freezing-point depression, and boiling-point elevation, are
precisely those which are conductors of electricity and suffer decom-
position by the passage of the electric current. Solutions which,
on the contrary, behave normally in respect to the three physical prop-
erties are nonconductors. Solutions of the former kind, consisting of
ionogens (p. 296) dissolved, usually, in water, are called electrolytes,
and the effect of an electric current upon them is named electrolysis.
Solutions of the same substances in toluene> chloroform, etc., do riot
conduct electricity, are not decomposed by the current, and, as we have
seen (p. 294), show no evidence of ionization. Substances like sugar,
methyl acetate, etc., which show no evidence of dissociation, are non-
conductors, whatever solvent we employ.
In endeavoring to connect these important facts, we must remem-
ber that the pure solvent by itself, and the pure substances which we
dissolve in the solvent by themselves are, at ordinary temperatures, all
but complete nonconductors of electricity. Dry salts, except when at
high temperatures and fused, and dry hydrogen chloride (p. 182), on
the one hand, and pure water (p. 95)\on the other, do not permit the
passage of the current. Yet a mixture of one of the former with the
latter conducts extremely well. Let us consider first the nature of
the decomposition which accompanies the conduction.
Chemical Changes Connected with Electrolysis. — When the
wires from a battery are attached to platinum plates immersed in any
electrolyte (e.g. Fig. 65, p. 169), we observe that the products appearing
at the two electrodes are always different. They may be of several
kinds physically, and will be secured for examination variously ac-
cording to their nature. When they are gases which are not too
soluble, they may be collected in inverted tubes filled with the solu-
310
ELECTROLYSIS 311
tion. Solids, if insoluble in the liquid, will either remain attached to
the electrode or fall to the bottom of the vessel as precipitates. Sol-
uble substances on the other hand will usually not be visible. They
may be handled by interposing a porous partition of some description
which will restrain the diffusion of the dissolved body away from the
neighborhood of the electrode, while not interfering appreciably with
the passage of the current. Surrounding one electrode with a porous
battery jar is a convenient method for effecting this.
When a current of electricity is passed in this fashion through a
solution of silver nitrate AgN03, we observe that at the negative
electrode metallic silver is set free and adKeres to the plate. At the
positive electrode a gas, oxygen, appears. Since these substances do
not account for all the constituents of the salt, we are impelled to ex-
amine the solution around each pole, and discover that nitric acid
(HN08) is being formed along with the oxygen at the positive end.
Although we have now found something more than the two parts of
the original molecule, we have little difficulty in explaining the pres-
ence of these two products on the assumption that the original
molecules were divided into the parts, Ag and N03. The latter, since
it is not a known compound, must have interacted with the water to
produce nitric acid and oxygen: 2NO3 + H20 — » 2HlSr03 + 0. The
atomic oxygen has subsequently united so as to form the gas (02).
The whole change may therefore be tabulated as follows :
Neg. Wire, Ag. < Ag.NO8 > O2 and HN08, Pos. Wire.
If we substitute cupric nitrate Cu(N03)2, we obtain a red deposit
of metallic copper on the negative plate, and at the positive plate
oxygen and nitric acid are formed. We infer therefore that the parts
of the original molecule are Cu and N08 :
Neg. Wire, Cu < Cu.(N03)2 > 02 and HNO8, Pos. Wire.
With a solution of potassium nitrate we find hydrogen and oxygen
appearing at the negative and positive electrodes respectively.
Litmus paper, however, shows the presence in the solution of a base
(potassium hydroxide, KOH) at the negative and an acid (nitric acid) at
the positive end. Secondary chemical changes have occurred at both
poles. We infer that the parts of the parent moleciiles are K and N03.
The former, in its customary manner (p. 99), instead of being
liberated, gave rise to free hydrogen and potassium hydroxide,
312 INORGANIC CHEMISTRY
2K + 2H20 -> 2KOH + H2 :
Neg. Wire, H2 and KOH< K.N03 >02and HN03, Pos. Wire.
We are confirmed in these conclusions when we employ a pool of mer-
cury in place of the negative wire. A portion of the potassium is
found to have dissolved in the mercury and escaped interaction with
the water.
When dilute sulphuric acid is electrolyzed (p. 95), the result is a
liberation of hydrogen and oxygen and an accumulation of sulphuric
acid round the positive electrode :
Neg. Wire, H2< H2.S04 > 02 and H2S04, Pos. Wire.
All acids give hydrogen alone at the negative electrode.
Of the various illustrations we have encountered, the electrolysis
of hydrochloric acid (p. 184) happens to be the only one which delivers
the two components (H and Cl) with the minimum of modification by
secondary interaction. The gases liberated are, of course, H2 and CLj.
The chlorides, bromides, and iodides of all such metals as do not
interact with water give equally simple results. Secondary actions are
equally worthy of notice, not only because they are common, but also
because they often play a part in industrial electrolytic processes.
On now comparing the chemical behavior of a large number of
ionogens with the results of electrolysis of the same substances, we are
led to identical conclusions. Acids contain hydrogen possessing
certain special properties (p. 281), and by electrolysis they divide so as
to give up this constituent alone at one electrode. Salts undergo
double decomposition easily and exchange radicals with other ionogens
(p. 283), and the current divides their molecules at the same point,
liberating the radicals.
Quite a crop of problems is raised by this discovery : Since a
solution may eventually be cleared of all the hydrochloric acid, for
example, which it contains, we should like to know how the con-
stituents in the center of the cell reach the electrodes. Then there is
the question of the quantity of electricity required to effect a given
amount of decomposition in a given ionogen. Finally, since the
details are, as usual, inscrutable, a formulative hypothesis will be
needed to explain the whole proceeding.
Ionic Migration. — The first of these questions is easily answered
by experiment. We have only to take an ionogen one of whose radi-
ELECTROLYSIS
313
cals is colored, and watch the movement of the colored material as it
drifts towards the electrode. Thus, in dilute cupric sulphate solution,
a freezing-point determination shows that the depression has practi-
cally double the normal value. In other words, the dissociation
CuS04 <=± Cu" + S04"
is almost complete.* Now, the blue color of this solution cannot be
due to the remaining molecules of CuS04, for anhydrous cupric sul-
phate is colorless. Nor is it due to the color of the S04" ion, for
dilute potassium sulphate and dilute sulphuric acid are both color-
less. On the other hand, all cupric salts, in dilute solution, have the
same tint. The color is therefore that of the cupric ion (Cu**). Simi-
larly the deep-yellow tint of potassium dichromate K2Cr2O7 (q.v.) is
that of the Cr207" ion. Hence either of these substances will serve
the purpose of showing how the ionic material moves.
One of the above salts is dissolved in warm water containing about
5 per cent of agar-agar, and the lower part of the U-tube (Fig. 77) is
charged with the mixture. After
this fluid has set to a jelly, a few
grains of powdered charcoal are
added on each side to mark the
present limits of the colored ions,
or strips of paper are pasted on
the outside. Then any colorless
electrolyte, such as potassium
nitrate solution, is added on each
side, and the electrodes are hung
one in each limb. The lower half
of the potassium nitrate solution
on each side contains agar-agar
also. The agar-agar does not offer
any appreciable resistance to the
motion of the ions, and is pre-
sumed to form a sort of open network in the solution. It is added to
prevent all motion of the water. Immersion of the whole tube in ice
and water prevents the melting of the jelly by the heat generated
by the current.
* In all ions the valence is indicated, as in this equation, by the number of
superior marks • or ' as the case may be. The use of this custom will appear
presently.
FIQ. 77.
314 INORGANIC CHEMISTKY
After a time, we observe that the blue cupric ions ascend above the
mark on the negative and descend away from it on the positive side.
With potassium dichromate the yellow ions move in the opposite
direction with reference to the poles. In each case there is no shad-
ing off in the tint. The motion of the whole aggregate of colored ions
occurs in such a way that, if the contents of the tube were not held in
place by the jelly, we should believe that a gradual motion of the
whole solution was being observed. With a current of 110 volts, and
a 16-candle power lamp in series with the cell, the effect becomes
apparent in a few minutes.
Although the SO/' ions are invisible, we may safely infer that they
are drifting towards the positive electrode. Indeed, this can be demon-
strated by interposing a shallow layer of jelly containing some barium
salt a little distance above the charcoal layer on the positive side.
When the SO/' ions reach this, barium sulphate begins to be precipi-
tated and the layer becomes cloudy. In a similar way the progress of
other colorless ions may be rendered visible.
It appears therefore that electrolysis is not a local phenomenon,
going on round the electrodes only, but that the whole of the dis-
sociated solute is set in motion. It is on account of this remarkable
property of traveling or migrating towards one or other of the elec-
trodes connected with a battery that the ions receive their name (Gk.
itav, going). The term was first applied by Faraday to the materials
liberated round the electrodes.
Relative Speed of Migration of Different Ions. — The speeds of
different ions may readily be compared. The cupric ion moves at the
same speed whatever salt of copper we employ. In fact, the speeds
of all ions are individual properties and are independent of the nature
of other ions that may be present. The speeds of all are increased
by using a current of greater electromotive force. Under similar con-
ditions, the relative speeds of most ions are in the neighborhood of
50 or 60, on the scale commonly used in expressing ionic velocities.
Thus, we have, K* 65.3, Cl' 65.9, Cu" 49. The speed of the hydrogen
ion is the greatest of all, 318, while that of hydroxyl ion (OH') comes
next, being 174.
The actual speeds of these ions in dilute sohitions at 18°, when
driven by a potential difference of 1 volt between plates 1 cm.
apart, expressed in cm. per hour is : K* 2.05, Cl' 2.12, Cu" 1.6, H* 10.8,
OH' 5.6.
ELECTROLYSIS
315
By an experiment similar to the last, and devised by A. A. Noyes, the relative
speeds of different ions may be demonstrated. The U-tube (Fig. 78, showing the
same tube A before the current starts, and B after it has been passing for some
time) is partly filled with agar-agar emulsion containing potassium chloride and
phenoiphthalein (see Indicators). On the right side, a few drops of potassium
hydroxide have been added to render the mixture pink. On the left, a few drops
of hydrochloric acid are present, and the mixture is colorless. Above the charcoal
layer, in the right limb, a mixture of hydrochloric acid and cupric chloride (i.e.,
H" and Cu"), and in the left limb potassium hydroxide solution (i.e. OH'), are
placed. The positive electrode is introduced on the right and the negative on the
left. The H" and Cu" ions drift away from the former down the tube towards the
H' & Cu" OH' g
1 i BI
Pink — OH
FIG. 78.
latter, the OH' ions away from the latter down the tube towards the former.
The motion of the H" is marked by the disappearance of the pink color, that of the
Cu" by the advance of a blue layer, that of the OH' by the progress of a pink
coloration.* By the time the H" ions have been displaced 5J cm., the Cu" ions
have moved 1 cm. and the OH' about 2£ cm. These distances indicate their rela-
tive speeds of migration.
Faraday's Laws. — Having learned that electrolytes furnish at
least two definite decomposition products by the action of the elec-
tricity, we naturally inquire next whether there is any chemical relation
between the quantities of the products set free by the same current.
Quantitative experiments in electrolysis show the most perfect adjust-
ment in this respect. Thus, in a single cell, the quantities of material
* Bases, on account of the OH' they give, turn phenoiphthalein solutions from
colorless to pink ; acids, on account of the H" they furnish, turn it from pink to
colorless (see Indicators).
316
INORGANIC CHEMISTRY
liberated at the two poles are invariably chemical equivalents of one
another. With hydrochloric acid, while 1.008 g. of hydrogen is being
liberated at one pole, 35.45 g. of chlorine are set free in the same time
at the other. While 63.6 g. of copper, are being deposited from cupric
sulphate at one pole, 96 g. of SO4 are being liberated at the other,
and, by interaction with the water, form 16 g. of oxygen and 98 g. of
sulphuric acid.
Again, the amount of any one substance liberated is proportional
to the quantity of electricity -which has traversed the cell. This is the
first part of Faraday's law.
Finally, the passage, of equal quantities of electricity through sev-
eral different acids liberates equal amounts of hydrogen from each.
This is true, whether the passage of the given quantity of electricity
is compressed into a brief time in one case and spread over a longer
time in another, or is uniform in all cases compared. It is irrespec-
tive of the state of dilution and of the temperature of each acid.
Thus two moles of hydrochloric acid are always decomposed for
every one of sulphuric acid by the same current. Similarly, if in dif-
ferent cells we place solutions of substances like sodium chloride (NaCl),
cupric chloride (CuCl2), antimony chloride (SbCl3), ferrous chloride
(FeCl2), and ferric chloride (FeCl3), equal amounts of chlorine are
liberated by currents of equal strength in the same time in each.
If we consider the relation of these facts to the equivalence of the
materials liberated in any one cell, it will be evident that when one
gram of hydrogen is liberated from each of the two acids mentioned
above, one equivalent
of chlorine will be set
free in the one cell, and
one equivalent of SO4,
or half the weight rep-
resented by the for-
mula, will be set free
in the other. Similarly,
with the chlorides of
the first three metals, while 35.45 g. of chlorine are being liberated in
each cell, the quantities of the metal set free will be, of sodium 23 g.,
of copper one-half of 63.6 g., and of antimony one-third of 122 g.
Finally, with the two iron salts, the quantities of iron liberated by the
same current will be one-half and one-third of 56 grams respectively.
The simplest way in which to insure the passage of precisely equal
KaCl CuCl
Cl H iS04Na Cl iCu Cl JSb Cl jFe Cl JFe
•ll'l'.
FIG. 79.
ELECTROLYSIS 317
amounts of electricity through all the cells is to arrange them in
series. We know, that in such circumstances the quantity of electri-
city traversing any section of the whole circuit must be the same as that
traversing any other. In a series of cells containing substances like
the above, therefore, during the time that 1.008 g. of hydrogen is being
set free, we shall have liberation of the equivalent quantities of each
of the other ions (Fig. 79). Thus the second part of Faraday's law
states that chemically equivalent quantities of ions are liberated by
the passage of equal quantities of electricity.
The Ionic Hypothesis. — The main facts in regard to electrolytic
conduction being now before us, the problem is to adapt the atomic and
molecular hypotheses to their explanation, that is, to their detailed de-
scription. There are two peculiarities to be accounted for :
How can the production of a conducting medium by mixing two
nonconductors be imagined to take place ? The solvent is a non-
conductor, and the ions, even if they are composed of conducting mate-
rial, are distributed through the liquid as independent particles and
cannot furnish a continuous medium for the stream of electricity.
This will be clear when we remember that although liquid mercury is
an excellent conductor, mercury vapor, composed as it is of conduct-
ing particles, is not a conductor.
Again, the conducting power of the solution is indissolubly con-
nected with the fact that the original molecules of the solute have been
broken up by the solvent into smaller molecules containing one or
more atoms. Why should this particular kind of sub-molecules be at-
tracted by electrically charged plates, which have been lowered into
the solution, when molecules of dissolved sugar, for example, are not
so attracted?
An answer to the second question readily suggests itself. The
only bodies which we find to be conspicuously attracted by electrically
charged objects are bodies which are already provided with electric
charges of their own. Thus we are led to add to the molecular hypoth-
esis the assumption that substances which undergo dissociation in
solution divide themselves into a special kind of electrically charged
molecules. Since the solution, as a whole, has itself no charge, equal
quantities of positive and negative electricity must be produced :
HC1 <=» H + Cl NaCl <± Na + 01 NaOH <=> Na + OH.
Wild as this supposition seems at first sight to be, it turns out that
O
O
o
e
e
e © © e
e © e Q ©
318 INORGANIC CHEMISTRY
no valid objection to it can be raised. That it furnishes an answer to
both of our questions must first be shown.
A battery is a machine which maintains two points, its poles, or
two wires connected with them, at a constant difference of potential.
One cell of a storage battery, for example, maintains a potential dif-
ference of 'two volts. When the wires are joined, directly or indi-
rectly, the poles are immediately discharged, but the cell continuously
reproduces the difference in potential by generating fresh electricity.
Now the effect of immersing two plates, one of which is kept by the
battery at a definite positive potential and the other at a definite nega-
tive potential, into a liquid filled with floating multitudes of minute
bodies, already highly charged, may easily be foreseen.
The figure (Fig. 80) will
Cathode + Anode v
cation = Ag convey some idea of the be-
anion — NO3 — > ~r •>
havior of the parts of a system
such as we have imagined.
The electrodes are marked —
and +. The negatively
charged plate attracts all the
positively charged particles in
""I'l'l'f the vessel, and, although these
FIG. go. particles are in continuous and
irregular motion, they never-
theless begin, on the whole, to drift toward the plate in question. On the
other hand, the negatively charged particles are repelled by this plate
and attracted by the positive plate, so that they drift in the bpposite
direction. Those which are nearest each plate, on coming in contact with
it, will lose their charges of electricity, turning thereby into the ordinary
free forms of the matter of which they are composed. The continuous
removal of the electrical charges of the plates through contact with ions
of the opposite charge furnishes occasion for recharging of the plate from
the battery, and thus gives rise to a continuous current in each wire.
Again, the continuous drifting of positively and negatively charged
particles in opposite directions through the liquid, constitutes what, in
the view of all external means of observation, appears to be an electri-
cal current also. A magnetized needle, for example, which is deflected
when brought near one of the wires of the battery, is influenced in the
same way by being brought over the liquid between the electrodes.
The illusion, so to speak, of an electric current is complete, although
in reality it is a convection of electricity that is taking place. Further-
ELECTROLYSIS 319
more, the quantity of electricity being transported across any section
of the whole system is the same as that across any other, whether
this section be taken through one of the wires, through the electrolyte,
or even through the battery at any point. As fast as the ions are thus
annihilated as such, the undissociated molecules are drawn upon for
the production of fresh ones, as in all chemical equilibria. Eventually,
by continuing the process long enough, if the substances set free are
actually deposited and do not go into solution again in any form, the
liquid can be entirely deprived of the solute which it contains.
The analogy to the transportation of a fluid like water is noticeable,
although not complete. Water may be transported in three ways. It
may flow through a pipe, it may pass by pouring freely from one con-
tainer to another, and it may be carried in vessels. ' Thus a stream of
water, essentially continuous, might be arranged, in which part of the
passage took place through the pipes, part by pouring from the pipes
into buckets, and part by the carrying of those buckets between the ends
of the pipes. The quantity of water passing a given point per minute
in this system would be the same at every part, although the actual
method by which the water was transported past the various points
might be different. In such a disjointed circuit we suppose the elec-
tricity to move when carried from a battery through an electrolytic
cell. It flows in the wire, passes by discharge between the pole and
the ion, and is transported upon the ions in the liquid. The parallel
is imperfect, however, because we have used the conception of two
electric fluids and because the ions are already charged in the solution,
and before any connection with the battery is made. They do not, so to
speak, transport the electricity of the battery, but their own.
Difficulties Presented by this Hypothesis. — The question was
raised (p. 295), as to how we can imagine separate atoms of sodium to
exist in water without acting upon it, as the metal sodium usually does.
But the ions of sodium in sodium chloride solution are not metallic
sodium. They bear large charges of electricity. They possess an
entirely different, and in fact, by measurement, much smaller amount
of chemical energy than free sodium. And, as we have seen, the
properties of a substance are determined as much by the energy it con-
tains as by the kind of matter. Metallic sodium and ionic sodium are,
simply, different substances.
We think of hydrogen chloride and common salt as exceedinglv
stable substances, and are averse to believing that precisely these c-om-
320 INORGANIC CHEMISTRY
pounds should be highly dissociated by mere solution in water. But it
must be remembered that in solution they undergo chemical change very
easily, and it is only in the dry form that they show unusual stability.
Again, why do not the ions combine in response to the attractions
of their charges ? The answer is that they do combine, but the rate
at which combination takes place is no greater than that at which the
molecules decompose, so that on the whole the proportion of ions to
molecules remains unchanged.
Finally, it might appear that the assumption that bodies could
retain high charges in the midst of water is contrary to all experience.
It must be remembered, however, that the molecular, pure water,
which separates the ions from one another, is a perfect nonconductor.
The moisture which covers electrical apparatus and causes leakage of
static electricity is not pure water, but a dilute solution containing
carbonic acid (p. 114) and materials from the glass of which the appa-
ratus is made. It conducts away the charge electrolytically, by means
of the ions it contains, and not by itself acting as a conductor.
Amounts of Electricity on the Ions. — Since one atomic weight
of chlorine (35.45 parts) and one atomic weight of hydrogen (1.008
parts) are simultaneously liberated in the electrolysis of hydrochloric
acid and the solution remains electrically neutral, it follows that the
ions of hydrogen and chlorine bear equal charges. Again, since the
same thing is true of one atomic weight of copper (63.6 parts) and two
atomic weights of chlorine (70.9 parts) in cupric chloride, it follows
that the cupric ion bears a double charge. Similarly, the antimony
(p. 316) ions bear triple charges. In ferrous and ferric salts the
amounts carried by the iron ions differ, being in the one case twice,
and in the other three times as. great as those carried by the hydrogen
or chlorine ions. We may sum this up, then, by saying that univalent
ions possess the same quantity of electricity, and other ions bear
quantities greater than this, in proportion to their valence. This rule
is simply a restatement of Faraday's law, which it brings into direct
relation to the ionic hypothesis.
In writing equations involving ions, the numbers of + and —
charges must always be equal :
+ — +++ —
KC1 <=> K + Cl SbClg <=i Sb + 3C1
CuCl2 «z» Cu + 2C1 FeCL, <± Fe + 2C1
» Cu + SO4 FeCls ^ Fe++ 3C1
ELECTROLYSIS 321
In order that this idea may be carried out consistently, the libera-
tion of any of these ionic materials at one electrode in electrolysis is
written as follows :
Here © and 0 represent the unit quantities of negative and posi-
tive electricity furnished by the battery to the electrodes and destroyed
by opposite charges upon the ions~.
The harmony between the quantity of electricity and the chemi-
cal valence of the material which it liberates is complete. The
picture which the process of electrolysis in a series of cells (p. 316)
presents to our minds is very interesting. The progress of the elec-
tricity through the series is accompanied by a simultaneous discharge
in all the cells of chemically corresponding numbers of atoms. For
eve'ry atom of antimony that is liberated in one cell, three atoms of
chlorine, three atoms of hydrogen, and one atom of ferric iron, are set
free at the same time. For two atoms of ferric iron, three atoms of
ferrous iron and three atoms of copper are deposited. Even in the
battery which generates the current, the chemical changes taking place
proceed atom for atom and valence for valence in unison with those in
the cells on the circuit. For example, if the battery contains zinc plates,
for every atom of zinc that dissolves, one of copper and two of chlorine
will be liberated in one of the cells. Our imaginary mechanism thus
puts all the processes going on in the circuit in the light of move-
ments of the parts of a perfectly adjusted and interlocked machine.
Resume and Nomenclature. — An ion may be denned as, an
atom or group of atoms bearing a positive or negative charge of
electricity, and formed through the dissociation of an ionogen by a
solvent like water.
Each molecule gives two kinds of ions with opposite charges.
These two are forthwith distinct and independent substances, save that
the attractions of the charges prevent separation by diffusion. They
differ from non-ionic substances of the same material composition when
such are known. The electrical charge is one of the essential con-
stituents, and when it is removed the properties alter entirely. Thus we
have two kinds of hydrogen, gaseous molecular hydrogen (H2), and ionic
+
hydrogen (H), with entirely different chemical properties (p. 295).
Since the writing of the -f- and — charges over the symbols
322
INORGANIC CHEMISTRY
occupies much space, we shall hereafter employ a dot for the former
and a little dash for the latter: IT, Cl', Cu", SO/', Fe", Fe*", NH\.
In the ions formed from one molecule, the number of dots and dashes
must be equal.
Since ionic hydrogen, ionic chlorine, etc., are entirely different in
physical and chemical properties from the corresponding free ele-
ments, they should receive separate names. When it is inconvenient
to say "ionic hydrogen," "nitrate ions" (NOS'), etc., the following,
based on Walker's system, will be used :
SYM-
BOL.
N AMK.
ANION OF
SYM-
BOL.
NAME.
CATION OF
SALTS OF
SO/'
SO/'
CIO/
CIO,'
Sulphanion
Sulphosion
Perchloranion
Chloranion
Sulphates
Sulphites
Perchlorates
Chlorates
Na'
Ca"
Cu"
K*
Natrion
Calcion
Dicuprion
Kalion
Sodium
Calcium
Cupric copper
Potassium •
el-
Chloridion
Chlorides
Fe"*
Triferrion
Ferric iron
s''
NO/
Sulphidion
Nitrauion
Sulphides
Nitrates
NH/
Fe-
Amrnonion
Diferrion
Ammonium
Ferrous iron
OH'
Hydroxidioii
Hydroxides (bases)
ll '
Hydrion
Hydrogen(acids)
Faraday distinguished the two kinds of material which proceed
with and against the positive current by name. His terminology is
still used. Ions which proceed in the same direction as the positive
current are called cations (G-k. Kara, down). Such are H*, Cu", K",
NH/. They are metallic elements, or groups which play the part of
a metal. The electrode (G-k. oSds, a path) upon which they are
deposited, the negative electrode, is spoken of as the cathode
(Gk. 17 KafloSos, the way down).
The particles which move in the direction of the negative current,
and against that of the positive, are named anions (Gk. dva, up). The
ions Cl', N08', SO/', MnO/ are of this kind. They are usually com-
posed of non-metals, although sometimes, as in MnO/, the components
may be partially metallic. They are set free at the positive electrode,
which is therefore named the anode (Gk. 17 dvoSos, the way up).
Chemists speak of metals and non-metals as positive and negative
elements, respectively (cf. p. 119), even when electrical relations are
not directly in question, and ions are not concerned.
Actual Quantities of Electricity Concerned in Electrolysis.—
The coulomb is the unit quantity of electricity : the liberation of one
ELECTROLYSIS 323
gram of hydrogen corresponds to the passage of 96,540 coulombs
round the circuit. In other words, one gram of hydrion (H*) carries
this quantity of electricity. Equivalent amounts of other ions, for
example, 108 g. of argention (Ag') and 9^- g. of sulphanion (SO/'),
carry the same quantity. Thus the unit quantity of electricity, one
coulomb, deposits, or is carried by ^£?TT g. of hydrogen or ^§f ^ g.
of silver.
In consequence of this, the deposition of silver or copper is used
as a means of measuring quantities of electricity. The increase in
weight of the negative electrode in a cell, called under such circum-
stances a voltameter, is a measure of the quantity of electricity which
passes around the whole circuit of which it forms a part.
It is more common to define the quantity of electricity in terms of
current strength. A current of such nature that one coulomb flows
through the system per second is said to have a strength of one
ampere. Such a current liberates ^£4^ g. of hydrogen or ^s£f tf S- °f
silver per second. A current which deposits twice as much, has two
amperes current strength.
From this it is evident that a current of one ampere would take
96,540 seconds, or twenty-six hours and 49 minutes, to deposit 1 g. of
hydrogen (about 11 liters), or 108 g. of silver, or 48 g. of sulphanion.
A current of five amperes would accomplish the same result in a fifth
of the time.
The Electrical Energy Required to Decompose Different
Compounds. — Chemical compounds are of very different degrees of
stability, and hence the quantities of energy, electrical or otherwise,
required to decompose them vary widely. Thus, hydrogen chloride is
very stable, while hydrogen iodide is easily decomposed by heating.
The disunion of equivalent quantities of these substances in aqueous
solution absorbs 39,300 cal. and 13,100 cal. of heat energy, respectively.
Hence, although equal quantities of electricity (96,540 coulombs in
each case) perform this office, very unequal amounts of electrical energy
are used up in the electrolysis.
The energy in a stream of water is represented by the product of
the quantity passing a given section and the pressure or head of water
available at that point. If the pressure is low, the work that can be
done will be small, even if the quantity flowing is great. So electrical
energy is expressed by the product of the current strength, or quantity
passing per second during a certain period of time, and the electro-
324 INORGANIC CHEMISTRY
motive force. The latter corresponds to pressure, and is denned by
the difference in potential of two points in the circuit between which
the energy is being used up.
Now, in the series of cells which was described (p. 316), each cell,
while being traversed by the same quantity of electricity as any of
the others, cuts down the electromotive force of the current in propor-
tion to the amount of energy consumed by the decomposition going on
within it. Hence, while a voltmeter will show no difference in poten-
tial between two neighboring parts of the heavy wires used as connec-
tions, for no work is being done in the wires, it will show a considerable
difference in potential between two points which are separated by one
of the cells.
A system of cars hauled by a cable is analogous to our set of cells
and more familiar. When clutched to the cable, all the cars move
with equal speed, but, being loaded with different numbers of passen-
gers, take very different amounts of power from the moving cable.
We should infer from this, that to decompose every electrolyte, a
current of a certain minimum electromotive force, sufficient to furnish
the fall in potential necessitated by the chemical change, which would
be different in different cases, would be required. This is found to be
the case. For the easy decomposition of sulphuric acid and liberation
of the products an electromotive force of at least 1.92 volts is necessary,
for hydrochloric acid 1.41 volts, for hydriodic acid 0.53 volts, for zinc
sulphate 2.7 volts, and for silver nitrate 0.70 volts. When we use a
current of electromotive force falling short of that specified, we find
that the flow of electricity is interrupted. The electrolytic cell practi-
cally forms a break in the circuit, (see Chap, xxxviii).
Polarization. — It is found that when plates of platinum, a metal
which is not acted upon by the liberated radicals, are used, the products
of electrolysis accumulate on the electrodes and tend to produce a re-
verse current (see Electromotive chemistry). The cell is said to be
polarized. Thus, after hydrochloric acid has been electrolyzed for a
few moments, hydrogen and chlorine adhering to the two platinum
plates set up this current.* If the battery is disconnected, the elec-
trolytic cell becomes for a brief time itself a battery, the re-ionization
of the hydrogen and chlorine (reproducing hydrochloric acid) furnish-
ing the energy. It is the continuous overcoming of this reverse current,
* If copper plates are used, cupric chloride is formed at the positive plate
(anode), and no polarization can occur at that plate.
ELECTROLYSIS 325
and prevention of the reunion, that demands the minimum electro-
motive force (here 1.41 volts) of which mention has just been made.
It is possible to arrange cells in which no polarization can take
place. Thus, when we electrolyze cupric sulphate between copper
electrodes, the copper is deposited upon one plate and the S04 removes
the copper from the other plate, forming cupric sulphate, thus restoring
the electrolyte to its original condition. The only difference is that a
portion of the copper has been deposited on one pole and an equivalent
amount has been removed from the other (see Copper refining). With
such cells, no minimum difference in potential is required to effect elec-
trolysis, for there is no polarization current to be overcome. The
feeblest electric current will produce continuous, if slow, chemical
change.
This result is extremely interesting, for it shows that the operation
of electrolysis in itself does not require the consumption of much
energy. If the molecules were actually torn apart by the electricity,
then all electrolytic operations would require a minimum electromotive
force for their maintenance. The fact just stated, therefore, is signifi-
cant, for it confirms the present views in regard to the theory of solu-
tions. It is in agreement with the belief that the actual production of
the ions is accomplished by the water in advance, and quite independ-
ently of the use of electricity, and that the sole function of the elec-
tricity in the process of electrolysis within the solution consists in the
pilotage of the ions in reverse directions according to their charges, an
operation which necessarily consumes but little energy. The friction
alone of the moving ions has to be overcome. It makes clear the fact
that it is only when the chemical change in the cell involves the actual
decomposition of some material, accompanied (as in the electrolysis of
hydrochloric acid) by the final delivery of the constituents in the free
state, that considerable consumption of electrical energy, proportional
to the extent of the chemical change, must take place.
Conductivity for Electricity. — The facility with which equi-
molar solutions of different substances conduct electricity, when they
are placed under like conditions, depends jointly on the degree of
ionization, on the speed with which the ions move, and on the valence
of the ions. When equivalent instead of equimolar amounts are com-
pared, the last of these factors drops out of consideration. The most
highly dissociated acids, as we should expect, since they give large
numbers of the speedy hydrogen ions, are the best conductors, The
326 INORGANIC CHEMISTRY
highly ionized bases, such as potassium and sodium hydroxide, come
next. The best conductors among salts fall considerably behind both
of these, because, although their degrees of ionization may not be less
than those of the best conducting acids and bases, their ions all move
more slowly than do hydrion and hydroxidion. On the other hand,
concentrated solutions all conduct badly, relatively to the number of
molecules originally used in making them, because only that propor-
tion of the substance which is ionized contributes to the conduction
(p. 318). All this is just what we should expect, in view of our
hypothesis, for the passage of the electricity must be dependent upon
the frequency with which discharges of the ions upon the electrodes
occur, and this, in turn, must depend upon both the concentration and
the speed of the ions. To return to an analogy used once before, the
rate at which a fluid can be transferred between two reservoirs must
depend upon the denseness of the array of buckets available and on
the speed with which they are moved and on their individual capacity.
Ordinarily, it is the resistance which a substance presents to the
passage of the electric current which is measured. Obviously, how-
ever, for the present purpose it is more convenient to give expression
to the reciprocal of this value, which we term the conductivity. In
order that the results may have chemical significance, we express them
in terms of the conducting power of one gram-equivalent of the com-
pound dissolved in water and placed in a narrow cell whose opposite
walls, of great area and situated one centimeter apart, form the
electrodes. Since the water is a nonconductor, the conducting power
of the solution intervening between the plates is a measure of the
capacity of the dissolved substance for facilitating the discharge
between the poles. Inasmuch as varying the amount of the solvent
will not affect the velocity of the ions of the substance, any alteration
in conducting power resulting from dilution must depend solely on
the change in the number of ions available for carrying off the elec-
tricity. The conductivities of solutions of the same substance in
different concentrations must therefore be proportional to the degrees
of its ionization.
A trough and amperemeter * (Fig. 81) may be used to illustrate
this principle. The electrodes are here long. strips of copper foil,
which pass down at the ends of the trough and are situated, not one
centimeter, but ten or fifteen centimeters apart, in order that the con-
* For these experiments an amperemeter of low resistance, O.'i-l ohm, must be
used, and a battery of one or two accumulator cells is sufficient.
ELECTROLYSIS
327
tents of the vessel may be more easily seeu. When the two instru-
ments are placed in circuit with a battery and very pure water is
poured into the cell, the amperemeter does not indicate the passage
of any current of electricity. When a shallow layer of concentrated
hydrochloric acid is substituted, the situation is that a definite amount
of hydrogen chloride dissolved in a small amount of water forms one
of the links in the electric circuit. The deflection of the needle in the
amperemeter indicates that a certain current of electricity is able to
pass through this acid. When now distilled water is gradually mixed
with the acid, the amount of conducting material, in this case hydrogen
FIG. si.
chloride, is not altered. If, therefore, its capacity for conducting
were not affected by the dilution, the deflection of the needle in the
amperemeter would not change. The wider dissemination of the hydro-
chloric acid, tending to diminish the specific conductivity of the solu-
tion, should he exactly compensated by the greater area of the
electrodes coming into service. What is actually observed, however,
is a very marked improvement in the conducting power of the solu-
tion.* At first the reading of the amperemeter increases very rapidly.
* The interpretation of the experiment becomes easier if the trough is first
nearly filled with distilled water and the absence of deflection noted. Then a layer
of concentrated hydrochloric acid is introduced below the water, by means of a
long-stemmed dropping funnel, and the deflection is read. Finally, the layers are
destroyed by stirring, and a great increase in the reading of the amperemeter ob-
served. Here ajl the hydrogen chloride and water are between the electrodes ar
the time of the second reading, and the greater value of the third reading cannot
be attributed to the use of a larger area of the poles, but solely to the redistribution
of the acid throughout a greater volume.
328
INORGANIC CHEMISTRY
Later, the effect of additional dilution becomes less marked, until
finally it becomes very slight indeed.
When a saturated solution of cupric chloride is substituted, dilu-
tion is accompanied by a similar improvement in conductivity. Here
we notice, besides, that the yellowish-green liquid with which we start
changes to a pale blue, as the molecules of cupric chloride are dis-
sociated and the color of the solution becomes more exclusively that
of the copper ions. When the solution has become perfectly blue,
further dilution is seen to affect the conductivity but slightly.
The approach to a maximum of conductivity reached in these two
cases indicates that practically the whole of the material has assumed
the ionic form. Theoretically the absolute maximum would be reached
at infinite dilution. The conductivity of the same amount of sub-
stance in more limited dilution is that of the proportion of ions corre-
sponding to this dilution, since the complete molecules, still present,
are without influence on conductivity. Thus the ratio of the conduc-
tivity at a given dilution to the maximum conductivity is equal to the
proportion of the whole material ionized at the given dilution. From
a series of measurements for a fixed amount of a substance at different
dilutions, after the results have been plotted, we can usually (see,
however, below) ascertain the limiting, maximum conductivity by
graphic extrapolation. If A,, is the conductivity of an equivalent of
the substance dissolved in v liters of water, and A^ the conductivity
of the same amount at infinite dilution, then A^/A^ is the proportion of
molecules completely ionized in the former solution. Xv is called
the equivalent conductivity at the dilution v.
The following numbers show the equivalent conductivities at 18°
of solutions of four different substances, expressed in the iinits always
employed for the purpose (which are reciprocal ohms). The symbols
AOJ, meaning 1 equivalent in 0.1 1. ; At, meaning 1 equivalent in 1 1. ;
and so forth, denote the concentrations.
\M
\
\o
\oo
XM
(Calc.)
Hydrochloric acid . .
Sodium chloride . . .
Sodium acetate . . .
Acetic acid .
64.4
0.05
301.0
74.4
41.2
1.32
351.0
92.5
61.1
4.6
370.0
103.0
70.2
14.3
384
110
78
(352)
It will be seen from inspection of these figures that the conductiv-
ity does not improve much in the case of the first three substances
ELECTROLYSIS 329
when a solution containing one equivalent in 10 1. is diluted ten times,
and that further dilutions, no matter how extensive, produce a still
smaller effect. On the other hand, acetic acid conducts very badly in
concentrated solution, and, while the conductivity improves with dilu-
tion, it is not possible experimentally to observe any approach to the
maximum conductivity. The conductivity, in cases like this, is still
far removed from the maximum at dilutions at which, with other sub-
stances, the maximum is nearly attained.
In cases like that of acetic acid, the conductivity at infinite dilu-
tion cannot be estimated by extrapolation. But fortunately another
method is available. The values 384, 110, and 78 for X in the cases
«
of the first three substances can be reached by extrapolation, and rep-
resent the conducting powers of equal numbers of ions, for there are
equal numbers of equivalents present and no molecules remain un ion-
ized. These values are unequal solely because of the differing speeds
of the ions concerned. Each of them derives its value from numbers
representing the relative speeds of the two ions present, and must be
the sum of these two numbers. If, therefore, we measure the relative
speeds of the two ions (p. 315), we can divide the value of X^ in this
proportion and learn the part which each ion contributes to the total.
Dividing 384 in this way we get the'speed of H* = 318 and of Cl' =
65.9 already given (p. 314). Dividing X^ for sodium acetate (78), simi-
larly, we get the speeds ISTa'= 44.4 and C2H302' = 33.7. The speeds of
Cl' and Na* together (110.3) must then equal X^ for NaCl, and, as we
see, they do. Similarly, the speeds of H* and C2H802' together (351.7)
must equal \x for HC2H302', although we cannot observe the latter
directly. This method can be applied to all of the less highly ionized
acids and bases, for their sodium and potassium salts belong invariably
to the class of substances which are most ionized, and for which, there-
fore, X^ can be determined accurately by extrapolation.
Degrees of Ionization of Common Substances. — The rule,
degree of ionization = X^/X^ (p. 328), enables us to calculate the
value for any dilution, when the necessary data are given. We need
only the values of the conductivity (XB ) for different dilutions (p. 328)
and those of the relative speeds of each kind of ions expressed in the
same units. The latter, when added, give X^.
Thus, hydrogen chloride in a solution containing 1 equivalent in
0.1 1. (365 g. per liter), which would be a rather concentrated lydro-
chloric acid, shows the degree of ionization ^ > ov 0,168 (= 16,8 per
330
INORGANIC CHEMISTRY
1.32
352
cent). Normal hydrochloric acid is ionized to the extent of ^"> or
0.784; normal sodium chloride, 7~> or 0.676; normal acetic acid,
0.004 (= 0.4 per cent).
Misapprehension easily arises in regard to the inferences that may be drawn
from a conductivity value. A single such value, say that for salt at 10 1. dilution
(92.5), gives no information about the extent of ionization. We must have the
value at infinite dilution as well, that is, we must have the other term of the ratio
corresponding to complete ionization, before the proportion of the molecules
ionized at the 10 1. dilution can be known. Further, we must have the values of
both for the same salt, at the same temperature and in the same solvent, for the
values at all dilutions change markedly when any one of these conditions is altered.
Thus the conductivity of normal sodium chloride solution at 50° is 120, and is
therefore actually greater than that at 18° when the dilution is infinite. But
at 50° the conductivity at infinite dilution is 185, so that at this temperature the
degree of ionization is ^f £ or 0.65, about the same as at 18°. On the other hand,
when a little alcohol is added to the aqueous solution, the conductivities all dimin-
ish. But that at infinite dilution diminishes also, so that the proportion of the
material ionized does not seem to be greatly affected. The chief effect of raising
the temperature is simply to diminish the friction opposing the motion of the ions
and, therefore, to increase the conductivity. The change is about 2 per cent for
each degree. Addition of alcohol, on the other hand, increases the friction and
diminishes the conductivity. There is, however, a real, though usually smaller,
change in the degree of ionization with change in temperature. When the tem-
perature is raised, the fraction ionized increases or diminishes according as the
heat of ionization is negative or positive (cf. p. 260), and conversely when the
temperature is lowered.
The following tables include the common reagents and give the
proportions of ionized molecules (total molecules = 1). Except where
otherwise specified, the data are for normal solutions at 18°. In the
case of acids containing more than one displaceable hydrogen unit,
the kind of ionization on which the figure is based is indicated by a
period. Thus H.HC03 means that the whole of the ionization is
assumed to be into H* and HCO/.
FRACTION IONIZED.
ACIDS.
Nitric acid 0 . 820
Nitric acid (cone., 62%) . .0.096
Hydrochloric acid . . . . 0 . 784
Hydrochloric acid (cone., 35%) 0.136
Sulphuric acid, H.H.SO4 . .0.510
Sulphuric acid (cone., 95%) . 0.007
Hydrofluoric acid . . . . 0 . 070
Oxalic acid, H.HC2O4 (N/10,
25°) 0.500
Tartaric acid, H.HT (N/10,
25°) 0.082
Acetic acid ...... 0.004
Acetic acid (N/10) . . . .0.013
Carbonic acid, H.HCO3 (N/10) 0.0017
Carbonic acid, H.HC03 (N/25) 0.0021
Hydrogen sulphide,H.HS(N/10) 0 .0007
Boric acid, H.H,B03 (N/10) .0.0001
Hydrocyanic acid (N/10) . . 0 . 0001
Permanganic acid (N/2, 25°) . 0 933
Hydriodic acid (N/2, 25°) . .0.901
Hydrobromic acid (N/2, 25°) . 0 . 8y9
Perchloric acid (N/2, 25°) . .0.880
Chloric acid (N/2, 25°) . . . 0 .878
Hydrochloric acid (N/2. 26°) . 0.876
Phosphoric acid, H.H.,PO4 (N/2,
25°) . . . ." . . .0.170
ELECTROLYSIS
331
BASES.
Potassium hydroxide .
Sodium hydroxide .
Barium hydroxide .
Lithium hydroxide .
Ammonium hydroxide .
Tetramethylammonium
droxide (N/16, 25°) .
hy-
0.77
0.73
0.69
0.03
0.004
0.96
Strontium hydroxide (N/64,
26°) 0.93
Barium hydroxide (N/64, 25°) .0.92
Calcium hydroxide (N/64,
25°) 0.90
Silver hydroxide (N/1783,
25°) . • 0.39
SALTS.
Sodium phosphate, Na2.HPO4
(N/32) 0.83
Cupric nitrate (N/16) . . . 0 . 80
Potassium chlorate (N/2) . . 0 . 79
Sodium tartrate (N/32, 25°) . (0 . 78)
Potassium chloride . . . 0 . 75
Ammonium chloride . . . 0 . 74
Sodium chloride . . . . 0 . 676
Sodium chloride (N/2) . . .0.734
Sodium chloride (N/10) . .0.839
Potassium nitrate . . . . 0 ; 64
Potassium acetate . .0.64
Calcium sulphate (N/100) . .0.63
Silver nitrate 0 . 58
Potassium sulphate . . . . 0 . 53
Sodium acetate 0.53
Sodium bicarbonate, Na.HCO3 (0.52)
Potassium carbonate . . . (0.49)
Sodium sulphate 0.446
Zinc sulphate 0.24
Zinc chloride 0.48
Cupric sulphate 0.22
Mercuric chloride . . . «0.01)
Mercuric cyanide Minute
Degree of Ionization of Water. • — If we consider a liter of water
as a normal solution in which 18 g. (one mole) represents the solute
and the rest stands for the solvent, the conductivity for complete
ionization into H* and OH' would be 318 + 174 = 492. The actual
ionization is one ten-millionth part of this. In other words, there is
only one ten-millionth of 1 g. of hydrion and the same fraction of 17 g.
of hydroxidion in a liter of water. A column of water 1 cm. long con-
ducts less well than a column of mercury of equal cross-section and
over 660,000 miles in length.
General Remarks on these Values. — It will be seen from in-
spection of the above numbers for acids that the proportion of the
molecules ionized in solutions of equivalent concentration varies
enormously. Roiighly, the acids might be divided into four groups :
those in which the ionization exceeds 70 per cent in normal solutions ;
those in which it lies between 70 and 10 per cent ; those in which
it lies between 10 and 1 per cent ; and those in which it is smaller
than 1 per cent. To the first class belong the acids which we
generally recognize as the most active in all their chemical relations,
namely, nitric acid, the halogen hydrides, and one or two others. To
the second class belong sulphuric acid and phosphoric acid, and they
are less active. Amongst the acids whose ionization lies between one
and ten per cent are such as hydrofluoric acid and acetic acid, and
332 INORGANIC CHEMISTRY
chemically they are weak. Carbonic acid and boric acid are of the
fourth class, and are feeble acids.
The bases, although less numerous, show that a similar division
might be made, although in the above list only two classes are
represented, — the strong bases, beginning with potassium hydroxide,
and the feeble bases, represented by ammonium hydroxide.
The salts show a much greater uniformity ; and, if the list had been
extended so as to include the hundreds of common salts in constant
use, the vast majority would have been found to show degrees of
ionization lying between 50 and 80 per cent. Only a few fall below
these limits. The salts of mercury are almost the only ones which
would belong to the class of least ionized substances. Salts, like zinc
sulphate and cupric sulphate, in which both ions are inultivalent, are
always much less highly ionized than are salts (e.g. ZnCl2) made up of
either of the same ions along with a univalent ion. Salts, however,
are ahnost never restricted in their degree of ionization, to an extent
sufficient to produce any noticeable effect on their chemical properties
(see, however, Cadmium iodide and Mercuric cyanide).
The relation between degree of ionization and prominence of acid
or basic chemical properties will be developed in the next chapter.
Comparison with the Results Obtained by Other Methods. —
The value for the degree of ionization as measured by the conductivity
method is coincident with that found for the same solution by a study
of the abnormalities in freezing- and boiling-points and in osmotic
pressure (Chap. xvii). The electrical method gives accurate results
more easily than do the others, however, and is therefore the one most
frequently used. It was Svante Arrhenius, a Swedish chemist, who,
in 1887, first noted the coincidence in the values and devised the ionic
hypothesis to account for it. From the appearance of his remarkable
memoir we date the great development which the study of solutions *
has undergone in recent years.
Exercises. — 1. Name (p. 322) the ionic materials furnished by the
dissociation of potassium bromate, silver bromide, sodium periodate
(NaI04), permanganic acid.
2. Give lists of other anions and cations which have been en-
countered.
* The Scientific ^femoirs, No. IV. (American Book Company), is a reprint
of the fundamental papers by Raoult, van 't Hoff, and Arrhenius.
ELECTROLYSIS 333
3. How many coulombs are carried by and will deposit : 20 g. of
silver, 15 g. of antimony, 30 g. of chlorine, 60 g. phosphanion
(P04)?
4. What current strength (in amperes) is required to deposit :
20 g. of silver in an hour, 100 g. of iodine in 5 minutes, 60 g. of anti-
mony in 3 hours ?
5. What is the percentage of molecules ionized in : deci-normal
(N/10) sodium chloride, centi-normal (N/100) acetic acid, centi-normal
hydrochloric acid (p. 328) ?
6. Give an experimental definition of the term ion. That in the
text (p. 321) is in terms of the hypothesis.
CHAPTER XX
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES
BEFORE considering the typical interactions of ionogens in solution,
we must have a clear conception of the peculiarities of these bodies
which are likely to affect their behavior. The facts on which such
a conception must be based have been given in preceding chapters, and
all that is now necessary is to collect and apply these facts. On
account of the coherence which they give to the subject, the figures
of speech of the ionic hypothesis will be largely employed. The
reader will, therefore, do well to exercise especial care to distinguish
fact from fiction.
In this discussion it must be made clear that aqueous solutions
of ionogens are mixtures containing several solutes. It must also
be shown that each kind of ions is a distinct substance with indi-
vidual physical and chemical properties. Next, salts being used for
illustration, the commonest kind of interaction, double decomposition
between ionogens, will be discussed. In this connection precipitation
brings up the peculiar state of equilibrium between the undissolved
solute and the complex of molecules and ions in solution. Applica-
tion of the same principle to special cases, such as those of acids and
bases, then follows.
, The discussion of systems in equilibrium in the present chapter
will be purely qualitative. The quantitative consideration of ionic
equilibria (cf. p. 297) is postponed until the study of the metals and
their compounds is taken up (see Chap, xxxiv).
Solutions of Ionogens are Mixtures. — We are accustomed to
regard a bottle of sodium chloride solution as containing but one thing,
aside from the water. We must now think of it as containing
at least three dissolved substances, any one of which might be alone
responsible for some property of the solution. The same idea must
accompany our use of every solution of an ionogen. Thus, in ordi-
nary experiments, in which solutions of concentration not far from
normal are commonly used, we have something like the following
proportions (p. 330) of the three main components in six typical cases :
334
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 335
>Na* + Cl' (68%), (78%)CuS04<=>Cu"+ SO/'(22%),
(22%)HC1 «=±H- + C1' (78%), (99.6%)HC2H3O2^H' + C2H302'(0.4%),
(23%)KOH<z»K'+OH'(77%), (96.6%)NH4OH<z±NH4' + OH' (0.4%).
In solutions made from salts, the greater part, and by far the most
active part, of the contents is almost always ionic. Cupric sulphate,
being a salt both of whose ions are bivalent (p. 332), practically
illustrates the lower limit as regards quantity of ions, for only a few
salts of mercury and cadmium fall far below it. The acids and bases
have a wider range, and a larger proportion of them are like acetic
acid and ammonium hydroxide respectively. Still, even when small
in amount, the ions of acids and bases are almost always much more
active than the molecules.
The presence of still other components in solutions of salts, arising
from interaction with the water, will be noted later.
While most of these facts are ascertained by more or less remote
inference from physical and chemical properties of the solutions,
some of them are evident to the eye in certain cases. Thus the prog-
ress of the ionization of a salt may be seen if one of the ions is
different in color from the molecules. Cupric bromide in the solid
form is a jet black, shining, crystalline substance. When treated
with a small amount of water it forms a solution which is of a deep
reddish-brown tint, giving no hint of resemblance to a solution of any
cupric salt. This doubtless represents the color of the molecules.
When more water is added, the deep brown gives place gradually to
green, and finally to blue. The latter is the color of the cuprion
(Cu"), and is familiar in all solutions of cupric salts. The colorless
nature of solutions of potassium and sodium bromides shows that
bromidion (Br') is without color. Hence, in the present instance it is
invisible. We are thus watching the progress of the action :
CuBr2 <=» Cu" + 2Br'.
If 1 g. of the solid is taken, it dissolves in about its own weight of
water, and independent measurement shows that there is relatively
little ionization. Hence the solution is deep brown. When 10 c.c. of
water has been added, 70 per cent of the salt is ionized, and the solu-
tion is green. With 40 c.c. of water, only 19 per cent remains in
molecular form, and the blue color of the cuprion entirely overbears
the tint of the molecules. If, at the green stage, we dissolve solid
potassium bromide in the liquid, the high concentration of bromidion
336 INORGANIC CHEMISTRY
which results causes au extensive reversal of the dissociation (cf. p.
250), and the molecules, with their brown color, become prominent again.
Sufficient final dilution with water, however, reduces the concentrations
of all the ions once more, the molecules dissociate, and the brown
color is displaced by the blue for the second time.
Each Kind of Ion in a Mixture Acts Independently. —
Numberless facts show that each kind of ion, for example cuprion, has
an individual set of physical and chemical properties and behaves in
many ways as if alone present in the solution. We shall meet with
much evidence of this in the sequel. Some facts tending to prove it,
that have already been given, may be recalled (cf. p. 295).
If, in comparing the migration speeds of any element, say copper,
in different salts (p. 314), they were the motions of substances like
Cu(NOg)2, CuBr2, CuS04, that we were comparing, all analogy teaches
us that the speeds with which they would move should vary widely.
That the blue color drifts always at the same pace shows that it is the
same substance, namely, cuprion (Cu"), that we are observing.
If, in solutions of the different permanganates, KMn04, NaMn04,
Ba(Mn04)2, and so forth, the dissolved bodies were different in each
case, we should confidently expect the purple colors of the solutions to
differ markedly in shade. But, for dilute solutions of equivalent con-
centrations, when strict examination is made, the tints are found to be
absolutely identical. We are therefore simply comparing different
mixtures all containing the same proportion of the same free, colored
body, Mn04'.
In phosphorus pentachloride vapor (p. 255), the fully liberated
trichloride and chlorine are prominent components. Diminishing the
volume of a fixed amount of this mixture, by compression, throws
more chlorine into combination and the total absorption (from which
the greenish-yellow color is derived) becomes less, the compounds of
phosphorus being both colorless. Increasing the volume, on the other
hand, promotes the dissociation and increases the total absorption. The
system of ions and molecules in equilibrium in a solution of cupric*
bromide, or any other ionogen, behaves in exactly the same way. The
components possess and exhibit individual properties, much like the
components of a gaseous mixture (p. 155), both in this and in other
respects.
All solutions of acids are sour in taste, irrespective of the nature
of the negative ion, while salts containing the same negative radical
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 337
are not sour at all. Hence in solutions of acids we are tasting the
same free substance, hydrion (H*). Similarly, in solutions of all
alkalies, we note the soapy taste of hydroxidion (OH').
These illustrations concern physical properties. In the next sec-
tion we shall learn that an ionic material, such as bromidion or cu-
prion, has specific chemical properties irrespective of the nature of its
concomitants.
SALTS, IONIC DOUBLE DECOMPOSITION, PRECIPITATION.
Salts. — We have already seen that salts differ much in solubility
in water, that some comjfcine with water to form solid hydrates (p. 120),
that they interact with acids, bases, and other salts by exchange of
radicals, reversibly (p. 281), and that they are all ionogens.
Both the positive and negative ions of salts may be simple or com-
posite, Na.Cl, Na.N03, NH4.C1, NH4.N08. The elements which can
form a simple positive ion are known in chemistry as metals (p. 119,
and see Chaps, xxiii and xxxii). Non-metals, like nitrogen, may be
present in a positive ion, as in NH4*, but never exclusively. In other
words, we know no such substances as nitrogen sulphate, or carbon
nitrate. Metals, on the other hand, are frequently found in the nega-
tive ion, but never constitute it exclusively. They are then usually
associated with oxygen, as in MnO/, and Cr,07". Some ionic mate-
rials are colored, Cu*" blue, Cr*" reddish violet, Co" pink, MnO/
purple, Cr207" orange, but most of them are colorless, K*, Na', Zn",
CF, I', NO/. The ions of salts do not affect litmus. They vary in
taste, some being salt, some astringent, some bitter. In dilute solu-
tions they are almost always numerous in comparison with the surviv-
ing molecules. They carry electricity, but relatively less well than
do hydrion and hydroxidion, on account of their slower migration. All
the known ionic materials are found in solutions of salts. The only
ions which are not characteristic of salts, although sometimes occur-
ring in their solutions (see Mixed ionogens), are hydrion H*, and
hydroxidion, OH'.
Double Decomposition of Salts in Solution. — When we mix
sodium chloride (NaCl) and silver nitrate (AgNO3), both in solution in
water (H20), there would seem to be many different possibilities of
union amongst the six elements represented in the mixture. But only
one sort of change occurs, and it takes place almost completely:
AgNO3 + NaCl fc> AgCl J + NaNO3. ( 1 )
338 . INORGANIC CHEMISTRY
So in general, we find that a single crosswise union of ions with oppo-
site electrical charges is by far the commonest kind of interaction
between ionogens.- There are indeed four other kinds of ionic chemical
change, as will be seen in the sequel, but for the present we shall dis-
cuss only the cases of double decomposition.
When solutions of two ionized substances are mixed, the first reflec-
tion which occurs to us is that each of these has been diluted by the
water in which the other was dissolved, so that the first effect will be
to increase the degree of ionization of both to a certain extent. The
next consideration is, however, that we have produced a mixture of
four ions, which must have at least some tendency to unite crosswise.
Thus potassium chloride and sodium nitrate in dilute solution are very
greatly ionized before mixing. The reversible actions, represented by
the horizontal pair of the following equations, have taken place exten-
sively. But, by mixing the liquids, we have brought into presence of
one another two new pairs of positive and negative ions. Hence, two
other reversible actions, the vertical ones,
KC1 *=> K- + Cl'
NaN03<=»N08' + Na-
il IT
KNO3 NaCl
will be set up and will proceed until a fresh equilibrium of all the ions
with all four kinds of molecules has been reached. Examination of
the solutions of these four salts, separately, shows that they are in an
equal degree extensively ionized in dilute solutions, so that in this par-
ticular case the whole quantity of molecules of all kinds will not be
very great. That this inference is correct is shown by much independ-
ent evidence, of which two samples may be given.
The change of ionic materials into molecular and vice versa is always
accompanied by absorption or liberation of heat (p. 330). Now it
was for long a matter of surprise that when dilute solutions of salts,
as distinct from pairs which included acids or bases (g'.v.), were mixed,
no heat-change was observable. This fact was called the thermoneu-
trality of salts, but the reason for it was unknown. Since salts are all
highly ionized, the reason is now apparent. Similarly, no changes in
.color or volume accompany the mixing of solutions of salts.
Again, any of the means which may be used for measuring the
number of molecules, including ions, in a solution, may be applied to
learning whether any appreciable proportion of the latter has disap-
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 339
peared. Osmotic pressure, freezing-point, boiling-point, and conduc-
tivity are all applicable. The last lends itself best to the purpose of
demonstration. The cell and amperemeter described in the last chap-
ter (Fig. 81) may be employed. We place in the cell a one-fourth
normal (N/4) solution of potassium chloride and introduce an equal
volume of N/4 sodium nitrate, in such a way that it forms a separate
layer beneath the other solution. We read the amperemeter, mix the
liquids by stirring, so as to permit chemical interaction to take place,
and then note the conductivity once more. There is no observable
change in the conducting power, and, therefore, no appreciable change
in the condition of the ionic substances has taken place.
We shall have occasion in a later paragraph to show that, where
the ions unite crosswise to an appreciable extent, an experiment of
this kind shows a marked diminution in the conductivity, correspond-
ing to the amount of combination that has occurred. In such cases,
also, heat is either liberated or absorbed.
It thus appears that when dilute solutions of salts are mixed, and
there is no visible evidence of chemical change, the little that has taken
place may be neglected. Practically, mixing of this kind is a physical
operation. Of course, in the event of one product being precipitated,
a diminution in the conductivity, corresponding to the amount of ionic
material removed, will be observed.
In view of the above explanation, the old question of whether such a solution
contains the first pair of salts, or the second pair, represented in the double decom-
position, KC1 + NaNO3 «=> KNO3 + NaCl, loses its whole point. The solution con-
tains neither the initial molecular substances nor the molecular products, in
appreciable amount.
Solution and Precipitation of Salts. — In a body which is dis-
solving or being precipitated, we are observing the progress of a rever-
sible physical operation (p. 153) :
Molecules (undiss'd) ^± molecules (diss'd).
If the molecules are sparingly soluble, the forward action is feeble,
while the backward one, when we start with the same material in solu-
tion, makes great progress. Conversely, when the substance is a sol-
uble one, the forward action comes to a standstill (or the reverse action
occurs) only when the concentration of dissolved molecules has be-
come very large.
Now it will be noted that this mechanical adjustment concerns
only the molecules, and that the ions, if there are any, are involved
340 INORGANIC CHEMISTRY
only indirectly. The ions are in equilibrium with the dissolved mole-
cules :
AB (undiss'd) <=± AB (diss'd) <-> A' + B'.
Hence, when a substance dissolves, it does so in molecular form, and
ions are subsequently generated from some of these molecules until
equilibrium is reached. Conversely, when molecules come out of so-
lution, as the result of cooling, for example, the diminished concentra-
tion of the central term of the chain enables more ions progressively
to unite until the whole system has adjusted itself to the new condi-
tions.
Now this has an important bearing on the result of mixing dilute
solutions of two soluble salts. We have seen that the concentration of
the molecules of the new pair of salts is never large. Yet it may
easily be in excess of the amount which the water can hold in solution,
if one of the salts is of the relatively insoluble class. This occurs,
for example, when a chloride is mixed with a salt of silver. The sys-
tem of equilibria, leaving out that of Na* and N03', and rearranging
(cf. p. 338), so as to save space, appears as follows :
The concentration of dissolved silver chloride which the solid can
maintain in solution being very minute, most of this salt is at once pre-
cipitated. The ions continue to unite because the requisite concen-
tration of molecules, whose dissociation should bring their union to a
standstill, has not been kept up. The new molecules are in turn pre-
cipitated. The system reaches a stable condition only when the con-
centration of chloridion (CF) and argention (Ag*) has fallen to that
which can be maintained by the mere trace of molecules which the in-
solubility of the substance permits to remain in solution. When the
system of equilibria is examined, we see at once what the result must
be. The removal of chloridion and argention enables the remaining
molecules of sodium chloride and silver nitrate to become completely
ionized. Thus the concentration of NaCl and AgN03, of Ag* and Cl'
and of the dissolved AgCl, all become practically zero. The system
finally contains only molecular, solid silver chloride and the three sul>-
stances, Na* + NO/ ±3 NaN03, in equilibrium, of which by far the
greater part is the ionic. There is therefore in the action much detail
which the first equation (p. 337) did not show.
THE CHEMICAL BEHAVIOK OF IONIC SUBSTANCES 341
The above explanation will apply to any case of precipitation
resulting from the interaction of ionogens. If the least soluble of the
four salts is more soluble than silver chloride, more concentrated solu-
tions are required to secure precipitation. The interaction of hydro-
gen chloride and sodium hydrogen sulphate (p. 179) is of this nature :
iss,dx _, NaC1
NaHS04 <=> HSO/+ Na' )
It should be noted that, strictly speaking, the only interaction tak-
ing place when the solutions are mixed is the production of the insol-
uble body. In the case of silver chloride, (1) and (2), the largest part
of the chemical action may be formulated thus :
Ag' + Cl'^AgCl. (3)
The chief change that has as yet befallen the ions of sodium nitrate is
that they have been transferred from two separate vessels into one.
Potentially the salt has been formed. But the actual union of its ions
to give the second product in the molecular condition :
Na' + N03' -> NaN08, (4)
comes about only when, at some subsequent time, if at all, the water is
evaporated away.
Individual, Specific Chemical Properties of Each Ionic
Material. — We wrote the equation for the formation of silver chlo-
ride (Ag* -f Cl' — » AgCl) as if argention and chloridion were the only
substances concerned in the action. Further study shows this to be
justifiable. Thus, hydrochloric acid, cupric chloride, and dozens of
other chlorides may be used instead of sodium chloride and give silver
chloride just as readily. The natrion had nothing to do with the
result. Of course we cannot get a solution containing chloridion
alone. Like a vessel in which to make the experiment, some positive
ion is required. But, like the rest of the apparatus, this ion may be
varied indefinitely, is not altered in the course of the change, and may
therefore be dispensed with in the equation. The nitranion (NO/)
which accompanied the argention is similarly a part of the apparatus,
for silver sulphate solution works just as well as silver nitrate.
That chloridion is a substance with specific chemical properties, is
easily demonstrated. It forms silver chloride whenever it encounters
argention. Other substances, even when they contain chlorine, lack
342 INORGANIC CHEMISTRY
this property. Chloroform (CHClg) and chlorobenzene (C6H5C1), in a
solvent in which ionogens are dissociated, do not interact when silver
nitrate is added. They give no chloridion, and, in fact, remain un-ion*
ized. Potassium chlorate (KC103) and perchlorate (KC104) and chlor-
acetic acid (HC2H2C102), with argention, fail likewise to give silver
chloride. They are ionized, but chloridion is not one of the ions
of any of them. The ions CIO/, CIO/, and C2H2C10/, have properties
of their own, and their compounds with argention are soluble.
Other chemical properties of chloridion are : That it unites also
with plumbion (Pb**) and monomercurion (Hg*), forming insoluble
chlorides (p. 185). It is discharged and liberated as free chlorine by
fluorine (p. 241) :
2H- + 2C1' + F2 -» 2H* + C12 + 2F'.
Since the hydrion is not affected and many chlorides behave in a
similar manner, the positive ion may be omitted :
2C1' + F2 -+ C12 + 2F'.
Finally, chloridion has relatively little tendency to unite with other
ions, or, in other words, the compounds of chloridion with most other
ions are highly ionized. Thus it combines with hydrion to the extent
of only 22 per cent (p. 330) in normal solution. In this respect it
differs markedly from free chlorine, just as hydrion differs from
hydrogen. The free elements unite with vigor and completely. Hy-
drogen chloride is easy to dissociate into ions, but difficult to dissociate
into its constituent elements. Nothing could show more strikingly
than this that the ionic materials have chemical properties of their
own.
Similarly, barium salts and ordinary siilphates give, when mixed, a
precipitate of barium sulphate. Here we encounter a property of ba-
rion (Ba") and sulphanion (SO/'). But potassium ethyl sulphate
(KC2H5S04), in spite of its name, will not give this reaction with
a barium salt. Here electrolysis shows that sulphanion is absent and
that the negative ion is C2H5SO/.
In the same way every other ionic material may be shown to be a
substance with an individual set of physical (p. 336) and chemical
properties. Each salt, when dissolved, gives two kinds (see, however,
below) of ionic materials. The solution is simply a mixture, and each
physical component forthwith behaves towards ions capable of uniting
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 343
with it, as if it were alone. The other materials, ionic and molecular,
which are present, may remain essentially unaffected throughout the
' change.
Application in Chemical Analysis. — Since the larger number
of ordinary chemical substances are ionogens, and the most rapid and
simplest chemical changes take place when they are in solution,
the various reactions of- their solutions are employed as tests for the
substances in question. An advantage of the use of the solutions is
that they contain a mixture of two independent materials, the anion
and the cation, and when these have been identified successfully the
salt from which they were formed is known. The simplicity to which
chemical analysis is thus reduced may be seen when we consider that
twenty-five common metals with twenty-five negative radicals might give
a total of over six hundred different salts. If the distinct properties of
each of these had to be considered, the identification of an unknown
substance would be very difficult. In solution, however, the problem
becomes much easier. Every solution made from a single salt will con-
tain but two substances (in the main ; see, however, below), and the
problem reduces itself to ascertaining which two, out of a total of
fifty, are present in any particular case.
As an example of the method, let us suppose that we look first for the positive
ion. Most systems of analysis begin by the addition of a solution containing
chloridion, generally dilute hydrochloric acid, to the liquid. If an ion is present
which in combination with chloridion gives an insoluble compound, a precipitate
will appear. Amongst the common positive ions but three are of this kind, namely,
argention, monomercurion, and diplumbion. So that the precipitate, if it appears,
is a chloride of one of these three metals, and the matter of distinguishing between
the three is quickly disposed of by further examination of its properties. If no
precipitate comes out, then these three metals are probably absent, and some fresh
ion capable of precipitating another set of positive ions is introduced (see Chap,
xxxvii). Thus by a process of elimination we quickly find out whether any metal
ion is present, and, if so, precisely which one it is.
The language of analysis is frequently somewhat loose. Thus we speak of the
addition of a silver salt to a solution as being a " test for chlorine." As a matter of
fact, it is not a test for chlorine. It is not intended as a test for free chlorine,
nor will it show the presence of chlorine in many states of combination. It is simply
a test for ionic chlorine (Cl'), and cannot give us information in regard to the presence
or absence of any other form of the element. So the wet-way tests for " copper,"
"silver," etc., so called, are tests for the ionic forms of these elements, and not
for the presence of the element in every form. Even the two kinds of copper and
mercury ions, Cu", Cti", Hg", Hg*, must be classed as distinct substances. Thus,
the last is precipitated by chloridion while the second last is not, mercuric chloride
(HgCLj) being soluble.
344 INORGANIC CHEMISTRY
Hydrolysis of Salts. — The natural ionization of water is very
slight, but there are cases in which its effects become noticeable, and
the interaction of its ions with those of dissolved salts cannot be neg-
lected. For example, an aqueous solution of pure cupric sulphate is
always acid and therefore contains hydrion :
CuSO^SO/'+Cu" >
2H.O t, 2H* + 20H' | ** Cu(°H>* <dM8>d>-
Cupric hydroxide, being a very feeble base, and comparable with water
itself in the small extent to which the solvent is able to hold its ions
apart, is formed to a small extent. The removal of some hydroxidion
by this means enables more of the water to dissociate. This, in turn,
furnishes the material for the production of more cupric hydroxide.
The action does not proceed very far, but it makes sufficient progress
to leave a perceptible excess of hydrion in the liquid and to give it,
therefore, an acid reaction. The hydrion combines slightly, but only
slightly, with the sulphanion, for sulphuric acid is a highly ionized
acid. This part of the action has, therefore, been left out of the dia-
gram. The ordinary equation for this change would be :
CuS04 + H20 z+ Cu(OH)2 + H2S04.
The hydrolysis is much, greater with sodium sulphide (q.v.) and anti-
mony trichloride (y.v.).
Again, soap solution is always faintly alkaline :
t, Na" + C15H31C02'
The sodium palmitate is highly ionized, but palmitic acid (HC15H31C02)
is hardly ionized at all. The final result is the production of a recog-
nizable amount of hydroxidion in the solution. Thus, a salt derived
from an acid and a base of very different degrees of activity, whether
it is the base (as Cu(OH)2) or the acid (as palmitic acid or hydrogen
sulphide, q.v.~) which is the weaker member, is likely to be more or
less hydrolyzed by water. In the former case the solution is acid, in
the latter basic in reaction. Other things being equal, salts containing
bivalent or trivalent radicals are more noticeably hydrolyzed than are
those composed only of univalent radicals.
Cases of this kind being common, we are thus compelled to enlarge
our list of possible components in the solutions of any salt. In addi-
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 345
tion to the molecules and ions of the salt, there are present, water and
its ions, and the molecules of the base and acid formed by the union
of the latter ions with the former. There are thus no less than eight
different components in the mixture.
ACIDS AXD BASES AND THEIR DOUBLE DECOMPOSITION WITH SALTS.
Hydrogen Salts. — The substances of the composition HC1,
H2S04, and so forth, are commonly called acids, and when more con-
venient we shall conform to this usage. But it is only when they
have been dissolved in water or some other ionizing solvent that they
show the properties characteristic of acids. In fact, in terms of the
ionic hypothesis, there is only one acid, hydrion (H*), although the
substances which give it by dissociation are many. The parent sub-
stances are salts of hydrogen, in which the element hydrogen plays
the part of a metal.
The properties of the hydrogen salts, that is, of the original iono-
gens, are the same as those of any other ionogens. They are distin-
guished from other ionogens by the fact that their positive radical
is always hydrogen and that in solution they yield hydrion. Their
negative radicals are all different, Cl, Br, I, C103, C104, Br03, N03, and so
forth. These radicals form the negative ions in solutions of hydrogen
salts. In such solutions all the properties of these ions are the same
as when they are furnished by dissolving other salts containing the
same radicals. Some hydrogen salts have oxidizing powers like hypo-
chlorous acid (p. 269). Usually, they exchange radicals with the
other ionogens. They often do this even when dissolved in non-dis-
sociating solvents. They frequently do it also in the absence of a
solvent, especially when heated. They differ from one another in the
matter of solubility in water, some being almost insoluble. By solu-
tion in water they give acids of very different degrees of activity (see
Activity of acids, below).
Hydrion. — Hydrion is a colorless substance which exists in water
and certain other solvents only. It is always associated with an equiv-
alent amount of some negative ion. It is sour in taste, and its pres-
ence is recognized by the fact that it turns blue litmus red and decolor-
izes pink phenolphthalein (see Indicators, below) solutions. It confers
a high conducting power upon solutions in which it is contained, on
account of its great speed of migration. It is univalent and combines
346 INORGANIC CHEMISTRY
with negative ions, such as hydroxyl and the negative radicals of salts.
It is displaced by metals like magnesium. In all these respects it differs
markedly from free hydrogen gas.
The displacement of hydrogen from dilute acids (p. 95) now ap-
pears in a new light. The action will be formulated thus :
Zn + 2H- + SO/' -> Zn" + S04"+ H2.
The sulphanion (SO/') , although zinc sulphate is somewhat less ion-
ized than sulphuric acid, is not much affected by the change and may
be omitted :
Thus, this action, which takes place in the same fashion with most
acids, is seen to be independent of the nature of the negative ion. It
consists simply in the transference of the electric charges from the
hydrogen to the zinc, whereby the latter becomes ionic. The dis-
charged hydrogen is liberated as gas. When the solution is evapo-
rated, the ionogen, in the above case zinc sulphate, is formed :
Zn"+S04"->ZnS04.
Modes of lonization of Acids. — An acid containing but one
unit of hydrogen in its molecule can give but two kinds of ions.
Thus, chloric acid gives only H* and CIO/. When more than one
hydrogen unit is present, however, more than two kinds of ions are
formed. Thus, sulphuric acid, H2S04, produces, in the first place, hydro-
sulphanion :
H2SO4<->H' + HSO/.
The latter is also an acid, but is considerably less active than sulphuric
acid. Hence, the further dissociation of this ion (HSO/<=± H*+ SO/')
lags considerably behind the primary dissociation. In concentrated
solutions of the acid there is, therefore, much HSO/ present. In very
dilute solutions, however, SO/' predominates. We know that HSO/
is a weaker acid, and is dissociated with greater difficulty by water,
because acid salts (see below), like KHS04, which give this ion, are
much weaker acids than are acids like HC1 and HC103, with which
the substance HSO/ might fairly be compared. This behavior is not
peculiar to sulphuric acid, but is shown by all acids containing more
than one hydrogen unit in the molecule (cf. Hydrogen sulphide).
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 347
Activity of Acids. — In solutions containing equivalent quantities
of hydrogen salts, and therefore equal amounts of combined hydrogen,
in equal volumes, the concentration of hydrion present at any moment
in each will be different. This concentration will be high or low
according to the extent to which water is able to dissociate the mole-
cules. Ifow the activity of the hydrion, that is, the speed with which
it will interact, like that of any other substance, depends on its con-
centration (p. 250). Hence the hydrogen salts furnish, on being dis-
solved, acids of all degrees of activity. Thus in normal hydrochloric
acid, the fraction dissociated is 0.78, and the hydrion is 0.78-normal,
whereas in normal acetic acid the hydrion is only 0.004-normal
(p. 330). Yet the amounts of hydrogen chloride and hydrogen acetate
per liter contain equal quantities of eombined hydrogen, namely, 1 g.
each. Both the solutions in fact are normal in respect to combined
hydrogen. But the normal acetic acid has only about one two-hun-
dredth of the activity of normal hydrochloric acid.
That a difference in the activity of different acids does exist may
be shown, roughly, by placing similar pieces of the same metal, say zinc,
in equal volumes of various normal solutions of acids, such as
hydrochloric, sulphuric, and acetic. The hydrogen is evolved more
rapidly by the first than by the second, and very much faster by
either than by the last. Naturally, the first is sooner exhausted, while
the third acts in its slow way for a very long time before being all
used up. In the third case few ions of hydrogen are at hand at any
one moment, but more are formed continuously from the molecules, to
take the place of those displaced. Thus the total amount of hydrogen
obtained from each acid is finally the same. It is the speed of evolu-
tion alone which is different and shows the differing concentrations of
the hydrion.
In cases of extremely small ionization, the presence or absence of
visible action on litmus may torm another means of estimating
activity. Thus, litmus is easily turned red by a deci-normal solution
of acetic acid or of any more active acid (p. 330), but hydrogen sul-
phide, in a solution of the same molecular concentration, contains only
one-twentieth as many hydrogen ions (p. 330), and affects litmus paper
but slightly. Paper dipped in Congo red exhibits differences in
the activity of acids by the different depths of the tints it assumes.
For example, it is much less markedly affected by acetic than by sul-
phuric acid' of the same concentration (see Indicators, below).
Many hydrogen salts are but slightly soluble. Thus, with silicic
348 INORGANIC CHEMISTRY
acid (<?.#.), the solid can keep only a small concentration of molecules
in solution: H2SiO3 (solid) ±=> H2Si08 (diss'd). So that, although some
ions are doubtless present, H2Si03 (diss'd) <=i 2H* + Si03", their con-
centration, being dependent on that of the molecules, is very minute
indeed. Still, even in the absence of an effect upon litmus, the sub-
stance can be recognized to be an acid. Thus, by the action of sodium
hydroxide, silicic acid can be made into sodium silicate, NagSiO^ which
is highly soluble and highly ionized. Hence, since Si03" is a negative
ion, we reach the conclusion indirectly that H2Si03 is an acid.
Substances like ammonia NH3, sugar, and alcohol, although they
contain hydrogen, are not hydrogen salts. They are not ionogens
(cf. p. 281), and give no hydrion. lonizable and non-ionizable hydrogen
may even be contained in the same compound. Thus, each molecule of
acetic acid (HC2H302) contains four hydrogen units, but gives only one
hydrogen ion. The other three are part of the acetanion (C2H802').
We infer this because metals can be substituted for one hydrogen unit
(NaC2H302), but not more.
Salts of Hydroxyl. — Substances like potassium hydroxide, am-
monium hydroxide (NH4OH), and zinc hydroxide (Zn(OH)2), are
commonly called bases. But it is only in their aqueous solutions that
the basic properties appear. There is only one base, namely, hydroxyl
(OH'), and these substances are simply the source of it. The parent
substances are salts of some metal, or group playing the part of a
metal (e.g. NH4), in which hydroxyl is the negative radical.
The more active bases are called alkalies, sometimes caustic
alkalies and, individually, often, 'caustic potash, and caustic soda.
The solutions are called lyes.
The name "base " was originally applied to the non-volatile, and therefore seem-
ingly more fundamental part of a salt that remained behind when the salt was
heated. Usually the negative radical is disintegrated, as in heating calcium carbon-
ate (<?.«.)• But, as a matter of fact, it is generally the oxide aiid not the hydroxide
of the metal that remains. Still, the oxide, formerly named the base, often readily
gives the hydroxide (cf. p. 119) of which the term " base " is now used, and behaves
similarly to it in. many interactions (cf. p. 186).
The salts of hydroxyl have the properties common to all ionogens.
They are distinguished from other ionogens by the fact that their
negative radical is always hydroxyl and that in solution they yield
hydroxidion. Their positive radicals are all different, K, Na,'NH4,
Zn, Cu, etc., and constitute the positive ions of solutions of the bases.
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 349
All the properties of these ions in such solutions are the same as
when these ions are formed in solutions of salts containing the same
radicals. Usually the salts of hydroxyl undergo double decomposi-
tion -with other ionogens, exchanging radicals, even in absence of a
solvent. Like hydrogen salts (p. 347) they are ionized to different
degrees when in solutions of equivalent concentration, giving solutions
with different concentrations of hydroxidion. Hence, they give
solutions of different degrees of basic activity.
The common bases, with the exception of the hydroxides of
potassium, sodium, barium, strontium, calcium, and ammonium, are
but slightly soluble in water. Hence, zinc hydroxide, for example,
although it dissolves sufficiently to enable chemical action to take
place slowly, does not give enough hydroxidion at one time to affect
litmus paper. Magnesium hydroxide and lead hydroxide turn red
litmus paper blue with difficulty. Doubtless the few molecules that
do dissolve are almost all ionized :
Mg(OH)2 (solid) <=; Mg(OH)2 (diss'd) <=» Mg" + 20H',
but all the dissolved materials put together (0.01 g. per 1.) will
scarcely be weighable unless a considerable volume of the solution is
evaporated.
Hydroxidion. — Hydroxidion is a colorless substance found only
in water and certain other solvents, and is always associated with an
equivalent amount of some positive ion. It possesses a soapy taste,
and turns red litmus blue and colorless phenolphthalein pink (see
Indicators, below). It confers great conducting power upon solutions
in which it is contained, on account of its speed of migration which is
second only to that of ionic hydrogen. It is univalent and unites with
positive ions.
Ionic Double Decomposition and Precipitation of Acids and
Bases. — Under this head, we shall discuss only the interaction of a
salt with an acid or base, reserving for a separate section, on neutral-
ization, that of acids and bases with each other.
When a highly dissociated acid is mixed with a salt, a reversible
action tending to form another acid and salt is set up (p. 264).
Such an action is that of nitric acid on a hypochlorite (p. 267) in
dilute solution :
KOC1 1=5 KN03 + HOC1,
350 INORGANIC CHEMISTRY
giving potassium nitrate and hypochlorous acid. In such a case, if
the products are both as highly ionized as the initial substances, the
result is similar to that of the interaction between potassium chloride
and sodium nitrate (p. 338). No decisive change takes place.
With hypochlorous acid, however, which is very slightly ionized,
the result is different :
This acid is promptly formed from its ions, and the final mixture con-
tains, mainly, K", N08' and molecular HOC1. Yet, since the substance
is soluble, no outward evidence that the action differs from that of
potassium chloride and sodium nitrate is visible. The conductivity of
the mixture (p. 339), however, is found to have been greatly reduced
by the removal of half the ions, including the most rapidly migrating
of the four, hydrion (p. 314).
When the molecules of the resulting acid are insoluble, then it may
be precipitated (cf. silicic acid), after the manner of silver chloride
(p. 340), or may escape if a gas (cf. hydrogen sulphide), irrespective
of its degree of ionization.
In the same way, when a salt and a base are brought together, a
base and a salt are produced. All that has been said in the preceding
paragraph applies to this case also. Thus ammonium hydroxide (q.v.),
which is a feebly ionized base (p. 331), is formed on this plan, by mix-
ing solutions of an ammonium salt and a strong base :
NH4C1 + NaOH t=> NaCl + NH4OH.
When the resulting base is insoluble, like zinc hydroxide, it is pre-
cipitated, and the action becomes nearly complete on this account and
irrespective of the degree of ionization.
Ionic Double Decomposition and Activity. — It is quite clear
that the complete formation of acids, bases, and salts by precipitation
is purely a result of mechanical details concerning solubility, and
shows nothing about the degree of affinity between the constituent
ions. Again, the union of ions to form feebly ionized substances only
shows the tendency of the ionic materials to unite and may be com-
plete where the free elements have little mutual activity, and vice
versa. Thus, hydrion and hypochlorosion (CIO') unite almost com-
pletely, while hydrion and chloridion hardly unite at all. Yet hypo-
THE CHEMICAL BEHAVIOK OF IONIC SUBSTANCES 351
chlorous acid is very unstable, while hydrogen chloride is just the
reverse. Ionic double decompositions, consequently, give no clue to
the activities of the free materials.
NEUTRALIZATION.
Neutralization. — When 80 per cent sulphuric acid is poured
upon solid potassium hydroxide, much heat is developed and clouds of
steam arise. The solid product, when freed from the rest of the water,
is potassium sulphate. The proportions of the materials used and pro-
duced are shown by the equation :
2KOH + H2SO4 1=> H20 \ + K2SO4.
With any other pair consisting of an acid and a base, a similar interac-
tion occurs (cf. p. 266), water and a salt being produced.
A double decomposition between ionogens is always reversible
(p. 281), and so we should expect that in dilute solution the interaction
of an acid and a base would be incomplete. We find, however, that this
particular sort of action almost always goes so near to completion
that it can be employed for exact measurement of the quantity of the
acid or base. This kind of action is called neutralization, because both
acid and base are completely consumed, and hydrion and hydroxyl are
alike impossible of detection in the resulting mixture. The solution
is neutral to litmus.
Acidimetry and Alkalimetry. — If the problem is to ascertain
the weight of hydrogen chloride in each liter of a specimen of hydro-
chloric acid, this can be done by neutralizing a measured portion of
this acid with a solution of an alkali of known concentration. The
volume of the latter which is required for the purpose is observed. If
the alkali is sodium hydroxide, the action taking place is :
HC1 + NaOH -> H20 + KaCl.
The volume of acid is measured out into a beaker by means of a
pipette (Fig. 82) of fixed capacity, which is filled "by suction to the
mark on the stem. Suppose the amount to be 25 c.c. The standard
(cf. p. 236) alkali solution is placed in a burette (Fig. 83), which is
filled down to the tip of the nozzle. A few drops of litmus solution
are now added to the acid, and the alkali is allowed to run in slowly.
After a time, the hydroxidion which this introduces will begin to pro-
352
INORGANIC CHEMISTRY
FIG. 82.
duce a blue color close to where the stream enters the liquid. This is
at first dissipated by stirring, and the whole remains red. Finally,
however, a point is reached at which the entire solution assumes
a tint intermediate between blue and red. With one drop less
of the base, it is distinctly red. With one drop more, it
would become distinctly blue. Litmus paper of either shade
dipped in this neutral solution remains unaffected. It is
needless to say that the acid might have been added to the
base with the same final result.
The standard solutions used in this work are usually
normal, and contain one equivalent weight of the alkali or
acid in one liter of the solution. For more delicate work,
deci-normal solutions
may be employed. The
concentration of such
a solution is called its
titer, and the operation
of neutralizing another
solution by means of
it, titration. The value
of standard solutions
lies in the fact that
when once the solution
has been prepared, and
the exact concentration
adjusted by quantita-
tive experiments, its
use does not require
any weighing, and the
measurements of volumes can
be carried out with great rapid-
ity. A process involving weigh-
ing need not again be under-
taken until the stock of the
standard solution is exhausted.
The calculation of the result
is also simple. One liter of
normal alkali contains 17 g. of
available hydroxyl, and one liter of normal acid 1 g. of available
hydrogen. Equal volumes of normal solutions will therefore exactly
FIG. 83.
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 353
neutralize one another, 18 g. of water being formed by interaction of
a liter of each. If, for the neutralization of the 25 c.c. of hydrochlo-
ric acid used above, 50 c.c. of normal alkali are required, the acid is
twice-normal (2N). When 15 c.c. are required, the acid is \% or f N.
If the actual weight of hydrogen chloride in the latter case has to be
calculated, we remember that there are 36.45 g. of the compound in
1 1. of a normal solution, and therefore 36.45 X f X T^§^ g- = 5.466 g.
in 25 c.c. of one which is f normal.
Methods of quantitative analysis in which standard solutions (cf.
pp. 148, 236, 307) are employed are known as volumetric methods, and
are much used by analysts and investigators. They occupy much less
time than gravimetric operations, in which numerous weighings have
to be made, and are often just as accurate. The substances, like
litmus, by whose change of color the completeness of the action is
made known, are called indicators (see below).
Theory of Neutralization. — The neutral mixture of the acid
and base gives no evidence of the presence either of the hydrogen ions
or of the hydroxyl ions. The characteristic tastes, and actions upon
indicators, of these two ions, and the interaction of the former of the
two with metals like magnesium, are all wanting. That this is due,
not simply to two opposing influences having destroyed each other's
effects, but to a real disappearance of the agencies themselves, may be
demonstrated by showing that the total number of ions is very much
smaller in the mixture than in the two substances taken separately.
The trough (Fig. 81, p. 327) is half-filled with a dilute solution (say, N/4)
of some active acid, such as hydrochloric acid. An equal volume of
a N/4 solution of some soluble base, such as sodium hydroxide, is
then allowed to flow in, below the acid. On completing the circuit we
find a considerable deflection of the amperemeter (say, 1.5 amperes).
When the interaction is now brought about by stirring, a very great
fall in the reading (say to 0.5 amperes) is observed.* The only
plausible explanation is that, not only have many of the ions assumed
a molecular form, but those which have suffered in this respect have
been the most rapidly moving and best conducting ones, namely, the
hydrion and hydroxidion.
* The experiment may be made more striking by adding a few drops of phenol-
phthalein solution to the acid and using a minute excess of the base. To prevent
the appearance of a pink layer at the interface, and before the stirring, a thin layer
of sodium chloride solution may be introduced below the acid, before the layer of
the base is added.
354 INORGANIC CHEMISTRY
The general plan of all interactions of acids and bases is as
follows :
HC1 fcj Cl' + H-
NaCl H20
The ionization of the hydrochloric acid reaches 0.785 in a normal
solution, and goes further when the acid is diluted with the water of
another solution. That of the sodium hydroxide similarly goes beyond
0.73. Thus the initial substances are almost entirely ionic. The
crosswise union, H* + OH' t^ H20, however, is all but complete, for
water is hardly ionized at all (p. 331). The materials on whose inter-
action with the Cl' and Na", respectively, the maintenance of mole-
cules HC1 and NaOH depends, being thus removed, the dissociation of
the acid and base promptly brings itself to completion, and the left
sides of the equations vanish. Practically all the hydrion and
hydroxyl become water. The Cl' and Na*, however, if the solution is
now semi-normal, unite to the extent of 0.266 only (p. 331). If
it is more dilute, this union forms a still smaller factor in the whole
change. Practically it is negligible. Now all that has been said of
this acid and base will apply mutatis mutandis whenever any active,
highly ionized acid and base come together. Thus we may write
one simple equation for all neutralizations of active acids and bases :
H- + OH' -» H20,
without omitting anything essential.
The ions of a salt are always left over from the main action, and
may be brought together, in turn, by evaporation (ef. p. 341).
The equations as commonly written :
NaOH + HC1 — » NaCl + H2O,
2NaOH + H2S04 — > Na2SO4 + 2H20,
apply to the interactions when water is absent. If used for neutralization in
dilute solution, it must be understood that they condense two changes into one
equation. The formation of water comes first, that of the salt afterwards. Some-
times neutralization is wholly misconstrued by the supposition being made that it
occurs in consequence of a great tendency to salt formation.
It will be seen that neutralization is the precise reverse of hydroly-
sis (pp. 181, 344). The former being almost always nearly complete,
the latter must be, as a rule, very slight.
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 355
Indicators. — Indicators- are substances which, in presence of
certain other substances, assume a very deep color, or change sharply
from one deep color to another. Thus, phenolphthalein (p. 349) is
colorless in presence of acids (i.e., hydrion), and red (when dilute
pink) in presence of alkalies (i.e., hydroxidion). Litmus, again, is red
with acids, and blue with alkalies. The change of color depends upon
a chemical interaction in each case, but since indicators are chosen for
their strong coloration, the quantity of the acid or base used up in chang-
ing the tint of the trace of the indicator is so small as to be negligible.
The common indicators are :
Phenolphthalein, C14H1004, a colorless substance and very feeble
acid. It is not perceptibly dissociated into its ions :
C14H1004 (colorless) <=> C14H90/ (red) + H',
and in neutral or acid solutions is, therefore, without visible color.
When a base is added gradually to an acid containing some of this
indicator, the acid is first neutralized. Then, and not till then, the
slightest excess of hydroxidion unites with the trace of hydrion from
the phenolphthalein, the above equilibrium is displaced forwards, and
a visible amount of the red negative ion is formed :
C14H1004 (colorless) <=; C14H,0/ (red)* 4- H' )
NaOH t=» Na' + OH' $ ~
This indicator shows the presence of an excess of alkali most sharply
when the alkali is an active one like sodium hydroxide, and should,
therefore, be employed only with strong bases. With a weak base
like ammonium hydroxide, a considerable excess must often be used
before the color appears.
Litmus is a natural dyestuff of unknown chemical structure.
Doubtless, however, one of its colors is that of the molecule, and the
other that of the ion.
Methyl orange, (CH8)2NC6H4.N : N.C6H4S08]S"a, is a complex organic
compound which gives, in acid solution, a red and in alkaline solution
a yellow color.
Congo red is the sodium salt of an acid of complex structure (see
Dyes). In neutral or alkaline solutions it is red; with acids it turns
* The ioii has this composition, but, in reality, has a different chemical struc-
ture from the corresponding part of the original molecule. An internal rearrange-
ment, not representable in the equation, accompanies the dissociation. The same
remark applies to the other indicators,
356 INORGANIC CHEMISTRY
«
blue. Paper dipped in Congo red differs from litmus paper in that it
shows gradations in color, the blue being much more distinct with an
active acid than with a relatively weak one like acetic acid (p. 347).
Litmus paper is equally red with all acids save the very feeblest.
Some special indicators have been mentioned. Thus, starch emul-
sion is used for recognizing the presence of traces of iodine (p. 235).
Potassium permanganate is itself so strongly colored that it is its own
indicator (p. 307).
Neutralization of Little Ionized Substances. — When concen-
trated solutions are employed, or acids and bases which are but little
ionized are involved, the mechanism of the change is still the same in
all respects. The only difference is that, since the acid or base is not
fully ionized to start with, its molecules must dissociate progressively,
in proportion as the hydrogen ions pass into combination. All the
hydrogen and hydroxyl capable of forming ions will pass through that
stage and ultimately become water before the solution can reach the
neutral condition.
From this it will be seen that the activity of acids and bases cannot
be measured by the quantity of base or acid required to neutralize
them. The full amount required by the equation is always needed
in every case. This is because neutralization uses up the hydrion or
hydroxidion at once, and so permits the rapid generation of a fresh
supply. The concentration of one of these ionic materials can only be
measured by some action which uses it up slowly or not at all, so that
ionic double decompositions are excluded. In the action of metals on
acids (p. 347) and in determining conductivity (p. 325) the consumption
of the ions is slow, and hence the measurement can be made in these
cases. Actions which consume no ions at all are also known, and are
used in measuring activity (see Carbohydrates and esters).
When the acid or base is but little soluble in water, as when zinc
hydroxide is treated with a dilute acid, one other link is added to the
network of equilibria. The acid proceeds to interact with the small
dissolved part of the base. As this is disposed of, solution goes on
progressively, and, through a train of equilibria :
Zn(OH)2 (solid) <=» Zn(OH)2 (diss'd) <=> Zn"+ 20H',
the supply of hydroxidion is maintained until all the molecules of the
base, solid and dissolved, are used up and the action is completed.
Heating hastens these, as it does all other changes.
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 357
If acid and base are alike insoluble, it is best, if the production of
'the salt is the ultimate object, to fuse the materials together at a high
temperature.
Thermochemistry of Neutralization. — The above interpreta-
tion of the phenomena of neutralization is confirmed by many facts.
Thus a considerable amount of heat is liberated in neutralization.
Now, when active acids (p. 347) and bases (p. 349) in dilute solution
are concerned, it is found that the quantities of heat for the neutraliza-
tion of the same amount of hydrion, or hydroxidion, are always the
same, namely, 13,700 cal. for equivalent weights. If the action con-
sisted primarily in the formation of a different salt from every pair,
we should expect the heat liberated to be different. Thus, the heats of
formation of dry potassium chloride and dry sodium iodide are 10,120
cal. and 7030 cal., respectively. But the heats of formation of their
solutions by neutralizing the proper acids and bases are identical. If,
however, in such cases, neutralization consists always simply in the
formation of water, we should expect the quantities of heat liberated to
be identical, as, in fact, they are :
H* + OH' -> H2O + 13,700 cal.
We are confirmed in these conclusions when we employ concentrated solutions,
or use less completely ionized, or insoluble acids and bases for neutralization. With
such substances, — and they are in the majority, — the heats of neutralization are not
alike, but different in every case. Thus, for dilute solutions of sodium hydroxide
and hydrofluoric acid, the latter a slightly ionized soluble acid, the thermochemical
equation is as follows :
NaOH + HF -» H2O + NaF + 16,270 cal.
Since the sodium fluoride is fully ionized, the only difference between this case and
the preceding one is that the hydrogen fluoride is largely in the molecular condi-
tion to start with, and that here, in addition to the union of hydrogen and
hydroxyl ions, we have a continuous dissociation of the hydrofluoric acid accom-
panying the neutralization. The fact that here the heat produced is much greater
than before, shows that the dissociation of this acid is associated with the produc-
tion of heat (cf. pp. 260, 330). When the same base is used with hypochlorous acid,
the divergence is in precisely the opposite direction and about the same in amount:
NaOH + HCIO — > H20 + NaCIO + 9840 cal.
Here again the salt produced, sodium hypochlorite, is fully ionized, so that the dim-
inished evolution of heat must be due to the fact that the feebly ionized hypochlo-
rous acid absorbs part of the heat of neutralization in passing into the ionic
condition. Applying this to bases, we find that the neutralization of ammonium
hydroxide, a feebly ionized base, with any active acid, gives a heat of neutraliza-
tion below the normal:
358 INORGANIC CHEM1ST11Y
NH4OH + HC1 -» H2O + NH4C1 + 12,200 cal.
Here again the salt produced is fully ionized. Thus the ionization of the ammo-
nium hydroxide must have consumed an appreciable part of the heat of neutraliza-
tion which would otherwise have reached the normal figure of 13,700 calories.
Volume Change in Neutralization. — When the volumes of the solu-
tions of active acids and alkalies are carefully measured before being mixed, and
compared with the volume of the neutral mixture, a contraction is always found to
have occurred. When one liter of a normal solution of each substance is taken at
starting, the volume of the mixture is always 20 c.c. less than that of the compo-
nent liquids. In other words, 18 g. of the ions of water occupy a volume greater
by 20 c.c. than that of the water they form. When less highly ionized acids and
bases are used, the alteration in volume is irregular, since it is affected by the
occurrence of other changes than the mere union of hydrion and hydroxidiou.
MlXED lONOGENS AND DOUBLE SALTS.
As a rale, a univalent ion, such as chloranion (C103'), unites with
one kind of cation and gives but one kind of salt (cf., however, p. 242).
The resxilt is called a neutral or normal salt, as KC10, or NaClO..
o o
The acid, chloric acid, is called a monobasic acid, for its molecule
reacts with but one molecule of a base. The possibilities are more
numerous, however, with an ion of higher valence. Thus :
CABBONANION MAY GIVE : CALCION MAY GIVE :
H2C08, the acid, Ca(OH)2, the base,
Na^COg, a neutral salt, CaCl2, a neutral salt,
NaHC08, an acid salt, Ca(OH)Cl, a basic salt,*
NaKCOg, a mixed salt, . CaCl(OCl), a mixed salt.
Carbonic acid is a di-basic acid, and calcium hydroxide' a di-acid base.
The last two compounds of each set are mixed ionogens. Their char-
acteristic is that they contain more than two kinds of radicals and
break up in solution, giving more than two kinds of ions.
Acid Salts. — The acid salts are obtained by using half that quan-
tity of the base which would be required fully to neutralize the acid, and
evaporating the resulting solution :
NaOH +H2S04 *=; H20 \ + NaHS04.
With a monobasic acid, say hydrochloric acid, this treatment gives
simply a mixture of the normal salt and the free acid, and not a single,
substance.
* This particular basic salt has not been isolated in a pure state.
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 359
Acid salts are also formed by the interaction of other salts with an
excess of the acid (pp. 178, 179, 227).
The acid salt is intermediate in composition between the acid itself
and the normal salt. All of the hydrogen of the acid has not been
displaced by the metal. It is named an acid salt on account of its
composition, but is not necessarily acid in its reaction towards
litmus. That depends on whether its solution contains a sufficient
amount of hydrion to affect indicators. Sodium hydrogen sulphate
gives the ions Na* and HSO/, but, even in moderately dilute solution,
the latter ion is further dissociated into H* and SO/' to a large
extent (p. 346). Its solution is therefore acid in truth. On the other
hand, sodium hydrogen carbonate, NaHC03, derived from carbonic acid,
H2C08, gives the ions Na * and HC03', and the amount of hydrion
formed by the latter is too small to be detected by indicators. This
acid salt gives therefore a solution which is actually neutral to litmus.*
Basic Salts. — Corresponding to the acid salts we have also basic
salts, about which statements parallel to the above might be made.
Thus, from sodium hydroxide but one salt can be formed. With lead
hydroxide, Pb(OH)0, however, the displacement of one hydroxyl by a
negative radical, without the disturbance of the other, is conceivable
and can be achieved. The half-chloride, for example, is called lead
oxychloride (Pb(OH)Cl). The basic salts are usually insoluble in
water, and therefore as a rule do not exhibit the basic reaction with
litmus.
Mixed Salts. — So-called mixed salts, like sodium-potassium car-
bonate KNaC08, (see Silicates), may be obtained by half neutralizing
the acid with the calculated amount of one base and then completing
the operation with the other. Corresponding treatment will give
mixed salts of a di-acid base.
It will be seen that 2KNaC08 is equivalent to K2C03,Na,C08, and
that 2CaCl(OCl) is equivalent to CaCl2,Ca(OCl)2. Since we have as
yet no general means of determining the molecular weights of solids,
there is no generally applicable way of deciding which formula is
preferable (see, however, Bleaching powder, p. 266, and under Calcium).
In solutions of these salts the ions which are found might come from a
* Because of hydrolysis (p. 344), the solution of the "neutral" salt, sodium
carbonate Na2CO3, is actually alkaline in reaction. The terms " acid," "basic," and
" neutral," applied to salts, refer simply to the composition and ignore the behavior.
360 INORGANIC CHEMISTRY
substance possessing either of the alternative formulae, so that no light
is thrown on the question by this means. Thus, most compounds of this
kind, with the exception of acid and basic salts,* are considered to be
molecular compounds (p. 123) of two salts and are classed as double salts.
Double Salts. — Substances similar to ferrous-ammonium sulphate
FeSO^NH^SO^GH^O, are very numerous. Because their formulae
can be written so as to show two complete salts, and because they are
easily formed by crystallization from a solution containing both salts,
they are called double salts. In solution they are resolved into their
constituent salts, and these, in turn, are ionized. Almost always the
acid radicals are identical (see, however, Kainite).
Each kind of ion of a double salt exhibits its own properties, irre-
spective of the nature of the numerous substances, ionic and otherwise,
which are present. Hence, when a solution of a particular ionic
material is required, solutions of such bodies are often used instead of
those of simpler ones, if for any reason the substitution is convenient.
The choice of the complex compound must be made in such a way
that the other ions shall not interfere with that particular reaction of
one of them which is in question.
The class of bodies known as salts of complex acids (q.v.) are
ionized like ordinary salts and not like double salts.
KINDS OF IONIC CHEMICAL CHANGE.
Five distinct varieties of chemical change are characteristic
of ionic materials. These are : (1) Disunion or combination of ions,
(2) displacement of the material of one ion by another substance, (3)
destruction or formation of a compound ionic material, (4) change in
the charges of two ionic materials, (5) charge or discharge of two ionic
materials, electrically. Every one of these kinds of action has been
illustrated, some of them very frequently, in the present and foregoing
chapters.
Disunion and Combination of Ions. — This sort of change is
illustrated every time an ionogen is dissolved in water (disunion) or a
solution of such a substance is evaporated (combination). Both of the
* The formulae of basic salts even are often written as if they were molecular
compounds, as Cu(OH)2,CuCl2, or even CuO,CuCl2,H2O, in place of Cu(OH)Cl
(see Copper).
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 361
directions of this sort of change occur also to a greater or less extent
whenever solutions of two ionogens are mixed. In the latter case :
(1) Two salts give two salts (pp. 337-341).
(2) An acid and salt give a salt and an acid (p. 349).
(3) A base and salt give a salt and a base (p. 350).
(4) An acid and base give water and a salt (neutralization).
(1) is complete only when at least one product is insoluble. (2)
and (3) are complete when at least one product is little ionized or
insoluble or both. (4) is almost always complete because water is
generally less ionized than any other substance in the system.
Displacement of One Ion : Electromotive Series of the Metals.
— When zinc, aluminium, and other metals are placed in a dilute acid,
hydrogen is liberated (pp. 95 and 346) :
Zn -f 2H* -» Zn" + H2.
This action takes place with all acids, because it concerns in reality
only the hydrion in the solution. Its speed for any one metal depends
on the concentration of hydrion (p. 347). Similarly, fluorine displaces
chlorine from chloridion (p. 342), chlorine displaces bromine from
bromidion (p. 228) :
Cl2+2Br'^2Cl'+Br2,
and bromine displaces iodine from iodidion (p. 236). Each of these
actions is independent of the nature of the other ion which accompa-
nies the one undergoing change.
The same sort of displacement occurs with all positive ions. Thus,
zinc will not only displace hydrogen, but also other metallic elements,
like iron, lead, copper, and silver, from the ionic condition in solutions of
their salts. Lead, in turn, will displace copper and silver, but not zinc or
iron. Copper will displace silver. Thus the metals can be set down
in an order such that each metal displaces those following it in the
list and is displaced by those preceding it. This list (see next page)
is known as the electromotive series of the metals, because in electro-
lysis of normal solutions of their salts, the electromotive force of
the current required to deposit each metal (cf. p. 324) is less than
that for the metal preceding in the list (see Electromotive chemis-
try).
This list embodies many facts in the behavior of the metals, and
362
INORGANIC CHEMISTRY
should be kept in mind as furnishing a key to the actions in which a
free metal is used or produced. For example, the chemical activity
of the free metals places them in the same order. The earliest ones
rust much more readily in air than do the later ones. Those following
copper do not rust. Conversely, the oxides of the metals down to
and including manganese, when heated in a
stream of hydrogen, may give lower oxides, but
are not completely reduced. The oxides of cad-
mium and succeeding metals are easily reduced.
The oxides of mercury and the last four metals
are decomposed by heating alone. The relations
of the metals in respect to combination with
elements other than oxygen are similarly ex-
pressed by the arrangement in this table.
The position of hydrogen is particularly sig-
nificant. It will be noted that none of the metals
preceding hydrogen are found free in nature as
ordinary minerals,* while all of the metals suc-
ceeding hydrogen, although occurring to some
extent in combination, are found also free. The
explanation of this is that, by prolonged action
upon ordinary water, containing, as it must, car-
bonic acid and other sources of hydrogen ions,
the metals preceding hydrogen must eventually
displace hydrion and pass into some form of
combination (cf. p. 346). The metals following
hydrogen do not displace hydrion and are much
less affected by the agencies which are most
active in the chemical transformation of minerals.
Hence they often remain in the free state.
The negative ions can be arranged in order in a similar way.f
ELECTROMOTIVE
SERIES OF THE
METALS.
Alkali metals (q.v.)
Alkaline earth
metals (q.v.)
Magnesium
Aluminium
Manganese
Zinc
Chromium
Cadmium
Iron
Cobalt
Nickel
Tin .
Lead
Hydrogen
Arsenic
Copper
Antimony
Bismuth
Mercury
Silver
Palladium
Platinum
Gold
* Free lead and tin do occur as rare minerals. Iron, with a little cobalt and
nickel, constitutes many meteoric masses.
t To avoid a common misconception, it must be noted that the electromotive
series has no bearing on the tendency of one radical to dislodge another in double
decompositions. The place of an element in the E.M. series defines its relative
activity when free, and has to do only with actions where one free element dis-
places (p. 99) another. The influences which determine a double decomposition
(cf. pp. 180, 257-260) have little to do with the chemical activity of the com-
pounds concerned (p. 350), and nothing at all to do with that of the free elements,
for these, in fact, are not present at all (see Halides of silver).
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 363
Destruction or Formation of a Compound Ion. — The de-
struction of a compound ionic material is observed in the action of any
reducing agent, such as hydrogen peroxide (p. 307), upon a dilute
solution of a permanganate. The compound ion MnO/ gives by
reduction Mn" and water. It was also encountered (p. 278) in the
reduction of I03' by hydriodic acid, in which free iodine and water
were formed:
H- + IO/ 4- 5H* + 51' -> 3H,0 4- 3I2.
The converse occurs when potassium cyanide is added in excess to
a solution of a salt of silver. First, silver cyanide is precipitated, and
then this compound unites with the excess of cyanidion :
Ag' + NO/ + K- + NO' <=> K' + N0'3 + AgNC J,
K* 4- CN'+ AgCNf-^ K- + Ag(CN)/.
The product is the soluble potassium argenticyanide. It is a salt of
the complex acid HAg(CN)2, and not a double salt (p. 360). It does
not decompose into potassium and silver cyanides and their ions when
in solution, for the second action, above, is not appreciably reversible.
Change in the Charges of Two Ions. — A reduction in the
charge in two ions probably occurs in the preparation of chlorine
(p. 171). The decomposition of the manganese tetrachloride, if it is
indeed formed, takes place by the simultaneous discharge of two
equivalents of electricity from the quadrivalent manganese ion and
two ions of chloridion :
Mn"" + 4C1' -» Mn" + 2C1' 4- C12.
Both this sort of change and its converse are common with ions of
metals such as iron and tin (q.v.), which have more than one valence.
Charge and Discharge of Two Ions, Electrically. — Dis-
charge of ions is brought about in every electrolysis (p. 318). Thus,
when hydrochloric acid is decomposed by the current, we have :
2H -f 2 0 -> H3 and 2C1 4- 2 0^ CL,.
The converse takes place when the polarization current (p. 324) is
allowed to flow. Both charge and discharge occur in every simple
battery, as when zinc dissolves in dilute sulphuric acid to give zinc
sulphate (p. 346) :
Zn+20->Zn and 2H420-+H,.
364 INORGANIC CHEMISTRY
The creation of the positive charges in the former of the two
equations leaves the rod of zinc negatively charged, that of the nega-
tive charges in the latter renders the platinum wire positive.
Two or More Kinds of Ionic Action Simultaneously. — In
the interaction of iodic acid and hydrogen iodide, both the decomposi-
tion of a compound ion (I03') and the change in charge on other ions
(I' giving I2, and H* giving water) occur simultaneously. Compli-
cations of this kind are not uncommon.
NON-IONIC MODES OF FORMING IONOGENS.
While ionogens may always be made by the union of the proper
ions, they must nevertheless, in the absence of the solvent, be regarded
as chemical substances which may be constructed out of their constit-
uents without reference to the ionic plane of cleavage. Thus we have
incidentally observed many ways in which acids, bases, and salts may
be prepared that do not involve a union of the constituent ions and
are probably not ionic.
Acids and Bases. — Oxygen acids can almost all be prepared
from the anhydrides, which are not ionogens, and water. Phosphoric
acid, sulphurous acid (p. 71), hypochlorous acid (p. 268), and many
other acids are so formed. The hydrogen halides are all producible by
union of the constituent elements. Many acids are formed from others
when the latter are heated ; for example, chloric acid from hypochlo-
rous acid and phosphoric from phosphorous acid (q.v.}.
Bases are formed by the union of oxides of metals with water
(p. 119).
Salts. — The dry ways of forming salts are very numerous. Thus,
many are formed by direct imion of the elements, as in the case of
chlorides (p. 175), sulphides (p. 12), and other simple salts. Many are
made by reduction or oxidation from other salts, as potassium chloride
from potassium chlorate (p. 65), or potassium perchlorate from the latter
(p. 275). Often a reducing or an oxidizing agent is used, as in making
potassium nitrite (g.v.) from the nitrate, and lead sulphate from lead
sulphide (q.v.}. Almost all oxygen salts can be obtained by the union
of two oxides, as calcium carbonate (q.v.) from calcium oxide and
carbon dioxide. Ammonium salts are formed by combination of am-
monia, which is not an ionogen, with acids (p. 183).
THE CHEMICAL BEHAVIOR OF IONIC SUBSTANCES 365
In manufacturing salts, methods like the above, as well as those
involving ionic actions, are very commonly used. In each case the
cheapest and most easily accessible materials are chosen, and the least
expensive operation is selected.
Neutralization is theoretically the simplest ionic way of getting a salt, because
the water can be removed by mere evaporation. Yet most of the salts which are
on the market are made by the use of other actions. In fact, the pure bases and
acids are usually too expensive to be utilized as sources of salts.
The commonest definition of a salt, as a substance formed by the neutraliza-
tion of an acid by a base, is open to many objections. It is logically defective
because it does not describe what a salt is, but one method of making a salt,
which is an entirely different matter. It is unfortunate in its choice amongst pos-
sible paralogisms, because neutralization is more significant as a method of forming
water than as a means of preparing a salt. And finally, as we have just seen, the
definition has not even the excuse of practical value, for most salts are manu-
factured by entirely different reactions.
Exercises. — 1. Using the data in regard to ionization (p. 330),
formulate other dissociations according to the models on p. 335.
2. Explain fully (cf. p. 297) the effect of potassium bromide upon
the ionization of cupric bromide (p. 335).
3. Give a list of all the colorless ionic substances you can think of
(p. 337).
'4. Formulate fully, according to the diagram on p. 340, the precip-
itation of barium sulphate (p. 342), of silicic acid from sodium silicate
(p. 350), of zinc hydroxide from zinc sulphate (p. 350), of silver chlo-
ride from silver sulphate, and the liberation of hydrogen chloride by
phosphoric acid (p. 179).
5. Give a list of the specific physical and chemical properties
(p. 341) of iodidion and of hydrion.
6. Formulate (p. 361) the displacement of iodine by bromine
(p. 236), and of bromine by chlorine (p. 232).
7. Explain the acid reaction of ferric chloride (FeCl3) solution
(p. 344).
8. Name all the physical components in aqueous solutions of po-
tassium hydroxide, hydrogen chloride, and sulphuric acid (cf. p. 322).
9. Name the anions and cations whose formulae are used on p. 337.
10. Formulate (p. 361) the actions of iron and of aluminium on
dilute hydrochloric acid.
11. What is the molar concentration (p. 250) of hydrion in N/10
hydrogen sulphide (p. 330) and in N/10 acetic acid, of natrion in
N/2 sodium chloride, and of cuprion in N cupric nitrate ?
366 INORGANIC CHEMISTRY
12. Combining the models on pp. 349, 354, and 356, formulate the
action of hydrochloric acid on magnesium hydroxide and on zinc
hydroxide.
13. Formulate (p. 338) and discuss the action of sulphuric acid
upon potassium permanganate (p. 306) and upon potassium dichro-
mate (p. 306).
14. Formulate (p. 354) the neutralizations mentioned on p. 357.
15. What do we infer (p. 359) from the fact that the solution of
sodium hydrogen sulphide (NaHS) is neutral ?
16. Can you invent an interaction of two soluble salts in which
both products shall be insoluble?
17. To which classes of ionic actions do those of iodine on hy-
drogen sulphide (p. 238), and of magnesium on cold water (p. 97),
belong ? Formulate the former according to the model on p. 340.
18. What metals, beside platinum, would be most likely to form
suitable electrodes for an electrolytic cell (p. 324) ?
19. How should you attempt to obtain (p. 363) a pure aqueous
solution of the acid HAg(CN)2?
20. Formulate (p. 363) the electrolysis of hydriodic acid and that
of cupric sulphate, the latter between copper electrodes (p. 325).
21. Give, for each of the following, two definitions, one in terms of
experimental facts, the other in terms of the ionic hypothesis : acid,
base, salt, neutralization, acid salt, mixed salt.
CHAPTER XXI
SULPHUR AND HYDROGEN SULPHIDE
Occurrence. — Free sulphur is found in volcanic regions, where it
is mixed with gypsum and other minerals and occupies the pores of
pumice-stone. Rocky materials accompanying a mineral in this way
are called the matrix. Here and there non- volcanic deposits, formed by
the action of bacteria, have been met with, as in Louisiana and in
Germany. There are many minerals, compounds containing sulphur,
which are chiefly important on account of their other constituents.
Sulphides of metals, such as pyrite (FeS2), copper pyrites (CuFeS2),
galena (PbS), zinc-blende (ZnS), and sulphates, like gypsum (CaS04,
2H20), barite (BaS04), and celestite (SrS04), are fairly plentiful. Sul-
phur is a constituent of albumin and other substances found in the
animal body.
Mamifacture. — Most sulphur is obtained by the simple process of
melting it away from the accompanying volcanic rock at a low tempera-
ture. The liquid sulphur is allowed to run into wooden molds, in
which it solidifies in the form of roll sulphur. For many purposes this
sulphur is sufficiently pure. To produce the best quality it is sub-
jected to distillation from earthenware retorts. The vapor passes into
a large brick chamber and condenses upon the walls and floor in the
form of a fine powder, sold as flowers of sulphur. When the chamber
has become heated, the sulphur condenses in the form of a liquid,
which is drawn off and cast in molds as before.
The greater part of the sulphur of commerce comes from Sicily,
where, in 1898, 447,000 tons were manufactured against 41,000 tons
elsewhere. A certain amount is obtained in Japan. It is found in
California, Nevada, and Louisiana. Within the last two years the
product from the last-named source has superseded Sicilian sulphur in
the American market. Sulphur is popularly known as brimstone.
Some sulphur is also made from poly sulphides (q.v.') (which give
precipitated sulphur), from the waste products of alkali manufacture
by Chance's process (q.v.\ and from the exhaiisted material used in
removing sulphur from illuminating-gas during its purification.
307
368 INORGANIC CHEMISTRY
Physical Properties. — The chief physical peculiarity of sulphur
is that, instead of existing in three physical states only, like water, it
possesses two familiar and perfectly distinct solid forms and two dif-
ferent liquid states of aggregation.
1. Native sulphur is yellow, has a sp. gr. 2.06 and melts at 114.5°.
It is almost insoluble in water, but dissolves freely in carbon disul-
phide (40 parts in 100 at 18°) and in sulphur monochloride (y.v.\
When good crystals are found, they belong to the rhombic system
(Fig. 1, p. 11). The solid sulphur obtained by evaporating a solu-
tion shows, as a rule, more perfect crystalline forms than does native
sulphur, but in all other respects is identical with it. Roll sulphur is
the same substance as these two, although the crystals in their growth
have interfered with one another, and the mass is crystalline, simply,
and not -well crystallized. This variety is called, from its form,
rhombic sulphur.
2. When a large mass of melted sulphur solidifies slowly, and the
crust is pierced and the remaining liquid poured out before the whole
has become solid, the interior is found to be lined with long needles.
This kind of sulphur is nearly colorless and has a sp. gr. 1.96, melts
at 119°, and is in all physical respects a different individual from
rhombic sulphur. The crystals are long, transparent prisms, almost
rectangular in section, and usually showing one large oblique face at
the end. This variety of the simple substance is named, from the
system to which its crystals belong, monoclinic sulphur.
Monoclinic sulphur can be kept only above 96° and below its melting-
point (119°). Every recently solidified mass of sulphur is composed of
it. But, below 96° the mass gradually becomes opaque, the change
usually spreading from one or two points and finally affecting the
whole mass. The opacity is due to the fact that the material has
turned into an aggregate of small particles of rhombic sulphur, each
of which occupies less space than the monoclinic sulphur from which
it was formed. Conversely, rhombic sulphur can be kept only below
96°. When heated above this temperature, but not as high as its
melting-point, it turns slowly into monoclinic sulphur. Contact with
a piece of monoclinic sulphur, or mere rubbing with a hard body, will
determine the point, at which the transformation shall begin, and the
expansion which accompanies this results in a spreading opacity as
before. The delay before the change starts and the effect of rubbing
and inoculation are familiar in connection with almost all changes of
state (cf. p. 159).
SULPHUR AND HYDROGEN SULPHIDE 369
Transitions, marked by definite points, like this one at 96°, are
attended by similar phenomena, whether they lie between two solid
states, or a liquid and a solid state (ice and water), or a gaseous and a
liquid state (steam and water). Heat is given out when we pass in one
direction and absorbed by passage in the other. The rate of change
of vapor pressure with change in temperature (cf. p. 163) is different
on each side of the transition point. A body which has two solid states,
and, therefore, two crystalline forms, is said to be dimorphous (two-
formed), and one with more than two such states polymorphous (see
Ammonium nitrate). But this term is not intended to imply that the
relation of two solid states to each other is essentially different
from that of two states of different kinds, such as solid and liquid,
although the term " polymorphous " is not applied to the latter.
3. When melted sulphur is heated, it undergoes another change at
160°. The formerly pale-yellow, mobile liquid (SA) suddenly becomes
dark in color and so viscous (S^) that the vessel may be inverted
without loss of material. Beyond 260° the viscidity becomes noticeably
less, and at 448.5° the liquid boils and passes into sulphur vapor.
The sulphur thermometer thus shows many more fixed points than
does the water thermometer. The latter consists of ice up to 0°, of
water from 0°-100°, and of steam above 100°. The former includes
rhombic sulphur up to 96°, monoclinic sulphur 96°-119°, mobile liquid
(SA) 119°-160°, viscous liquid (SM) 160°-448.5°, and vapor beyond.
Insoluble, Amorphous Sulphur. — When sulphur which has
been exposed to the air, and particularly sulphur to which a trace
of iodine has been added, is boiled and then allowed slowly to cool,
the product is crystalline and soluble in carbon disulphide, as
before. But when such impure sulphur is boiled and then suddenly
chilled by pouring into cold water, it is at first semi-fluid. After
several days this elastic sulphur, as it is called, becomes hard. It is
then found to contain rhombic sulphur mixed with a large proportion
of another variety of free sulphur. This is almost insoluble in any
solvent, and so may be secured by washing the mixture with carbon
disulphide.
This insoluble sulphur, being without crystalline structure, is called
also amorphous (Gk. d priv., p.op(f>rj form) sulphur. Now amorphous
bodies (see Glass) are always supercooled liquids, that is, liquids still
existing as such at a temperature at which the solid, crystalline form
is the stable one. They have been brought, by cooling, so rapidly
370 INORGANIC CHEMISTRY
through their freezing-point (in general, transition point from one
physical state to another) that crystallization has not had time to
begin (cf. p. 144) and a general rigidity only has supervened. Now
amorphous sulphur is viscous, liquid sulphur (SM) which, by sudden
chilling, has been carried past both the transition to mobile liquid
sulphur (SA) at 160°, and the crystallization as well, without under-
going either of these changes. It is supercooled S^. This accounts
for the fact that it is obtainable only by rapid cooling. Often over
50 per cent of the whole may thus be supercooled. Once the mixture
has been obtained by chilling, the insoluble sulphur reverts very slowly
to the soluble variety, and years are required for the completion of
the reversion at ordinary temperature. At 100° the reversion is com-
pleted in an hour. The capacity of the SM to be supercooled at all
seems to depend on the presence of traces of foreign bodies. Of these,
iodine is the most efficient. The sulphur dioxide or sulphuric acid
produced by prolonged exposure to the air (see below) is the agent
usually responsible for the supercooling. Freshly recrystallized sul-
phur gives no elastic sulphur and no insoluble sulphur.
Insoluble sulphur is found in flowers of sulphur and in sulphur
formed by precipitation from thiosulphates (q.v.) in presence of acids.
Chemical Properties. — When the density of sulphur vapor is
determined at low temperatures and under reduced pressures, the
molecular weight corresponds closely to the formula S8. As the tem-
perature is raised, however, the vapor expands very rapidly, and at
800° the molecular weight is 64.2, and the formula therefore S.,.
Intermediately mixtures of S8, S6, S4, and S2 exist. At higher tempera-
tures no further dissociation seems to take place. The formula of
dissolved sulphur, as measured by freezing- and boiling-point methods
(p. 162), is S8.
We do not ordinarily think of sulphur as a very active chemical
substance, but this is largely due to the fact that its solid condition
interferes with the attainment of close contact with the body upon
which it acts. When finely divided metals, with the exception of gold
and platinum (cf. p. 362), are rubbed together with powdered sul-
phur, union takes place and sulphides are produced. We have seen
that sulphur when heated combines with great vigor with iron and
copper, as it does indeed with most of the metals. Sulphur unites
also with many of the non-metals. Thus with oxygen it produces
sulphur dioxide, and even sulphur trioxide (S03). It unites also with
SULPHUR AND HYDROGEN SULPHIDE 371
chlorine directly. When sulphur is treated with oxidizing agents in
presence of water, no trace of sulphur dioxide (or sulphurous acid)
is formed; the only product is sulphuric acid. Even free oxygen gas
in the air is capable slowly of oxidizing moist powdered sulphur and
producing sulphuric acid, 2S + 211,0 + 302 -» 2H2S04.*
Uses of Sulphur. — Large quantities of crude sulphur are em-
ployed for making sulphur dioxide, which is used in the manufacture
of sulphuric acid, in bleaching feathers, straw, and wool, and in
making alkali sulphites for employment in the bleaching industry.
The manufacture of carbon disulphide consumes a considerable amount
also. The purified sulphur is employed in the manufacture of gun-
powder, fireworks,- matches, and, by combination with rubber, of
vulcanite. Flowers of sulphur is used in vineyards to destroy fungi,
which it does by virtue of the traces of sulphuric acid it yields by
oxidation.
HYDROGEN SULPHIDE.
This compound is found in some mineral waters, which in conse-
quence are known as sulphur waters. It is produced in the decompo-
sition of animal matter containing sulphur, when air is excluded, and
the distinctive odor of rotten eggs is due in part to its presence.
Preparation. — 1. Hydrogen and sulphur do not unite percep-
tibly in the cold. At 310° almost complete union occurs, but about
168 hours are required for the change.
2. Sulphides of metals, being salts, are acted upon more or less
easily by dilute acids (p. 349), and give hydrogen sulphide. Ferrous
sulphide, the least expensive of those easily affected, is generally used :
FeS + 2HC1 -+ H2S + FeCL,.
For hydrochloric acid we may substitute an aqueous solution of any
active, non-oxidizing acid. A Kipp's apparatus (p. 97) is commonly
employed. Since ferrous sulphide is but slightly soluble in water, the
action proceeds by a rather complex series of equilibria :
FeS (solid) <=? FeS (diss'd) ±=> Fe" + S" \
* 2C1' + 2H' ^ H«b (dlss d> ^ H«
* The paragraph on the chemical relations of the element (see end of this
chapter) should be read at this point.
372 INORGANIC CHEMISTRY
The dissolved hydrogen sulphide is very feebly ionized, and main-
tains a smaller concentration of sulphidion (S") than does ferrous
sulphide, in spite of the comparative insolubility of the latter. Hence,
the S" formed from the FeS is continuously removed by union with
the hydrion furnished by the acid, S"+ 2H* t^ H2S, and all the other
equilibria are constantly displaced forwards on this account. The
action is therefore, in essence, like neutralization (p. 354). It will be
observed that the action takes place rather on account of the feeble
ionization of the weak acid than by reason of the activity of the other
acid. We should therefore prefer to say that the weak acid with-
draws, and not, as is sometimes done, that the strong acid drives it
out.
Since the action is an ionic one, the acids must be employed in
dilute form. This is true especially of oxygen acids. Thus, concen-
trated sulphuric acid has little action upon ferrous sulphide in the
cold, and when the substances are heated the oxygen of the sulphuric
acid comes into play, and sulphur dioxide (q.v.~) and free sulphur are
formed.
3. Hydrogen sulphide is- the invariable product of the extreme re-
duction of any sulphur compound. Thus, it is formed by the action
of hydrogen iodide upon concentrated sulphuric acid (p. 237). Even
sulphur itself is reduced by dry, gaseous hydrogen iodide :
2HI + S -> H2S + I2.
The action appears to be just the reverse of that which takes place in
aqueous solution (p. 238), but in reality is quite different. Iodine and
gaseous hydrogen sulphide will not produce free sulphur and gaseous
hydrogen iodide, for this action would involve a considerable increase
in energy in the system. But, in water, they do give hydrion and
iodidion, for these bodies contain very much less energy than does
hydrogen iodide :
2H' + S"+I2^2H' + SJ+2I' or S"+ I2 -> S| + 21'.
Physical Properties. — Hydrogen sulphide is a colorless gas
with a characteristic odor. When liquefied, it boils at — 60.4° (755
mm.), and in solid form melts at — 82.9°. At 12° the liquid exerts a
pressure of 15 atmospheres. The solubility in water at 10° is 360 vol-
umes in 100, and becomes less as the temperature is raised. The gas
can be driven out completely by boiling the solution (cf. p. 182). The
SULPHUE AND HYDROGEN SULPHIDE 373
gas is very poisonous, one part in two hundred being fatal to mammals,
while it is stated that one part in fifteen hundred produces death
in birds.
Chemical Properties of the Gas. — When heated, the gas dis-
sociates :
H2S <=> H2 + S.
At 310° the decomposition, although very slow, affects a small but
perceptible proportion of the gas before coming to rest. The dissoci-
ation, like all thermal dissociations, is accompanied by an absorption of
heat and is therefore greater at higher temperatures (cf, p. 260).
Hydrogen sulphide forms a solid hydrate with water, HSS, 7H2O, the behavior
of which resembles that of chlorine hydrate (p. 174). The vapor tension of water
and the gas amount, together, to 1 atmosphere at 5°. The compound, therefore,
decomposes rapidly at 5°, and more slowly at lower temperatures.
The gas burns in air, forming steam and sulphur dioxide. The
temperature of the mantle of flame surrounding the gas, as it issues
from a jet, being far above 310°, the gas in the interior is dissociated
before it meets with any oxygen. Hence a cold dish held across the
flame receives a deposit of free sulphur, and a part of the hydrogen
also escapes unburnt. It may be remarked that dissociation of this
kind probably precedes the combustion of most gaseous compounds
(see Flame).
The metals, down to and including silver in the electromotive
series, when exposed to the gas, quickly receive a coating of sulphide.
That the gas should thus behave like free sulphur shows its instability.
This instability is shown also in the fact that its hydrogen reduces
substances, such as sulphur dioxide, which are not affected by free
hydrogen :
2H2S + S02 -> 2H20 + 3S.
This action takes place much more rapidly when the gases are moist
than when they are dry, and is retarded by dilution with indifferent
gases (cf. p. 252). Native sulphur is probably produced by this action,
as both of these gases are found issuing from the ground in volcanic
neighborhoods. Sulphur is deposited also when hydrogen sulphide
undergoes a partial combustion with a restricted supply of oxygen :
2H2S -f- 02 -> 2H,0 + 2S.
Its formation in nature is sometimes to be accounted for in this way.
374 INORGANIC CHEMISTRY
When hydrogen sulphide gas is led through concentrated, or even,
simply, normal sulphuric acid, the acid is reduced, sulphur dioxide
escapes, and sulphur is deposited :
H2S + H2S04 -> S + 2H20 + S02.
The sulphuric acid may be written H20, S03. In furnishing SO2, there-
fore, each molecule can give one unit of oxygen and therefore oxidize
one molecule of H2S (see p. 389). On account of this action the gas
cannot be dried by means of concentrated sulphuric acid. Calcium
chloride is likewise inapplicable, since a partial interchange takes
place, resulting in the production of calcium sulphide and hydrogen
chloride gas. Only a dehydrating agent, such as phosphoric anhy-
dride, with which it cannot interact, is suitable for drying the gas.
Chemical Properties of the Aqueous Solution of Hydrogen
Sulphide. — While the gas itself is not an acid, its solution in water
gives a feeble acid reaction with litmus. In an aqueous solution
(N/10), only .0007 (0.07 per cent) of the substance is ionized accord-
ing to the equation (cf. p. 346) :
H2S -> H- + HS'.
Some S" ions are present, but hydrosulphidion (HS') is less dissoci-
ated than is water itself, and the amount of sulphidion is therefore
very small.
By the action of oxygen from the air upon an aqueous solution of
hydrogen sulphide, the sulphur is slowly displaced and appears in the
form of a fine white powder :
0, + 2H2S -> 28 + 2H2O.
This is an action similar to the displacement of ionic iodine by free
chlorine (p. 236). On the other hand, the hydrogen may be displaced
by metals, particularly the more active ones, but the small degree of
ionization makes the action very slow.
The solution of the gas is a reducing agent, as its action upon
iodine shows (p. 238). So, also, in presence of an acid, it removes
oxygen from dichromic acid (produced by the action of an acid upon
potassium dichromate) :
K2Cr207 + 2HC1 <=> H2Cr207 + 2KC1 (1)
H2Cr.A + 6HC1 -» 4H20 + 2CrCls(+ 30) (2)
(30)' + 3H2S -^ 3H2O + 3S _ (3)
_
Adding: K2Cr2O7 + 8HCl + 3H2S -> 2KC1 + 2CrCl?+7H2O + 38
SULPHUR AND HYDROGEN SULPHIDE 375
The first partial equation (cf. p. 228) represents the regular interaction
of two ionogens, but the second interaction does not take place unless an
oxidizable body (here the hydrogen sulphide) is present to take posses-
sion of the oxygen which it is capable of delivering (cf. p. 307).
As an acid, the solution of hydrogen sulphide, sometimes known as
sulphydric acid, may be neutralized by bases.
Sulphides. — As a di-basic acid (p. 358), hydrogen sulphide gives
both acid and neutral (or normal) sulphides, such as NaHS and Na2S.
The former are obtained by passing the gas in excess into solutions
of soluble bases :
H2S + NaOH -4 H20 + NaHS,
and are neutral in reaction. Their negative ion, HS', gives practically
no hydrion.
By adding to the above solution an amount of sodium hydroxide
equal to that used before, and driving off the water by evaporation,
the second unit of hydrogen is displaced :
NaOH + NaHS s=> ]Sa2S + H20 \ .
This action is wholly reversed when dry sodium sulphide is dissolved
in water, the salt being completely hydrolyzed to the acid salt :
S
H0 «
The HS' gives a lower concentration of hydrion than the water, and
hence uses up in its formation the ions of hydrogen produced by the
latter until an amount of hydroxyl equivalent to half the sodium is
formed.
Many sulphides are insoluble in water, and these may be divided
roughly into three classes :
1. The sulphides of silver, copper, mercury, and some other
metals are exceedingly insoluble, and, therefore, do not interact with
dilute acids as does ferrous sulphide (p. 371). These may therefore
be made by leading hydrogen sulphide into a solution of any of their
salts :
±=5 CuSJ + H,SO4.
The acid produced has scarcely any effect upon the sulphide, and al-
most no reverse action is observed.
376 INORGANIC CHEMISTRY
2. The sulphides of iron, zinc, and certain other metals are insol-
uble in water, but not so much so as the last class. Hence they are
decomposed by dilute acids, and the reverse of the above action takes
place almost completely. These sulphides must therefore be made,
either by combination of the elements, or by adding a soluble sulphide
to a solution of a salt :
FeS04 + (NH4)2S ±5 FeS J + (NH4)2S04.
No acid is produced in this sort of interaction, and the considerable
insolubility of the sulphide of iron or zinc in water renders the
change nearly complete.
3. The sulphides of barium, calcium, and some other metals (q.v.),
although insoluble in water, are hydrolyzed by it, and give soluble
products (see p. 604) :
2CaS'+ 2H2O <=> Ca(OH)2 + Ca(SH)2.
They may be prepared by direct union of the elements, and from the
sulphates by reduction with carbon. (For the application of the above
differences in solubility amongst sulphides in chemical analysis, see
end of Chap, xxxvii.)
The soluble acid sulphides are oxidized in aqueous solution by at-
mospheric oxygen :
2NaSH + 02 -> 2NaOH + 2S.
The sulphur is not precipitated, but combines with the excess of the
sulphide, forming polysulphides (see below).
Poly sulphides. — When sulphur is shaken with a solution of a
soluble sulphide or acid sulphide, it dissolves, and evaporation of the
solution leaves substances varying in composition from Na^ to Na^j.
Whether those containing less than five units of sulphur are mixtures
of the pentasulphide with the ordinary sulphide (Na2S) in different
proportions, or are single sxibstances, has not been determined.
When an acid is poured into sodium pentasulphide solution, minute
spherules of rhombic sulphur are precipitated :
2HC1 -» 2NaCl + H2S f
When the order is reversed, sodium pentasulphide being thrown
into concentrated hydrochloric acid, no hydrogen sulphide is evolved.
Hydrogen pentasulphide, H2S5, a yellow oil, falls to the bottom of the
vessel.
SULPHUR AND HYDROGEN SULPHIDE 377
When purified, the compound is nearly colorless. Sulphur dis-
solves in it freely, giving a yellow sohition. It can be kept without
change when dry, but contact with water causes it to decompose into
hydrogen sulphide and sulphur.
The Chemical Relations of the Element. — In combination with
metals and hydrogen, sulphur is bivalent, forming compounds like
H2S, FeS, CuS, and HgS. In combination with non-metals, however,
the valence is frequently greater, the maximum being seen in sulphur
trioxide, where we must assume that the sulphur is sexivalent. Its
oxides are acid-forming, and it is, therefore, a non-metal.
Sulphur is regarded as resembling oxygen more closely than any
of the other elements we have studied so far. Both unite directly with
most metals and non-metals. In this they are like chlorine. But hy-
drogen chloride is highly ionized by water, while the hydrogen com-
pounds of oxygen and sulphur are feebly ionized. The formulae of the
compounds of oxygen and sulphur with metals are similar, CuO and
CuS, NaOH and NaSH, and so forth, but this is in part due merely to
the fact that both elements are bivalent. The chemical resemblance of
sulphur to selenium and tellurium (q.v.) is much more striking than
its resemblance to oxygen.
Exercises. — 1. The freezing-point of pure sulphur is found to
vary from 119° down to 114°, depending upon the temperature to
which the liquid has been heated and the speed with which it has been
cooled. To what should you suspect this variability to be attributable ?
2. How could the decomposition of hydrogen sulphide at 310° be
rendered (a) more complete, (b) less complete ? Would the percentage
decomposed be affected (a) by reducing the pressure, (b) by mixing
the gas with an indifferent gas ?
3. To what classes of ionic actions (p. 360) do the interactions of
hydrogen sulphide solution and (a) oxygen (p. 374), (b) acidified
potassium dichromate (p. 374), (c) sodium hydroxide (p. 375), (d)
iodine (p. 372), belong ?
4. Why is normal sodium sulphide only half hydrolyzed by water ?
CHAPTER XXII
THE OXIDES AND OXYGEN ACIDS OF SULPHUR
FOUR oxides of sulphur, represented by the formulae S203, S02,
S08, and S2O7, are known. Of these, however, the first and the last
are much less familiar substances than the other two. The dioxide
and trioxide of sulphur are not only important in themselves, but
their relation to the acids H2S03 and H2S04, which may be obtained
from them by the addition of water, makes them doubly so to the
chemist.
Preparation of Sulphur Dioxide. — When sulphur burns in air
or oxygen, sulphur dioxide is produced (p. 67). While some of the
sulphur dioxide used in commerce is prepared in this manner, the
larger part is probably obtained by the roasting of sulphur ores.
Pyrite (FeS2), for example, which is a familiar yellow, metallic-look-
ing mineral, burns when it has been raised to the kindling tempera-
ture, on account of the large amount of sulphur which it contains :
4FeS2 + 1102 -» 2Fe203 + 8S02.
It should be noted in passing, that heating and roasting are dis-
tinct processes in chemistry. The latter term always assumes the
access of the air and employment of its oxygen ; the former, in the
absence of modifying words, assumes the exclusion or the chemical in-
difference of the air.
Sulphur dioxide is also set free by the action of acids upon
sulphites. Sulphuric acid and a strong solution of sodium sulphite
may be used :
S03" J „ Qri •'.
H SO ±? SO " + 2H- j ** 2 8 ^ H2° + s°2 (diss'd) <=» S02 (gas).
The sulphurous acid being only moderately ionized, its molecules are
formed in considerable amount. Being also very unstable, it decom-
poses spontaneously into water and sulphur dioxide, and the latter
escapes when sufficient water for its solution is not present.
378
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 379
In the laboratory, sulphur dioxide is frequently made by the
reduction of concentrated sulphuric acid. Copper is the metal most
commonly employed, because by its means very pure sulphur dioxide
can be obtained. More active metals, such as iron and zinc, although
cheaper, cannot be used, since they reduce the sulphuric acid to
hydrogen sulphide. The undiluted hydrogen sulphate consists en-
tirely of molecules, and, at the high temperatures at which alone the
action is vigorous, is an oxidizing agent (cf. p. 97). A part of it loses
oxygen to form water with the hydrogen of another molecule:
H2S04 -r» H20 + S02 ( + 0) (1)
(0) + H2S04 + Cu -> H2O + CuS04 (2)
2H2SO4 + Cu -> 2H20 + S02 + CuSO4
Some easily oxidized non-metals, such as carbon and sulphur, act in
the same way :
C + 2H2S04 -> 2H20 + 2S02 + C02,
S + 2H2S04 -» 2H20 + 3SO2.
Physical Properties. — Sulphur dioxide is a gas possessing a
penetrating and characteristic odor. This is frequently spoken of as
the "odor of sulphur," but it should be remembered that sulphur
itself has scarcely any smell at all. The weight of the G.M.V. of
the gas (65.55 g.) shows it to be more than twice as heavy as air.
The critical temperature is 156°. By means of a freezing mixture of
ice and salt, the gas is easily condensed to a transparent, mobile fluid,
which boils at — 8°. At 20° the vapor tension of the liquid is 3.25
atmospheres. The liquid may be frozen to a white solid, melting at
— 76D. It ionizes substances dissolved in it as well as does water. The
solubility of the gas in water, 5000 volumes in 100, is very great.
Unlike solutions of the hydrogen halides (p.- 182), however, the liquid
is completely freed from the gas by boiling.
Chemical Properties. — Sulphur dioxide is stable, being decom-
posed only by the use of a very high temperature.
It unites with water to form sulphurous acid, H2S03. Although
the gas itself sometimes receives this name, it is not acid : it is simply
the anhydride (p. 71) of the acid.
Since the maximum valence of sulphur is 6, sulphur dioxide, in
which but four of the valences of sulphur are used, is unsaturated.
It is therefore still able to combine directly with suitable elements,
380 ' INORGANIC CHEMISTRY
such as chlorine and oxygen. When it is mixed with chlorine in sun-
light, a liquid, sulphuryl chloride SC^CLj is produced.
Liquefied sulphur dioxide is now sold in tin cans, and is employed
for bleaching straw, wool, and silk. As a disinfectant it has been
displaced to a large extent by formaldehyde.
Preparation of Sulphur Trioxide. — Although the formation of
sulphur trioxide S03 is accompanied by the liberation of much heat,
sulphur dioxide and oxygen, even when heated together, unite very
slowly. Ozone, however, combines with the former readily.
The interaction of sulphur dioxide and oxygen is hastened by
many substances, such as glass, porcelain, ferric oxide, and, more
especially, finely divided platinum, which remain themselves un-
changed and simply act as catalytic agents. The "contact process,"
as this is called, has been rendered available for the commercial
manufacture of sulphur trioxide by Knietsch (1901). The chief
features of the process are : (1) The complete removal of arsenious
oxide and other impurities derived from the burning of crude pyrite,
the minutest traces of which poison the catalytic agent and soon
render it absolutely inoperative. (2) The preliminary passage of the
cold mixture of gases over the outside of the pipes containing the
contact agent. This removes part of the heat generated by the ac-
tion, S02 + 0 <=± S08 + 22,600 cal., going on inside, and keeps the
temperature of the interior at 400°. Below 400° the union is too
slow ; above 400°, being reversible, it is incomplete. At 400°, 98-99
per cent of the materials unite ; at 700°, only 60 per cent, at 900° none.
Twice the quantity of oxygen theoretically needed is employed. The
vaporous product is condensed by being led into 97-99 per cent
sulphuric acid, and the concentration of the liquid is constantly
maintained at this point by the regulated influx of water. The
trioxide is thus chiefly used for immediate conversion into sulphuric
acid.
Formerly sulphur trioxide was obtained by the distillation of
impure ferric sulphate, Fe2(S04)3 — > Fe208 + 3S08. It may also be
prepared by repeated distillation of concentrated sulphuric acid with
a powerful drying agent, like phosphoric anhydride.
Physical Properties. — Sulphur trioxide S03 is, at ordinary tem-
peratures, fluid. The crystals, obtained by cooling, melt at 14.8°.
The liquid boils at 46°, and is, therefore, exceedingly volatile at ordi-
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 381
nary temperature. It fumes strongly when exposed to the air, in con-
sequence of the union of the vapor with moisture and the production
of minute drops of sulphuric acid.
A white crystalline variety of the substance, which in appearance
closely resembles asbestos, is obtained when a trace of water has
gained access to the oxide. It is a polymer (p. 242) of the above
liquid, having twice its molecular weight, and therefore the formula
(S08)2. When heated at 50° this polymer begins to dissociate and
passes into vapor of S03 without melting. This white solid is the
more stable and more familiar form of the trioxide.
Chemical Properties. — The vapor of sulphur trioxide disso-
ciates above 400° into sulphur dioxide and oxygen.
Sulphur trioxide is not itself an acid, but it is the anhydride of
sulphuric acid. When placed in water it unites vigorously, causing
a hissing noise due to the steam produced by the heat of the union.
In consequence of its great tendency to combine with water, the liquid
variety, which is the more active, removes the elements of this sub-
stance from materials which contain them in the proper proportions.
Thus paper, which is largely cellulose, (CgHjgOg),,., and sugar, C12H22On,
are charred by it, and carbon is set free.
Just as sulphur trioxide unites with water to give hydrogen sul-
phate, so it combines vigorously with many oxides of metals, produ-
cing the corresponding sulphates:
The union of an oxide of a non-metal with the oxide of a metal, in
this fashion, is a general method of obtaining salts (cf. p. 364).
Oxygen Acids of SwlpJiur. — Sulphurous and sulphuric acids
have been mentioned frequently already. Next to them in importance
come thiosulphuric acid and persulphuric acid. The compositions of
the acids show their relationships :
Hyposulphurous acid, H2S204 Thiosulphuric acid, H2S908
Sulphurous acid, H2S03 Persulphuric acid, H2S208.
Sulphuric acid, H2S04
Thiosulphuric acid (Gk. Oeiov, sulphur) is so named because it con-
tains one unit of sulphur in place of one of the units of oxygen of
sulphuric acid. Besides the above we have also : Dithionic acid H2S2Ot,
382 INORGANIC CHEMISTRY
trithionic acid H2S306, tetrathionic acid H2S406, and pentathionic acid
H2S606.
On account of its commercial importance and the interest attach-
ing to its method of manufacture and to its properties, we may first
discuss sulphuric acid. We shall then be able to dispose of the re-
maining acids in a much briefer fashion.
SULPHURIC ACID.
Although salts of sulphuric acid, such as calcium sulphate, are ex-
ceedingly plentiful in nature, the preparation of the acid by chemi-
cal action upon the salts is not practicable. The sulphates, indeed,
interact with all acids, but the actions are reversible. The comple-
tion of the action by the plan used in making hydrogen chloride
(p. 180), involving the removal of the sulphuric acid by distillation,
would be difficult on account of the in volatility of this acid. It boils
at 330°; and active acids, less volatile still, which might be used to
liberate it, do not exist. We are therefore compelled to build up sul-
phuric acid from its elements.
The union of sulphur dioxide and oxygen by the contact process,
and combination of the trioxide with water (p. 380), is the best method
for making a highly concentrated acid. For obtaining ordinary " oil
of vitriol," however, an entirely different plan is still used extensively.
History of Sulphuric Acid Manufacture. — Impure forms of
sulphuric acid have been known for many centuries. In the fifteenth
century it was made by distilling ferrous sulphate with sand. The
product, however, contained much water and sulphur dioxide. The
first successful preparation of the substance commercially was made
by Ward at Richmond-on-the-Thames (1758). The process consisted
in burning a mixture of sulphur and saltpeter (KIS"03) in a ladle
suspended in a large glass globe partially filled with water. The
gases which were evolved contained large quantities of sulphur
dioxide and oxides of nitrogen, which by interaction with atmospheric
oxygen and water (see below) produced the sulphuric acid. The solu-
tion which was obtained, although it could be prepared of any desired
concentration by the burning of a sufficient number of charges, was
far from pure and was expensive, bringing thirteen shillings ($3.25)
per pound. Subsequently a chamber lined with lead was substituted
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 383
for the glass vessel. This reduced the price to about two shillings
and sixpence ($0.60) per pound. The same principles are used in
the modern " chamber process."
Chemistry of the Chamber Process. — The gases, the interac-
tions of which result in the formation of sulphuric acid, are : water
vapor, sulphur dioxide, nitrous anhydride N208 (q-v-), and oxygen.
These are obtained, the first by injection of steam, the second usually
by the burning of pyrite, the third from nitric acid, and the fourth by
the introduction of air. The gases are thoroughly mixed in large
leaden chambers, and the sulphuric acid condenses and collects upon the
floors. In spite of elaborate investigations, instigated by the exten-
sive scale upon which the manufacture is carried on and the immense
financial interests involved, some uncertainty still exists in regard to
the precise nature of the chemical changes which take place. Accord-
ing to Lunge the greater part of the product is formed by two succes-
sive actions, the first of which yields a complex compound that is
decomposed by excess of water in the second :
H20 + 2S02 + N208 + 02 -» 2S02 < ° ~ ^Q (1)
The group — NO is found in many compounds. Here, if it were dis-
placed by hydrogen, sulphuric acid would result. Hence this com-
pound is called nitrosylsulphuric acid:
(2)
The equations (1) and (2) are not partial equations for one inter-
action, but represent distinct actions which can be carried out sepa-
rately. In a properly operating plant, indeed, the nitrosylsulphuric
acid is not observed. But when the supply of water is deficient, white
" chamber crystals," consisting of this substance, collect on the walls.
The explanation of the success of this seemingly roundabout
method of getting sulphuric acid is as follows : The direct union of
sulphur dioxide and water to form sulphurous acid is rapid, but the
action of free oxygen upon the latter, 2H2S08 + 02 — > 2H2S04, is ex-
ceedingly slow. Reaching sulphuric acid by the use of these two
changes, although they constitute a direct route to the result, is not
feasible in practice. On the other hand, both of the above actions,
(1) and (2), happen to be much more speedy, and so, by their use,
384 INORGANIC CHEMISTRY
more rapid production of the desired substance is secured at the ex-
pense of a slight complexity. It may be added that the heat finally
given out in the formation of one formula- weight of sulphuric acid is
exactly the same in amount whether nitrous anhydride intervenes or
not (cf. p. 78).
The progress of the first action is marked by the disappearance of
the brown nitrous anhydride, and, on the introduction of water, the
completion of the second results in the reproduction of the same sub-
stance. It would thus seem as if the nitrous anhydride should take
part an indefinite number of times in these changes and so facilitate
the conversion of an unlimited amount of sulphur dioxide, oxygen,
and water into sulphuric acid, without impairment of its quantity.
In practice, however, certain subsidiary actions take place, such as,
for example, the reduction of some nitrous anhydride to nitrous oxide
(N20), which permanently remove a part of the material from parti-
cipation in the cycle.
The supply of nitrous anhydride is maintained by the introduction
of nitric acid vapor into the chamber. This acid is secured by the
action of concentrated sulphuric acid upon commercial sodium nitrate
NaN08+H2S04 +± HN03| +NaHS04.
On account of the volatility of the nitric acid, a moderate heat is
sufficient to remove it from admixture with the other substances, and
its vapor is swept along with the other gases into the apparatus. The
initial action which the nitric acid undergoes may be represented by
the following equation :
H20 + 2S02 + 2HN03 -> 2H2S04 +
We may write this in the form :
-1- 2S02 + H2O,N205 -> 2H2S04
The two molecules of water, one actually, the other potentially, pres-
ent, with the two molecules of sulphur dioxide, can furnish two
molecules of sulphurous acid (H2S03). The N205 in passing to the
condition N203 gives up the two units of oxygen required to convert
this sulphurous acid into sulphuric acid.
It should be remarked that we shall later find nitrous anhydride
to be an unstable substance, which, in the gaseous form, is composed
largely of the products formed by its dissociation, N2O3 <z± NO2 + NO.
THE OXIDES AND OXYGEN ACIDS OF SULPHUR
385
For the purpose of understanding its present use, however, it is not
necessary to consider this particular fact. The whole behaves as if it
contained nothing but N203.
Details of the Chamber Process. — The sulphur dioxide is
produced in a row of small furnaces A (Fig. 84), the structure of
which depends upon the nature of the substance employed to yield
this fundamental constituent of sulphuric acid. When good pyrite
is used, the ore burns unassisted (p. 42), while impure pyrite and
zinc-blende ZnS have to be heated, to a greater or less degree, artifi-
cially to maintain the combustion. The gases from the various furnaces
pass into one long dust flue, in which they are mingled with the proper
FIG. 84.
proportion of air, and have an opportunity to deposit oxides of iron
and of arsenic and other materials which they transport mechanically.
From this flue they enter the Glover tower G, in which they acquire
the oxides of nitrogen. Having secured all the necessary constituents,
excepting water, and having been reduced very considerably in tem-
perature, the gases next enter the first of the lead chambers. These
are large structures, from three to five in number, constructed com-
pletely of sheet lead. They vary in size, measuring as much as 100
X 40 x 40 feet, and sometimes having a total capacity of 150,000 to
200,000 cubic feet. As the gases drift through these chambers they
386 INORGANIC CHEMISTRY
are thoroughly mixed, and an amount of water considerably in excess
of that actually required, is injected in the form of steam at various
points. The temperature in the first chamber is maintained at 50° to
65°, while in the last chamber it is about 15° above that of the out-
side air. The acid, along with the excess of water, condenses and
collects upon the floor of the chamber, while the unused gases, con-
sisting chiefly of nitrous anhydride and a very large amount of nitro-
gen, derived from the air originally admitted, find an exit into the
Gay-Laissac tower L.
This is a tower about fifty feet in height, filled with tiles, over which
concentrated sulphuric acid continually trickles from a reservoir at
the top. The object of this tower is to catch the nitrous anhydride
and enable it to be reemployed in the process. This is accomplished
by a reversal of action (2) above. The acid which accumulates in the
vessel at the bottom of this tower contains the nitrosylsulphuric acid,
and by means of compressed air is forced through a pipe up to a
vessel at the top of the Glover tower G. When this " nitrous vitriol "
is mixed with dilute sulphuric acid from a neighboring vessel, by
allowing both to flow down into the tower, the nitrous anhydride is
once more set free by the interaction of the water in the dilute acid
(action (2)). The Glover tower is filled with broken quartz or tiles,
and the heated gases from the furnace acquire in it their supply of
nitrous anhydride. Their high temperature causes a considerable
concentration of the diluted sulphuric acid ds it trickles downward.
The acid, after traversing this tower, is sufficiently strong to be used
once more for the absorption of nitrous anhydride. To replace the
nitrous anhydride inevitably lost by reduction to nitrous oxide and
otherwise, fresh nitric acid is furnished by small open vessels N,
containing sodium nitrate and sulphuric acid, placed in the flues of the
pyrite-burners. About 4 kg. of the nitrate are consumed for every
100 kg. of sulphur.
The immense size of the chambers is necessitated by the fact that
the chemical action, although much quicker than the direct oxidation
of sulphurous acid, is after all rather slow. The presence of the large
amount of atmospheric nitrogen, which diminishes the concentration of
all the interacting substances, partly accounts for this slowness. The
acid which accumulates upon the floors contains but 60 to 70 per cent
of sulphuric acid, and has a specific gravity of 1.5-1.62. The excess of
water is used to facilitate the second action. It is required also in
order that the acid upon the floor may not afterwards absorb and
TIIK OXIDES AND OXYGEN ACIDS OF SULPHUR 387
retain the nitrous anhydride, for this substance combines with an acid
containing more than 70 per cent of hydrogen sulphate.
This crude sulphuric acid is applicable directly in some chemical
manufactures, such as the preparation of superphosphates (q.v.). In
most cases, however, a more concentrated sulphuric acid is required.
The concentration is effected in the first place by evaporation in pans
lined with lead, which are frequently placed over the pyrite-burners in
order to economize fuel. The evaporation in lead is carried on until a
specific gravity 1.7, corresponding to 77 per cent concentration, is
reached. Up to this point the sulphate of lead formed by the action
of the sulphuric acid produces a crust which protects the metal from
further action. The insoluble sulphate of lead, however, becomes
more soluble in sulphuric acid the more concentrated it is, and the
higher, therefore, its boiling-point. For some applications of the acid,
concentration beyond this point is not needed. When a stronger acid
is required, the water is usually driven out by heating the sulphuric
acid in vessels of glass or platinum. Commercial sulphuric acid, oil of
vitriol, has a specific gravity 1.83-1.84, and contains about 93.5 per
cent of sulphuric acid.
In modern sulphuric acid plants, the concentration beyond the
specific gravity 1.77 is frequently carried out in vessels of cast iron, as
that metal is attacked to an appreciable extent only by acids of
lower specific gravity.
Physical Properties. — Pure hydrogen sulphate has a sp. gr.
1.85 at 15°. When cooled, it crystallizes (m.-p. 10°). At 150°-180°
the acid begins to fume, giving off sulphur trioxide. At 330° it boils.
The 100 per cent acid gives off sulphur trioxide, and the dilute acid
water, until a mixture of constant composition (98.33 per cent) and
constant b.-p. 330° is formed. This distils with unchanged compo-
sition (r£. p. 182). The vapor, however, contains 34 per cent of the
material in the form of water and sulphur trioxide, which reunite
when cooled.
When hydrogen sulphate is mixed with water a considerable evolu-
tion of heat takes place. This heat receives progressively diminishing
increments as more water is added, until a very great dilution has been
reached. The total is 19,400 cal. This heat of solution has not been
accounted for quantitatively (cf. p. 165), but a part of it is due to the
heatgiven out in connection with the ionization of the hydrogen sulphate.
The solution is thus much more stable (>>., it contains much less energy)
388 INORGANIC CHEMISTRY
than the pure substance, and hence the latter absorbs water greedily.
Many substances containing hydrogen and oxygen are deprived of
equivalent amounts of these elements by sulphuric acid, just as they are
by sulphur trioxide (p. 381). The same tendency is enlisted to promote
chemical actions in which water is formed, particularly in connection
with the manufacture of nitroglycerine (q.v.), gun-cotton (q.v.), and
esters (q.vS).
Impurities. — Commercial sulphuric acid is frequently brown in
color on account of the presence of fragments of straw which have
become charred and finally completely disintegrated. It contains also
lead sulphate, which appears as a precipitate when the acid is diluted,
as well as arsenic trioxide and oxides of nitrogen in combination, and
many other foreign substances in small quantities. The pure sul-
phuric acid employed in chemical laboratories has received special treat-
ment for the removal of these ingredients.
Chemical Properties of Hydrogen Sulphate. — The compound
is not exceedingly stable, for dissociation into water and sulphur triox-
ide begins far below the boiling-point (cf. p. 387), and is practically
complete at 416°, as is shown by the density of the gas. When
raised suddenly to a red heat it is broken up completely into water,
sulphur dioxide, and oxygen.
When mixed with a small quantity of water and cooled strongly it
gives a crystalline monohydrate H2S04, H2O, which melts and is resolved
into its constituents at 8°. Possibly this should be regarded as a differ-
ent sulphuric acid, H4S05 (cf. p. 278).
When sulphur trioxide is dissolved in hydrogen sulphate, disul-
phuric acid H2S207, a solid compound, is obtained. Hydrogen sulphate
containing 80 per cent of disulphuric acid is known as " oleum," and is
employed in chemical industries. The old " fuming," or "Nordhau-
sen " sulphuric acid contained 10-20 per cent of extra sulphur triox-
ide. The salts of disulphuric acid may be made by strongly heating
the acid sulphates, for example :
2NaHS04 <=» Na.SA + H20 |.
In view of this mode of preparation by the aid of heat, they are fre-
quently known as pyrosulphates (Gk. irvp, fire). When they are dis-
solved in water, the acid sulphates are reproduced.
On account of the large quantity of oxygen which hydrogen sul-
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 389
phate contains, and its instability when heated, it behaves as an oxi-
dizing agent. This property has already been illustrated in connection
with the action of the acid upon carbon and sulphur (p. 379) and upon
copper (p. 379), hydrogen sulphide (p. 374), zinc (p. 96), and, particii-
larly, hydrogen iodide (p. 237) and hydrogen bromide (p. 231). The
sulphuric acid is in consequence reduced to sulphur dioxide, and even
to free sulphur or hydrogen sulphide. The metals, from the most
active down to silver (p. 362), are capable of reducing it. Gold and
platinum alone are not attacked, and hence their use in making sul-
phuric acid stills. Free hydrogen itself is oxidized to water when
passed into hydrogen sulphate at 160° :
S02(OH)2 + H2 -» S02 + 2H20.
With salts which it does not oxidize, hydrogen sulphate reacts
by double decomposition and sets free the corresponding acid. The
actions are always reversible ones ; but where the new acid is volatile,
as in the case of hydrogen chloride (p. 180), we are furnished with
one of the cheapest means of preparing acids.
Since hydrogen sulphate is dibasic, that is, since it has two units
of hydrogen which may be replaced by a metal, it forms both acid
and neutral salts.
Chemical Properties of Aqueous Hydrogen Sulphate. — The
solution of sulphuric acid is a mixture whose components are : undisso-
ciated molecules H2S04, hydrion H', hydrosulphanion HSO/, and sul-
phanion S04". The chemical properties shown by the solution are
those of one or other of these components, according to circumstances.
Except in concentrated solutions (normal or stronger) the oxidizing
effects of the undissociated, molecular substance are not encountered.
The temperature of the diluted acid, even when boiling, is not high
enough for the purpose. In fairly strong solutions hydrosulphanion
is plentiful and shows itself in the results of electrolysis (see p. 397).
The presence of hydrion is shown by all its usual properties (p. 345).
In the following table the proportion of the whole of the hydrogen
existing in the form of hydrion (column five) and its concentration
(column six), taking a normal solution of hydrion containing 1 g. per
liter as standard, are shown (cf. pp. 149 and 329). The first three
columns give the concentration of the sulphuric acid as a whole, in
terms (first column) of the volume of liquid containing one equivalent
(^H2S04 = 49 g.), in terms (second column) of a normal solution as.
390
INOKGAMC CHEMISTRY
standard, and by per cent (third column), respectively,
column shows the conductivity (p. 328).
The fourth
V
H2S04
PEK CENT
H2SO4.
A.
**/*,
H'
0.1
ION
38.00
70
0.18
1.8N
1
N
4.79
198
0.51
0.51 N
10
0.1 N
0.48
225
0.58
0.058 N
100
0.01 N
0.05
308
0.79
0.0079 N
1000
0.001 N
0.005
361
0.93
0.00093 N
00
0
0.00
388
1.00
0.00
Column 5 thus states that in a normal solution 51 per cent, and in a
centi-normal solution 79 per cent of the hydrogen is ionic.
Sulphanion SO/' unites with all positive ions. The product, when
insoluble, appears as a precipitate. The introduction of barium ions,
for example, by adding a solution of barium nitrate or chloride, is em-
ployed as a means of recognizing the presence of sulphanion :
Ba"4-S04"<i±BaSOJ.
Naturally this test is given by solutions of all soluble acid and normal
sulphates. Since there are other barium salts which are insoluble in
water, but no common ones which are not decomposed by acids, dilute
nitric acid is first added to the solution supposed to contain the sul-
phanion. The other ions, if present, then give no precipitate with
barion.
Sulphates. — The acid sulphates, known also a,s'l>isulphates, may
be produced either by semi-neutralization of sulphuric acid with a
base, followed by evaporation : NaOH -f H2S04<^ H2O + NaHS04, or
by actions in which another acid is displaced, as in making hydrogen
chloride (p. 179).
The neutral (or normal) sulphates are obtained by complete neutral-
ization and evaporation, or by the second of the above methods when
a sufficient amount of the salt and a higher temperature are used :
NaCl + NaHS0 <
They are often made also by precipitation, by oxidation of a sulphate
at a high temperature, PbS + 202— > PbS04, or by addition of sulphur
trioxide to the oxide of a metal (p. 381).
Acid sulphates, when heated, yield pyrosulphates (p. 388). Normal
sulphates of the heavy metals decompose at a white heat, giving off
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 391
sulphur trioxide (p. 380). The sulphates of potassium, sodium, and
others of the more active metals, however, are not affected by heating.
When a sulphate, or indeed any salt of a sulphur acid, is heated
strongly with carbon, the oxygen is removed and a sulphide remains :
Na^SOj -f 4C — > Na^S -}- 4CO. Upon this is founded a general test for
the presence of sulphur in any substance. The material to be tested is
mixed with sodium carbonate. A small amount of the mixture is
placed on the end of a match, which has been charred and rendered
partially incombustible by previous application of sodium carbonate.
When the end of the match is now held in the reducing part of the
Bunsen flame, the compound of sulphur, if it contains oxygen, is
reduced to the form of sulphide. This, by interaction with the carbon-
ate, gives sodium sulphide, NaJS. W^hen the product of the reduction
is placed upon a silver coin and moistened, the sodium sulphide, if
present, produces a black stain of silver sulphide. This is known as
the hepar test, hepar being an old name for a sulphide.
Constitution of Hydroyeu Sulphate. — The formula which we
assign to sulphur trioxide is 0 = S ^ • It is in general our desire
to use the smallest possible valence, but here no reduction can be
effected below the value 6 for the sulphur, unless we join the oxygen
units to one another, as in the formula 0 = S V . This, however, would
\O
Q _ TT
suggest a relationship to hydrogen peroxide, i , which is not con-
O H
firmed, for hydrogen peroxide cannot be made from sulphuric acid.
Assuming, therefore, the above formula for sulphur trioxide, the
addition of the elements of water to it in the simplest fashion results
in the structures :
O . i . O H-0 O
*S' or ^Sf
0X/IXX0 H-(/ NX0
H
The second of these two modes of disposing of the water is the one
which, in parallel cases, is usually most feasible. Hardly any alter-
native to it is possible, for example, in representing the action in
which quicklime is slaked (p. 119) :
392 INORGANIC CHEMISTRY
This represents the change with little derangement of the original
structure and without alteration in the valence, while the first unwar-
rantably increases the valence to ten. There are other objections to
the first formula. In it the hydrogen is supposed to unite more
immediately with the sulphur, whereas, when the free elements are
concerned, hydrogen actually combines more readily with oxygen, and
forms a more stable compound with it than with sulphur. Again,
compounds like hydrogen sulphide, H— S— H, in which the hydrogenis
undoubtedly united to sulphur, are but slightly ionized, and are feeble
acids, while hydrogen sulphate is highly ionized.
Another fact is more satisfactorily accounted for by the second
formula. The addition of chlorine to sulphur dioxide must be shown
thus :
^° C1\ *°
< +•<*•-+ X »
^0 C\/ ^0
for chlorine has a much greater tendency to unite with sulphur than
with oxygen. When the product, sulphuryl chloride, is brought in
contact with water, sulphuric acid and hydrogen chloride are pro-
duced. Since water has the formula H— 0 — H, and two molecules of
water are used, this action is most simply accounted for, with the
minimum of disturbance in both molecules, by imagining the operation
to take place as follows :
H-0-; H Cl !<
• ; N o
H-O-j H Cl /
The hydrogen chloride is eliminated, and the other units of hydrogen,
originally without doubt attached to oxygen in the water, may be pre-
sumed to be still connected with that oxygen when they enter the
molecule of hydrogen sulphate.
This illustration shows the sort of reasoning, based upon the
chemical properties and the modes of formation of a substance, which
lead us to the devising of an appropriate graphic or structural formula
(cf. p. 224). The latter is not supposed in any sense to represent the
actual physical structure of the molecule, but simply to be a diagram-
matic representation of the chemical relations of the constituents and
of the chemical behavior of the whole. Formulae of this kind are in
continual use in the study of the compounds of carbon, but are seldom
required outside of that region.
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 393
Uses of Sulphuric Acid. — Sulphuric acid is employed in- almost
all chemical industries, and in some of them (q.v.) the quantities
required are very great. It is used, for example, in the first stage of
the Le Blanc process for the manufacture of soda, in the refining
of petroleum, in the manufacture of fertilizers, in the preparation of
nitroglycerine and gun-cotton, and in the production of coal-tar dyes.
OTHER ACIDS OF SULPHUR.
Hyposulphurous Acid. — The zinc salt of this acid crystallizes
out when zinc acts upon a solution of sulphur dioxide in absolute
alcohol :
Zn + 2S02->ZnS204.
Whether the acid is monobasic or dibasic, and its formula HS02 or
H2S204, is still uncertain.
Commercially a solution containing the sodium salt is made by
the interaction of zinc with a solution of sodium bisulphite charged
with excess of sulphur dioxide :
2NaHS03 + S02 + Zn -> Na,S204 + ZnS03 + H/).
The acid and its salts are very rapidly oxidized by the air, the
former giving sulphurous acid and then sulphuric acid. The above
solution of sodium hyposulphite is used in indigo dyeing, on account of
its high reducing power. Indigo, which is insoluble, is reduced by the
salt to indigo-white which passes into solution. When cloth saturated
with the mixture, however, is exposed to the air, the salt is rapidly
oxidized, the indigo-white likewise undergoes oxidation, and blue, in-
soluble indigo is formed once more (see Dyeing).
Sulphurous Acid. — This term is applied to the solution of sul-
phur dioxide in water. A portion of the sulphur dioxide remains dis-
solved physically, while another portion is in combination with the
water, forming sulphurous acid. This in turn is ionized, and chiefly,
after the manner of the weaker dibasic acids, into two ions, H* and
HSO/. A little S03" is formed from the latter. There are thus in such
a solution four mutually dependent equilibria :
S02(gas)<=±S02(diss'd)+ H20<=»H2S03<->H' + HS03'<=»H' + SO/'.
When the solution is heated, uncombined sulphur dioxide is disengaged
as a gas. The equilibria being thus disturbed, the ions of the acid
394 INORGANIC CHEMISTRY
unite, the acid molecules decompose, and soon all the above actions are
completely reversed and the whole of the gas passes off. Conversely,
when a base furnishing hydroxyl ions is added to the solution of the
acid, the hydrogen ions disappear, forming water, and the above actions
all proceed in a forward direction until, with a half -equivalent of the
base, the whole of the material has been converted into the form
HSO/, in association, of course, with the positive ions of the base.
With a full equivalent, neutralization follows and S08" is the product.
Properties of Sulphurous Acid. — The acid is so unstable that
it cannot be obtained at ordinary temperatures excepting in solution
in water. Chemically it is a comparatively weak acid. As a reducing
agent, it is slowly oxidized by free oxygen, and rapidly by oxidizing
agents, turning into sulphuric acid. Thus, when free halogens are
added to the solution, sulphuric acid and the hydrogen halide are formed:
H2S03 + H20 + I2 <=» H2S04 + 2HI.
In the particular case of iodine this action takes place only in very
dilute solution, since concentrated sulphuric acid decomposes hydrogen
iodide (cf. p. 237) and the action is reversed. This interaction is used
in chemical analysis as a means of estimating the quantity of sulphur-
ous acid in a liquid (cf. p. 236).
Hydrogen peroxide, potassium permanganate, and other oxidizing,
agents convert the substance into sulphuric acid likewise. It should
be noted that in these oxidations we have, not an addition of oxygen to
S02, but to the S08" or HSO/ ion of the acid, whereby it passes into
the S04" ion of sulphuric acid. The ion is much more easily oxidized
than is free sulphur dioxide itself.
Sulphurous acid has the power of uniting directly with many or-
ganic coloring matters and, since the products of this union are usually
colorless, it is employed as a bleaching agent. It is especially useful
with materials like silk, wool, and straw, which are likely to be de-
stroyed by chlorine. Exposure to sunlight causes the dissociation of
these colorless compounds, and so, with use, straw hats slowly recover
their original color. As a disinfectant it acts by addition likewise.
As a dibasic acid, sulphurous acid forms normal salts like Na2S03,
and acid salts like NaHS03.
Illustration of the Effect of Concentration on Speed of In-
teraction. — The oxidation of sulphurous acid by iodic acid may be
used to show the effect of concentration on the speed of an action
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 395
(p. 252). The iodic acid may most readily be made by dissolving po-
tassium iodate and sulphuric acid together in water, in such quantities
as would give a N/2 solution of each. When 1 c.c. of this N/2 iodic
acid is added to 100 c.c. of filtered starch emulsion, and the whole is
mixed with an equal volume of water containing 1 c.c. of N/2 sulphur-
ous acid, the blue color produced by the liberated iodine appears sud-
denly after the lapse of a minute or more :
2HI03 + 5H2S03 -» 5H2SO4 + H20 + I2.
With double the above quantities in the same amount of water, that is,
with double concentrations, the speed of the action is greatly increased
and the iodine becomes visible in less than half the time.
Sulphites. — The acid sulphites of the alkali metals (i.e., of potas-
sium and sodium) are acid in reaction, owing to the appreciable dissocia-
tion of the ion HSO/. The acid being a weak one, however, solutions
of the normal salts, Na^SOg, etc., are alkaline towards litmus (p. 344).
The sulphites are readily decomposed by acids giving free svilphurous
acid, and the latter partly decomposes, yielding sulphur dioxide.
Calcium bisulphite solution, Ca(HSO3)2, is used to dissolve the lig-
nin out of wood employed in the manufacture of paper. About 30
per cent of the wood is lignin. The rest is cellulose (C6H1006).,., and
constitutes the prepared pulp.
When heated, sulphites undergo decomposition. The sulphates,
being the most stable of all the salts of sulphur acids, are formed when
the salts of any of those acids are decomposed by heating. The nature
of the particular salt determines what other products shall appear.
Here, one molecule of the sulphite furnishes three atoms of oxygen,
sufficient to oxidize three other molecules, and leaves one molecule of
sodium sulphide behind :
4Na,SO, -» Na2S + 3Na2SO4.
The acid sulphites (bisulphites) first lose sulphurous acid, before chang-
ing in this way. Thus, sodium hydrogen sulphite begins by decompos-
ing as follows :
2NaHSO8 -» NaoSOs + H2SO3 (or H2O + S02).
The acid salts of volatile acids, when heated, all decompose in this
way (cf. pp. 241, 388).
The sulphites are as readily oxidized as is the acid itself. They are
396 INORGANIC CHEMISTRY
slowly converted, both in solution and in the solid form, by the influ-
ence of the oxygen of the air, into sulphates. It is interesting to note
that the addition of sugar or glycerine to a solution of a sulphite re-
duces the speed of oxidation by free oxygen very markedly. These
substances act as catalytic agents ; and the present case shows that
agents of this kind may not only increase the speed of actions, which
is their usual function, but may also have a restraining influence.
Thiosulphuric Acid. — This acid is not known in the free con-
dition, but its salts are in common use in the laboratory and commer-
cially. The sodium salt, for example, is prepared by boiling a solu-
tion of sodium sulphite with free sulphur. The action is something
like the addition of oxygen to sulphurous acid :
Na.SOs + S -> Na2S203 or S03" + S -> S203".
The product, thiosulphate of sodium, is used in photography as a
solvent for salts of silver, and is commonly known as hyposulphite of
soda. Its solution in water forms the fixing bath of the photographer.
By the addition of acids to a solution of sodium thiosulphate, the
thiosulphuric acid is set free, but the latter instantly decomposes:
, + 2HC1 +± H2S208 + 2NaCl,
Even carbon dioxide from the air, giving carbonic acid (q-v.\ pro-
duces this effect slowly in fixing solutions. The actions being rever-
sible, preliminary addition of a sulphite to the solution helps to sustain
the reverse action, in which sulphurous acid is a factor, and so pre-
serves the solution. The delay in the appearance of the precipitate
of sulphur in dilute solutions is due to the temporary existence of a
supersaturated solution (cf. p. 160) of the free element.
Iodine acts upon sodium thiosulphate solution, giving sodium
tetrathionate :
2Na2S208 + I2 -» 2NaI + Na2S4O6.
This action is used, by employment of a standard solution of sodium
thiosulphate, for estimating quantities of free iodine in analysis. The
disappearance of the color of the latter indicates that a sufficient
amount of the salt has been employed. When chlorine is used, the
oxidation is- more complete. The products are sodium sulphate, sul-
phuric acid, and hydrochloric acid :
, + 5H20 + 4C12 -» Na^SO.-f H2SO4 + 8HC1.
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 397
In consequence of the very great amount of free chlorine which the
sodium thiosulphate is thus able to transform, it is employed, as
antichlor, for the purpose of removing chlorine from bleached fabrics.
Persulphuric Acid. — This, like the other acids just mentioned,
is unstable, and can be kept only in very dilute solution. Its salts,
however, are coming into use for commercial purposes and for " redu-
cing " negatives in photography. When a discharge of electricity is
passed through a mixture of sulphur trioxide and oxygen, drops of
liquid are formed which appear to have the composition S207, and
when dissolved in water give dilute persulphuric acid, S2O7 + H20 — >
H2S208. More significant of its relations is its formation, to some
extent, when concentrated sulphuric acid and a strong solution of
hydrogen peroxide are mixed :
2H.SO, + H202 <=* H2S208 + 2H20.
This action is reversible.* Under some circumstances, monopersul-
phuric acid is formed: HgSC^+HgOgjp* H2S05+H20. Interesting in
its way, also, is its production in the electrolysis of aqueous sulphuric
acid :
2HS04 + 2 0 -> H2S208.
This action is most conspicuous in rather concentrated solutions in
which hydrosulphaniou is plentiful (cf. p. 389) and when a small anode,
resulting in severe crowding of the HS04 radicals as they are liber-
ated, is employed. The salts are prepared by electrolyzing sodium
hydrogen sulphate NaHS04 in concentrated solution. The persul-
phuric acid, formed by the union of the negative ions in pairs, under-
goes double decomposition with the excess of sodium bisulphate, and
the less soluble sodium persulphate crystallizes out. The other salts
are made by double decomposition from this one.
* This action and the next are not easily classifiable under any of the ten
kinds formerly discussed (p. 187). They consist in the union of H and OH to form
water :
HSO4; H + HO : OH + H |S04H — > 2H20 + (S04H)2.
Neutralizations (p. 351) they are not, because the interacting substances are both
acids. Just as the loss of water from one acid gives an anhydride, so here, the
loss of water between two acids gives a mixed anhydride (see Chlorosulphuric
acid, below).
398 INORGANIC CHEMISTRY
The persulphates decompose readily when heated, yielding pyro-
sulphates and oxygen :
2K2S208 -i» 2K2S207 + 02.
The solution of the acid is an active oxidizing agent :
+ H20 -+ 2H2S04 ( +• O).
Polythionic Acids. — Di-, tri-, tetra-, and pentathionic acid (p.
381) are all formed simultaneously when sulphur dioxide and hydro-
gen sulphide gases are passed alternately into water, although the
gases themselves (p. 373) interact to produce simply free sulphur and
water :
H2S + 3S02 -» H2S406,
2H2S + 6S02 -> H2S3O6 + H2S506,
3H2S + 9SO2 -> H2S206 + 2H2S506.
Most of these acids and their salts are of minor interest and need not
be discussed.
The production of sodium tetrathionate by the action of iodine
upon sodium thiosulphate has already been mentioned.
When manganese dioxide is treated with sulphurous acid, it inter-
acts very rapidly and a solution of manganous dithionate is obtained :
Mn02 + 2H2S03 -> MnS206 + 2H20.
The salts of these acids are in many cases fairly stable, but the
acids themselves decompose readily when set free.
COMPOUNDS OF SULPHUR AND CHLORINE.
Sulphur Monochloride. — When chlorine gas is passed over
heated sulphur it is absorbed, and a dark reddish-yellow liquid, boiling
at 138°, is obtained. The molecular weight of this substance, as shown
by the density of its vapor, indicates that it possesses the formula
S2C12. When thrown into water it is rapidly hydrolyzed, producing
sulphur dioxide and free sulphur :
2S2C12 -f 2H2O -» SO2 + 4HC1 + 3S.
Sulphur itself dissolves very freely in the monochloride, and the
solution is employed in vulcanizing rubber.
THE OXIDES AND OXYGEN ACIDS OF SULPHUR 399
By surrounding sulphur monochloride with a freezing mixture,
and treating it with excess of chlorine, a liquid tetrachloride SC14 is
formed. This gives up chlorine, and turns into the monochloride
again, when allowed to become warm.
Thionyl Chloride. — By the action of sulphur dioxide gas upon
phosphorus pentachloride, part of the oxygen in the former is replaced
by chlorine :
S02 + PC15 -> SOCLj +POC18.
The products are thionyl chloride and phosphorus oxychloride. The
former is a colorless liquid, boiling at 78°, and is separated from the
latter (b.-p. 107°) by fractional distillation (see Petroleum). It is de-
composed immediately on contact with water :
SOC12 + H20 -> SO2 + 2HC1.
Sulphuryl Chloride. — Sulphur dioxide and chlorine gases unite
when exposed to direct sunlight to form a liquid known as sulphuryl
chloride S02C18. When camphor is introduced into the vessel the
union takes place much more rapidly, owing to some catalytic effect
of this substance. The compound is a colorless liquid, boiling at 69°.
With water it gives sulphuric acid and hydrogen chloride (p. 392).
When a strictly limited amount of water is supplied, a partial action
of the same nature occurs, and the product is known as chlorosulphuric
acid :
S02CL, + H20 -» S02 < + HC1.
This intermediate compound may be formed also by the addition of
hydrogen chloride to sulphur trioxide.
Exercises. — 1. What ground is there for assigning the formula
S02 instead of S204 to sulphur dioxide (p. 193) ?
2. Explain why nitric acid is completely displaced by the action of
sulphuric acid on sodium nitrate (p. 384).
3. How many times, on an average, does a molecule of nitrous
anhydride go through the cycle of changes by which sulphuric acid is
produced before it is eliminated in some other form (p. 386) ?
4. Make a list of, and classify, the various applications of sulphuric
acid to the liberation of other acids.
400 INORGANIC CHEMISTRY
5. 'Formulate the behavior of the hydrosulphanion (p. 389) when
a solution of barium chloride is added to a rather concentrated solu-
tion of sulphuric acid.
6. Can you give any reasons for preferring to regard KHS04, and
substances like it, as acid salts rather than double salts of the form
ig304, H2S04 (p. 358)?
7. From a consideration of the physical conditions, how can you
account for the fact that hyposulphurous acid is oxidized by free
oxygen first to sulphurous acid and then to sulphuric acid (p. 393),
while moist free sulphur, even with oxidizing agents, gives sulphuric
acid directly (p. 371) ?
8. Why were not disulphuric acid and monopersulphuric acid
placed in the list on p. 381 (cf. p. 278) ?
9. Write in ionic form the equation for the interaction of sodium
thiosulphate and iodine in aqueous solution.
10. Restate the import of the following sentence in such a way as
to avoid the use of the hypothesis of ions : Liquefied sulphur dioxide
" ionizes substances dissolved in it as well as does water " (p. 379).
CHAPTER XXIII
SELENIUM AND TELLURIUM: THE PERIODIC SYSTEM
ALONG with sulphur, chemists group two other elements, selenium
(Se, at. wt. 79.2) and tellurium (Te, at. wt. 127.6). If the nature of
the chief compounds of sulphur is kept in mind, the close analogy
between the nature and chemical behavior of the three elements and
their corresponding compounds will be noticed at once (see Chemical
relations of the sulphur family, below).
SELENIUM.
Occurrence and Properties of the Element. — Selenium (Gk.
o-eX^vi;, the moon) occurs free in some specimens of native sulphur,
and in combination often takes the place of a small part of the sulphur
in pyrite (FeS2). It is found free in the dust-flues of the pyrite-
burners of sulphuric acid works. The familiar forms are, the red pre-
cipitated variety, which is amorphous and soluble in carbon disulphide,
and the lead-gray, semi-metallic variety, obtained by slow cooling
of melted selenium, which is insoluble, and melts at 217°. In the
latter form it has some capacity for conducting electricity, which is
increased by exposure to light in proportion to the intensity of the
illumination. It boils at 680°, and at high temperatures has a vapor
density corresponding to the formula Se2.
The element unites directly with many metals, burns in oxygen to
form selenium dioxide, and unites vigorously with chlorine.
Hydrogen Selenide. — Ferrous selenide, made by heating iron
filings with selenium, when treated with concentrated hydrochloric
acid gives hydrogen selenide :
FeSe + 2HC1 <=> H2Sef + FeCl2.
The compound is a poisonous gas, which possesses an odor recalling
rotten horse-radish, and is soluble in water. The solution is faintly
acid in reaction, and deposits selenium when exposed to the action of
401
402 INORGANIC CHEMISTRY
the air (cf. p. 373). Other selenides, which, with the exception of
those of potassium and sodium, are insoluble in water, may be pre-
cipitated by leading the gas into solutions of soluble salts of appro-
priate metals (cf, p. 375).
Selenium Dioxide and Selenious Acid. — The dioxide (Se02)
is a solid body formed by burning selenium or evaporating a solution
of selenious acid H2Se03. The latter may be made directly by oxidiz-
ing selenium with boiling nitric acid or aqua regia (q.v.). Unlike
sulphur (p. 371), the element gives little of the higher acid H2Se04 by
this treatment. In aqueous solution the acid is easily reduced to
selenium. In fact, sulphurous acid can be oxidized by it :
H2Se03 + 2H2S08 -> 2H2S04 + 11,0 + Se.
Selenic Acid. — No trioxide is known. Selenic acid (H2Se04) is
made by using the most powerful oxidizing agents with an aqueous
solution of selenious acid. It can be isolated as a white solid. It is
itself a powerful oxidizing agent, and, even in dilute solution, liberates
chlorine from hydrochloric acid :
H2Se04 + 2HC1 -> I^SeO, + H20 + CL,.
Sulphuric acid (cf. p. 389), on the other hand, is an oxidizing agent
only in somewhat concentrated form, and even then it can oxidize
hydrobrornic acid (p. 231), but not hydrochloric acid.
Chlorides of Selenium. — The chlorides are formed by direct
union of the elements. The tetrachloride, a yellow crystalline sub-
stance, is formed when the monochloride is heated : 2Se2Cl2 — > 3Se +
SeCl4, the behavior in the case of sulphur chlorides being just the
inverse of this (p. 399).
TELLURIUM.
Tellurium (Lat. tellus, the earth) occurs in sylvanite in combina-
tion with gold and silver. It is a white, metallic, crystalline sub-
stance, melting at 462°. When formed by precipitation it is a black
powder. It conducts electricity to some extent. The vapor density
corresponds to the formula Te2.
The free element unites with metals directly, and burns in air to
form the dioxide.
SELENIUM AND TELLURIUM: THE PERIODIC SYSTEM 403
Hydrogen telluride (H2Te) is made by the action of acids on
metallic tellurides, and its aqueous solution is rapidly oxidized by air
with precipitation of tellurium. The tellurides of the alkali metals
are soluble in water, the others are insoluble.
Tellurious acid (H2Te03) is formed by oxidizing the element with
nitric acid, and is a crystalline solid, little soluble in water. It is a
feeble acid, of which many salts have been made. It is also some-
what basic, a sulphate (2Te02, S08) and a nitrate (Te208(OH)N08) being
known. In this respect it differs markedly from sulphurous acid.
Telluric acid is made by oxidizing tellurious acid in aqueous solu-
tion with chromic acid (p. 374). It is difficultly soluble in water. It
does not affect indicators, and is therefore actually more feebly acidic
than is hydrogen sulphide. It is obtained as a solid having the com-
position H6Te06 (or 3H20, Te03) on evaporating the solution. When
heated, this body loses water, some of the trioxide (Te03) being
formed. The last is a yellow solid, which shows no tendency to
recombine with water, in this respect resembling silica (q.v.). Tel-
lurates of the alkali metals may be made by heating the tellurites
with potassium or sodium nitrate : K2TeOs + KN08 — > K2Te04 +
KNO2.
Tellurium forms two very stable chlorides, TeClj, and TeCl4, which
are decomposed by water. The second, however, exists in solution
with excess of hydrogen chloride: TeCl4+3H20 <=± H2Te08 + 4HC1,
showing the tellurious acid to be basic in properties and the element
tellurium to be, to a certain degree, a metal (see Chap, xxxii).
The Chemical Relations of the Sulphur Family. — It will be
seen that sulphur, selenium, and tellurium are bivalent elements when
combined with hydrogen or metals. In combination with oxygen they
form unsaturated compounds of the form XIV02, while their highest
valence is found in S03, Te03, and H2SeO4, where they must be sexi-
valent. The general behavior of corresponding compounds is very
similar. At the same time, there is in all cases a progressive change as
we proceed from sulphur through selenium to tellurium. The elemen-
tary substances themselves, for example, become more like metals,
physically, and they show higher and higher melting-points. The
affinity for hydrogen decreases, as is shown by the increasing ease
with which the compounds H2X are oxidized in air. The affinity for
oxygen likewise decreases, for the elements become increasingly diffi-
cult to raise to the highest state of oxidation. On the other hand,
404 INORGANIC CHEMISTRY
the tendency to form higher chlorides becomes greater. We note also
that the compounds H2X04 become less and less active as acids, and a
basic tendency begins to assert itself.
THE PERIODIC SYSTEM.
Classification, or the arrangement of facts on the basis of likeness,
is part of the method of science. In chemistry the multitude of facts
is not less than in other sciences, and the necessity of arrangement
equally urgent. It is needed, to make possible the systematic descrip-
tion of the ascertained facts, and to furnish a guide in investigation, by
suggesting stochastic hypotheses, and so pointing out directions in
which new facts of interest may be found. Thus, we have treated the
halogens as a group ; and chemists, knowing how hypochlorous acid
(HC10) and perchloric acid (HC104), and their salts, are made, have
been led to attempt to obtain related substances, like HIO and HBr04
and their salts, by methods suggested by analogy.
At first sight, the most definite method of classification would
appear to be the grouping of elements of like valence. But this brings
together sodium and chlorine — an element whose hydrogen compound
is unstable and without markedly characteristic properties, and whose
hydroxyl compound is an active base, with an element whose hydrogen
compound is an active acid and whose hydroxyl compound is, in a
feeble degree, an acid also. This method homologates similar and con-
trasting elements indiscriminately.
Metallic and Non-Metallic Elements. — Thus far we have found
the division into metallic and non-metallic elements very serviceable for
classification in terms of chemical relations (p. 177). This distinction
we shall continue to employ. The metals, or positive elements (p. 119),
(1) form positive ions containing no other element (cf. p. 337). Thus
the metals give sulphates, nitrates, carbonates, and other salts, which
furnish a metallic ion together with the ions SO/', NO/, and CO/'. (2)
Their hydroxides, KOH, Ca(OH)2, etc., give the same metallic ion, and
the rest of the molecule forms hydroxidion. That is to say, their
hydroxides are bases and their oxides are basic. They often enter
with other elements into the composition of a negative ion, as is the case
with manganese in K.Mn04, with chromium in K2.Cr2O7, and with silver
in K.Ag(CN)2.* But the most definitely metallic elements form with
* The mode of division into ions is shown by the position of the period in the
formula.
SELENIUM AND TELLURIUM: THE PERIODIC SYSTEM 405
oxygen such a negative ion only while exhibiting a different valence
from that which they possess when acting as positive elements. Thus,
manganese when a positive element has the valences two and three,
MnO and Mn208, Mn.Cl2 and Mn.Cl8, Mn.S04 and Mn2.(S04)8, etc.,
while in permanganates we have potentially the oxide Mn207, (K20,
Mn207 = 2KMn04), in which it is heptavalent. The graphic formulae
express the difference in valence :
Cl /C1 ^°
Mn/ Mn-Cl K-0-Mn=0
\C1
Manganese forms no heptachloride or heptanitrate (corresponding to
the heptoxide) in which manganese alone would form a heptavalent
positive ion on the one hand, and no compound in which bivalent or
trivalent manganese with oxygen form the negative ion, on the other.
Chromium (q.v.~) is bivalent or trivalent when a metal, and sexivalent
when acting as a non-metal. The various hydroxides of these elements
are bases or acids in accordance with this distinction. If we knew only
the compounds in which manganese and chromium are heptavalent and
sexivalent, respectively, we should regard them as non-metallic ele-
ments pure and simple, the metallic appearance of the free elements to
the contrary notwithstanding.
The non-metals or negative elements, (1) are found chiefly in nega-
tive ions. They form no nitrates, sulphates, carbonates, etc., for they
could not do so without themselves constituting the positive ion. We
have no such salts of oxygen, sulphur, carbon, or phosphorus, for ex-
ample. (2) Their hydroxides, like C102OH, P(OH)3, S02(OH)2, furnish
no hydroxyl ions, as this would involve the same consequence. These
hydroxides are divided by dissociation, in fact, so that the non-metal
forms part of a compound negative radical, and the other ion is hydrion,
C103.H, P03H.H2, S04.H2. (3) Their halogen compounds, like PC18
(p. 181) and SgCLj (p. 398), are completely hydrolyzed by water, and the
actions are not, in general, reversible. The halides of the typical
metals are not hydrolyzed (see Chap, xxxii).
The distinction is not perfectly sharp, however. Thus, zinc (q.v.~)
gives, both salts like the sulphate, Zn.S04, and chloride, Zn.CLj, and com-
pounds like sodium zincate (p. 99), ZnO^Na^ showing the same valence
in both classes :
7/Cl 7 /0-Na
Zn\Cl Zn\0-Na
406 INORGANIC CHEMISTRY
Its hydroxide ionizes in two ways, Zn.(OH)2 and Zn02.H2. Similarly
tellurious acid, H2Te08, acts both as acid and base (p. 403). We shall
find this double behavior conspicuous in the compounds of arsenic and
antimony. In spite of the partial merging of the two classes of ele-
ments, however, the general distinction is worth preserving (see Chap,
xxxii).
Classification by Atomic Weights. — A closer discrimination
than that furnished by these two categories is required, however, and
a study of the order into which the elements fall when arranged accord-
ing to their atomic weights has provided this.
The first indication of a significant relation between the atomic
weights and the properties of the elements was given by a fact noted
by Dobereiner (1829). He drew attention to the existence of closely
similar elements in sets of three (triads), where the central element
was intermediate in properties between the two others, and the atomic
weight of the central element was almost the exact arithmetical mean
of the weights of the other two. The three following triads illustrate
this relation :
Chlorine . . . 35.45 Sulphur . . . 32.06 Calcium . . . . 40.1
Bromine . . . 79.96 Selenium . . . 79.2 Strontium . . . 87.6
Iodine . . . 126.97 Tellurium . . . 127.6 Barium .... 137.4
Mean of Cl and I, 81.1 Mean of S and Te, 79.8 Mean of Ca and Ba, 88.7
Newlands (1863-4) called attention to the existence of a surprising
regularity when the elements were placed in the order of ascending
atomic weight. Omitting hydrogen (1) the first seven are : lithium (7),
gluciuum (9), boron (11), carbon (12), nitrogen (14), oxygen (16),
fluorine (19). These are all of totally different classes, and include first
a metal forming a strongly basic hydroxide, then a metal of the less
active sort, then five non-metals of increasingly negative character, the
last being the most active non-metal known. The next element after
fluorine (19) is sodium (23), which brings us back sharply to the
elements that form strongly basic hydroxides. Omitting none, the next
seven elements are : sodium (23), magnesium (24.4), aluminium (27),
silicon (28.4), phosphorus (31), sulphur (32), chlorine (35.5).
In this series there are three metals of diminishing positiveness,
followed by four non-metals of increasing negative activity, the last
being a halogen very like fluorine. On account of the fact that each
element resembles most closely the eighth element beyond or before it
in the list, the relation was called the law of octaves. After chlorine
SELENIUM AND TELLURIUM : THE PERIODIC SYSTEM 407
the octaves become less easy to trace. Potassium (39) follows chlo-
rine and corresponds satisfactorily to sodium, but it is not until seven-
teen successive elements have been set down that we reach one closely
resembling chlorine, namely, bromine.
That this periodicity in chemical nature is more than a coincidence
is shown by the fact that the valence and even the physical properties,
such as the specific gravity, show a similar fluctuation in each series,
and recurrence in the following one. In the first two series the com-
pounds with other elements are of the types :
LiCl, G1C1., BC1,, CC1,;
Thus the valence towards chlorine or hydrogen ascends to four and
then reverts to one in each octave. The highest valence, shown in
oxygen compounds, ascends from lithium to nitrogen with values one
to five, and then fails because compounds are lacking. In the second
octave, however, it goes up continuously from one to seven.
Again, the specific gravities of the elements in the second series,
using the data for red phosphorus and liquid chlorine, are :
Na 0.97, Mg 1.75, Al 2.67, Si 2.49, P 2.14, S 2.06, Cl 1.33.
Of greater significance chemically are the related numbers represent-
ing the volumes in cubic centimeters occupied by a gram-atomic weight
of each element (the atomic volumes) :
Na 24, Mg 14, Al 10, Si 11, P 14, S 16, Cl 27.
A similar regular fluctuation is shown by all the physical properties of
corresponding compounds.
Mendelejeff's Scheme. — In 1869 Mendelejeff published an im-
portant contribution towards adjusting the difficulty which the ele-
ments following chlorine presented, and developed the whole concep-
tion so completely that the resulting system of classification has been
connected with his name ever since. Almost simultaneously Lothar
Meyer made similar suggestions, but did not urge them with the same
conviction or elaborate them so fully. The following table, in which
the atomic weights are expressed in round numbers, is a modification
of one of Mendelejeff's.
408
INORGANIC CHEMISTRY
OS
® CO
PH "3
>
s*
o
t3 CO
•-'
02 CO
U co
SELENIUM AND TELLURIUM: THE PERIODIC SYSTEM 409
The chief change from the arrangement in simple octaves is that the
third series, beginning with potassium, is made to furnish material for
two octaves, potassium to manganese and copper to bromine, and is
called a long series. The valences fall in with this plan fairly well.
Copper, while usually bivalent, forms also a series of compounds in
which it appears to be univalent. Iron, cobalt, and nickel cannot be
accommodated in either octave, as their valences are always two or
three. At the time Mendelejeff made the table, three places in the
third series had to be left blank, as a trivalent element [Sc] was lack-
ing in the first octave, and a trivalent [Ga] and a quadrivalent one [Ge]
in the second. These places have since been filled, as we shall pres-
ently see. The first two, (the short) series have been split in the table,
as lithium and sodium closely resemble potassium, while the remaining
members of these series fall more naturally over the corresponding ele-
ments of the second octave of the third series.
The fourth series (long) is nearly complete. It begins with an
active alkali metal, rubidium, and ends with iodine, a halogen.
The rule of valence is strictly preserved throughout the series, and in
general the elements fall below those which they most closely resemble.
The fifth, sixth, and seventh (long) series are incomplete, but the
order of the atomic weights and the valence enable us satisfactorily to
place those elements which are known. The chemical relations to ele-
ments of the fifth series justify the position assigned to each. Caesium,
for example, is the most active of the alkali metals ; barium has always
been classed with strontium, and bismuth with antimony.
In two cases a slight displacement of the order according to atomic
weights is necessary. Cobalt is put before nickel because it resembles
iron more closely. Tellurium and iodine are placed in that order to
bring them into the sulphur and halogen groups respectively. Their
valence and other chemical relations both require this. The general
agreement, however, is very remarkable.
General Relations in the System. — In every octave the valence
towards oxygen ascends from one to seven, while that towards hydro-
gen, in the cases of the four last elements (when they combine with
hydrogen at all), descends from four to one. The atomic volume disre-
gards the subsidiary octaves in the long series. It descends towards
the middle of each series (long or short), and ascends again towards
its initial value. The other physical properties fluctuate within the
limits of each series in a similar way. The values of each physical
410 INORGANIC CHEMISTRY
constant for corresponding members of the successive series do not
exactly coincide, however. A progressive change, as we descend each
vertical column, is the rule. Thus the specific gravities (water = 1)
of the alkali metals rise from lithium (0.59) to caesium (1.85). In the
same group the melting-points descend from lithium (186°) to caesium
(26.5°).
It must be stated that as yet no exact mathematical relation be-
tween the values for any property and the values of the atomic weights
has been discovered ; only a general relationship can be traced. An-
ticipating the discovery of some more exact mode of stating the rela-
tionship in each case, and remembering that similar values of each
property recur periodically, usually at intervals corresponding to the
length of an octave or series, the principle which is assumed to un-
derlie the whole, the periodic law, is stated thus : All the properties of
the elements are periodic functions of their atomic -weights.
That the chemical relations of the elements vary just as do the phy-
sical properties of the simple substances is easily shown. Thus, each
series begins with an active metallic (positive) element, and ends with
an active non-metallic (negative) element, the intervening elements
showing a more or less continuous variation between these limits.
Again, the elements at the top are the least metallic of their respec-
tive columns. As we descend, the members of each group are more
markedly metallic (in the first columns), or, what is the same thing,
less markedly non-metallic (in the later columns ; cf. p. 403).
In the first series boron is the first non-metal we encounter. In the
second series silicon is the first such element. In the third there is
more difficulty in deciding. Titanium, vanadium, and germanium are
usually, though with questionable propriety, classed as metals. Sele-
nium is undoubtedly a non-metal. Arsenic is, on the whole, a non-
metal. In the fourth series tellurium is commonly considered to be
the first non-metal. Thus a zigzag line, shown in the table, separates
all the non-metals from the rest of the elements, and confines them in
the right-hand upper corner.
A more compact form of the table, also based upon one of Mendele-
jeff's, may now be given (Table II). The only difference between
this and the other is that the two octaves of each long series have been
placed in the same set of seven main columns. The irregular sets of
three elements, consisting of the iron, palladium, and platinum groups,
occupy a column on the right of the main columns, and are often
called collectively the eighth group. For completeness, the newly
SELENIUM AND TELLURIUM: THE PERIODIC SYSTEM 411
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412 INORGANIC CHEMISTRY
discovered elements, found chiefly in the air, have been placed at the
left-hand side. Since they do not enter into combination at all, their
valence may appropriately be given as zero. With the exception of
argon, the values of their atomic weights agree well with this assign-
ment. Hydrogen is the only element whose place is still in debate.
Many rare elements have also been omitted, and tantalum has been
placed immediately after cerium. The valence is shown by the
general formulae at the head of each column.
Applications of the Periodic System. — The system has found
application chiefly in four ways :
1. By leading to the prediction of new elements.
2. By furnishing a comprehensive classification of the elements,
arranging them so as to exhibit the relationships among the physical
and chemical properties of the elements themselves and of their com-
pounds. Constant use will be made of this property of the table in
the succeeding chapters.
3. By enabling us to decide on the correct values for the atomic
weights of some elements, when the eqtiivalent weights have been
measured, but no volatile compound is known (cf. pp. 201 and 212).
4. By suggesting problems for investigation.
The Prediction of New Elements. — Mendelejeff (1871) drew
attention to the blank then existing between calcium (40) and titanium
(48). He predicted that an element to fit this place would have an
atomic weight 44 and would be trivalent. From the nature of the
surrounding elements, he very cleverly deduced many of the physical
and chemical properties of the unknown element and of its com-
pounds. He named it eka-boron (Skr. eka, one). In 1879 Nilson
discovered scandium (44), and its behavior corresponded closely with
that predicted for eka-boron.
In the same paper Mendelejeff described two other elements,
likewise unknown at the time. They were to occupy vacant places
between zinc and arsenic, and were named eka-aluminium and eka-
silicon. In 1875 Lecoque de Boisbaudran found gallium, and in 1888
Winkler discovered germanium, and these blanks were filled. Other
possible elements, such as eka-rnanganese and . eka-antimony, were
described, but still remain to be discovered.
Application to Fixing Atomic Weights. — Atomic weights have
been fixed with the assistance of the periodic system on several occasions.
SELENIUM AND TELLURIUM : THE PERIODIC SYSTEM 413
Thus the atomic weight of uranium was thought to be 120 until it was
observed that no place near to antimony (120) remained unoccupied.
With the value 240 (now 238.5), the element was accommodated at
the foot of the column containing those which it most resembled.
Again, the equivalent weight of indium was 38, and, as the element was
supposed to be bivalent, it received the atomic weight 76. It was quite
out of place near arsenic (75), however, being decidedly a metal. As a
trivalent element with the atomic weight 115, it fell between cadmium
and tin. Later work fully justified the change. Still again, glucinum,
with equivalent weight 4.5, resembled aluminium so strongly that
it was thought to be trivalent, like that element, and to have the
atomic weight 13.5. But the only vacancy in the first series then
existing was between lithium (7) and boron (11), and subsequent in-
vestigations showed that the properties of glucinum placed it most
fittingly in that position as a bivalent metal with the atomic weight 9.
Finally, and quite recently, radium (q.v.~) has been discovered, and
found to have the equivalent weight 112.5 and to resemble barium.
If, like barium, it is bivalent, it occupies a place under this element,
in the last series.
The System Suggests Lines of Investigation. — The periodic
system has been of constant service in the course of inorganic research,
and has often furnished the original stimulus to such work as well.
For example, the atomic weights of the platinum metals at first placed
them in the order, Ir (197), Pt (198), Os (199), although the resem-
blance to iron and ruthenium would have led us to expect that osmium
should come first. For similar reasons platinum should have come last,
under palladium. A reinvestigation of the atomic weights, suggested
by these considerations, was undertaken by Seubert, and the old values
were found in fact to be very inaccurate. He obtained : Os = 191,
Ir = 193, Pt = 195.
Again, the atomic weight of tellurium bore the value 128 when the
table was first constructed, and it was confidently expected that re-
examination would bring this value below that of iodine (then 127,
now 126.97). The problem has proved more difficult than usual, and
several most careful studies of the subject have been made by chemists,
using different methods. Absolute certainty has not yet been reached,
but it seems probable that the real value of the atomic weight is not
far from Te = 127.6, and therefore more than half a unit greater than
that of iodine. Since, however, mathematical correspondence is found
414 INORGANIC CHEMISTRY
nowhere in the system, the existence of marked inconsistencies like
this need not shake our confidence in its value when it is used with
due consideration of the degree of correspondence to be expected.
Originally lead, although it fell in the fourth column, possessed
only one compound, PbO2, in which it seemed to be undoubtedly quad-
rivalent. Search for salts of the same form, however, speedily yielded
the tetrachloride (PbCl4), tetracetate, and many others. As has been
stated already, there are many elements between lanthanum and
tantalum, about which sufficient knowledge has not yet been accumu-
lated to make possible their definite allotment to places in the table.
The existence of osmic acid (Os04), and a corresponding compound
of ruthenium, suggests that other compounds of the elements of the
eighth group, displaying the valence eight, may be capable of prepa-
ration. The collocation of copper, silver, and gold, in the same column
with the alkali metals, is not at present perfectly satisfactory, and sug-
gests the advisability of strengthening their position, if possible, by
further investigation.
The System as a Guide in Classification. — Having disposed
of the halogen and sulphur families, situated, respectively, in the
seventh and sixth columns of the table, we shall next take up nitrogen
and phosphorus from the right side of the fifth column. This family
includes nitrogen, phosphorus, arsenic, antimony, and bismuth, and
brings us for the first time in contact with a group which is partly
non-metallic and partly metallic. Then from the fourth column, we
shall select carbon and silicon, and from the third boron, leaving the
other, more decidedly metallic elements for later treatment.
Exercises. — 1. Can you explain the presence of free selenium in
the flues of pyrite-burners (p. 385) ?
2. Why does the existence of tellurium tetrachloride in solution
in aqueous hydrochloric acid show tellurium to be somewhat metallic
in chemical properties ?
3. How should you attempt to obtain HIO and HBr04, or their
salts ?
4. Make a list of bivalent elements and criticize this method of
grouping as a means of chemical classification.
5. Write down the symbols of the elements in the fourth series
of Table I (that beginning with rubidium, and ending with iodine).
Record the valence of each element toward oxygen, using for reference
the chapters in which the oxygen compounds are described.
CHAPTER XXIV
NITROGEN AND ITS COMPOUNDS WITH HYDROGEN
WHEN the oxygen of the air is removed (p. 62), a gas remains
which is largely nitrogen. When first discovered, the most conspicu-
ous property this gas was observed to have was indifference ; it did not
support combustion or life. The latter fact led to its being named
azote (Gk. £O>TIKOS, life). While this name is still preserved in the
French language, in English the name of the substance is derived
from the fact that it is an important constituent of saltpeter KN08
(Lat. nitrum).
Tlie Chemical Relations of the Element. — In some compounds
nitrogen is trivalent, while in others, particularly those containing
oxygen and other negative elements, it is quinquivalent.
The compounds of nitrogen are often extremely active and inter-
esting. Those of them which we have to discuss in inorganic chem-
istry are ammonia (NHg) and nitric acid HN03, and several
related substances. The organic compounds containing nitrogen are
very numerous and possess highly characteristic properties. Some,
like nitroglycerine (q.v.~) and gun-cotton, are violently explosive;
others, like antipyrine, show great physiological activity ; still others,
such as the aniline and other organic dyes, provide us with beautiful
and useful coloring matters.
Occurrence. — Apart from the presence of free nitrogen in the
air, the element is found in many forms of combination. The nitrates
of potassium and sodium are found in Bengal and Peru respectively.
Natural manures, such as guano, contain large quantities of nitrogen
compounds, and owe part of their value as fertilizers to this fact.
Nitrogen is an essential constituent of vegetable and animal matter.
The albumins, for example, which often make up a considerable part
of such matter, contain on an average about 15 per cent of combined
nitrogen.
415
416 INORGANIC CHEMISTRY
Preparation. — Nitrogen containing about one per cent of argon
(q.v.) is easily obtainable from purified air when the oxygen of the
latter is removed. For this purpose phosphorus is frequently burned
in the air, or the air is passed over heated copper.
When pure nitrogen is required, it must be obtained from chemical
compounds, in order that it may be free from argon. The simplest
method is to heat ammonium nitrite. This substance decomposes
readily, even at temperatures little above the ordinary, according to
the following equation :
NH4N02-+2H20 + N2.
In practice, since ammonium nitrite cannot easily be kept, a mixture
of an ammonium salt with some salt of nitrous acid is employed.
Thus, when strong solutions of ammonium chloride and sodium nitrite
are mixed, a double decomposition results in the formation of ammo-
nium nitrite, NH4C1 + NaN02 <=± NH4N02 + NaCl, and this breaks
up when heat is applied, giving nitrogen.
In certain cases it is convenient to prepare nitrogen by the oxida-
tion of ammonia (NH3), by passing the latter over heated cupric oxide,
or by the reduction of nitric oxide (NO) by passing this gas over heated
copper.
Physical Properties. — Nitrogen is a colorless, tasteless, odorless
gas, as we should expect from the fact that air possesses these proper-
ties. When sufficiently cooled, it forms a colorless liquid, boiling at
— 194°. By further cooling, this liquid freezes to a white solid, which
melts at —214°. One liter of pure nitrogen weighs 1.2507 grams.
The solubility of nitrogen in water : 1.6 volumes in 100, is less than
that of oxygen.
Chemical Properties. — The density of the gas shows the formula
of free nitrogen to be N2.
Nitrogen unites with few common chemical elements directly. At
ordinary temperatures it is almost absolutely indifferent. When
passed through a tube over strongly heated lithium, calcium, magne-
sium, or boron, it forms definite chemical compounds, known as
nitrides, in which it is trivalent. These have the formulae Li8N, Ca3N2,
Mg3N2, and BN respectively. When the gas is mixed with oxygen or
hydrogen, and the mixture is heated, little chemical action takes place.
When sparks from an induction coil are passed between platinum
NITROGEN AND ITS COMPOUNDS WITH HYDROGEN 417
wires through the mixtures, small amounts of nitrogen tetroxide
N204, and ammonia NHS, respectively, are produced. The indifference
of free nitrogen is doubtless due to the fact that its molecules (N2) are
extremely stable.
One case of direct union of nitrogen is of economic importance.
The supply required by plants is obtained partly from nitrogen com-
pounds contained in fertilizers, or equivalent substances already pres-
ent in the soil, and partly from ammonium nitrite and nitrate, which
are washed down from the air by the rain. It appears, however, that
plants belonging to the order leguminosce, such as peas, beans, clover,
etc., may be associated constantly with certain bacteria which flourish
in nodules upon their roots. These bacteria have the power of taking
free nitrogen from the air, which penetrates the soil, breaking up its
molecules, and producing compounds containing nitrogen. The masses
round the roots often contain over five per cent of combined nitrogen
which has been acquired in this way. These compounds are chiefly
albumins, which are afterwards digested and absorbed by the roots of
the plant.
Compounds of Nitrogen and Hydrogen. — The commonest and
longest known of these substances is ammonia NH3, which was first
described by Priestley (1774) and named " alkaline air." Curtius
discovered hydrazine N2H4 in 1889, and hydrazoic acid HN~8 in 1890.
Hydroxylamine NH80, discovered by Lessen in 1865, is similar to
ammonia in chemical behavior.
AMMONIA.
Ammonia is liberated in connection with the putrefaction of animal
matter.
Preparation. — 1. When sparks from an induction coil are passed
through a mixture of nitrogen and hydrogen, small proportions of the
materials unite to form ammonia (see below).
2. When water is added to the nitride of magnesium or calcium
(p. 416), ammonia is given off and the hydroxide of the metal formed :
6H20 -» 3Mg(OH)2 + 2NH3.
3. When parts of animals, particularly the horns, hides, and feathers,
which contain complex compounds of carbon, nitrogen, hydrogen, and
oxygen, are heated strongly, much of the nitrogen is driven out as
418 INORGANIC CHEMISTRY
ammonia. Hence the name spirit of hartshorn, applied to the aqueous
solution.
4. Coal often contains as much as 2 per cent of combined nitrogen,
and when heated in absence of air, gives off much ammonia. The
entire commercial supply is obtained as a by-product from operations
like the manufacture of illuminating-gas and of coke, in which the
destructive distillation of coal takes place. The crude mixture of gases
passes first through water, in which some of the tar is condensed and
most of the ammonia dissolves. This ammoniacal liquor is heated with
a little slaked lime, and the escaping gas is led into dilute hydrochloric
or sulphuric acid, with which it combines to form the chloride or sul-
phate of ammonium.
5. In the laboratory, a mixture of slaked lime and some salt of
ammonium, such as ammonium chloride, either with or without water,
is heated in a flask or retort provided with a delivery tube :
Ca(OH)2 + 2NH4C1 £5 CaCL, + 2NH4OH,
By the double decomposition usual with ionogens, ammonium hydroxide
is formed, and this, being unstable, immediately breaks up into water
and ammonia.
6. Warming the aqueous solution also gives a steady stream of the
'gas. Since the gas is very soluble in water, it must be collected over
mercury or by upward displacement of air. The moisture which it may
contain is removed by passage through a tower or wide tube filled with
quicklime. Calcium chloride cannot be used, as the gas combines with
it forming a compound CaCl2, 8NH3, similar in properties to a hydrate
(pp. 120-123).
Physical Properties. — Ammonia is a colorless gas with a pun-
gent, characteristic odor familiar in smelling-salts. The G.M.V. of the
gas weighs 17.29 g., so that the density is little more than half that of
air (cf. p. 215). When liquefied it boils at — 34° and exercises a pressure
of 6.5 atmospheres at 10°. The solid is white and crystalline and melts at
— 77°. The gas is very soluble in water, one volume of the latter dissolv-
ing 1148 volumes of the gas at 0°, 764 volumes at 16°, and 306 volumes
at 50°. The 35 per cent aqueous solution, saturated at 15°, and sold
as " concentrated ammonia," has a sp. gr. 0.882. The whole of the
dissolved gas may be removed by boiling (cf. p. 379).
NITROGEN AND ITS COMPOUNDS WITH HYDROGEN 419
Liquefied ammonia is used in refrigeration. The gas is liquefied by
compression, and the heat which is thus liberated is removed by flowing
water which surrounds the pipes. The liquid then flows into other
pipes immersed in calcium chloride brine, and is there allowed to evapo-
rate, the gas returning to the compressor. The heat of vaporization,
260 calories per gram, is taken from the brine, which is thus partially
•frozen. The resulting freezing mixture of ice and calcium chloride
solution is then distributed to the localities to be cooled. The ammonia
and brine remain within their respective closed systems of pipes,
and are used over and over again. In many cases the cooling by the
liquefied ammonia is used directly, and the intermediate freezing mix-
ture is dispensed with.
Chemical Properties. — The discharge from an induction-coil
decomposes ammonia (to the extent of 94-98 per cent) into nitrogen
and hydrogen. Under the same circumstances, union of the constitu-
ents also occurs (up to 2-6 per cent) :
2NH3 <=> N2 + 3H2.
The behavior of a system in chemical equilibrium may be illustrated by
inclosing dry ammonia in a tube over mercury and allowing sparks to
pass between platinum wires (Fig. 85). After a time
the volume is practically doubled, every two molecules
of ammonia giving four of the products (cf. p. 208).
When, now, a little sulphuric acid is admitted above
the mercury, the remaining trace of ammonia combines
with the acid and is removed. This disturbs the equil-
ibrium. As sparks continue to pass, the action now
proceeds steadily backwards. The ammonia which is
formed is no longer subject to decomposition, for, once
it has combined with the acid, it never returns to the
neighborhood of the discharge. Thus the whole of the
gases finally combine :
(NH4)2S04 <- H2S04 + 2NH3
3H2,
and ammonium sulphate, dissolved in the excess of acid, pIG. 85.
alone remains. The action, therefore, first goes almost
completely in one direction, and then quite completely in the other,
while no change has taken place in the conditions to which the gas
is subjected at the point where the interaction is occurring. The sole
420 INORGANIC CHEMISTRY
difference is that a little of an acid has been introduced into a rela-
tively remote part of the space.
When passed over reducible oxides, such as heated cupric oxide,
ammonia is oxidized to water and free nitrogen :
3CuO + 2NH3 -* 3Cu + 3H2O + Na,
and burns in pure oxygen with the same result. In air, heat is used
up, not only in decomposing the compound, but also in raising the
temperature of the nitrogen of the atmosphere. This causes a contin-
uous drain on the heat given out by the action, and prevents the main-
tenance of the kindling temperature (p. 73).
When dry ammonia is passed over heated potassium or sodium, one
unit of hydrogen is displaced :
2K + 2NH3 -> 2KNH2 + H2,
and a solid of metallic appearance remains behind. Substances con-
taining the group — NH2 are called amides, and this one is therefore
named potassamide.
Chlorine and bromine combine with the hydrogen and liberate the
nitrogen of ammonia. This action may be used for obtaining a stream
of nitrogen, provided excess of chlorine is avoided (see Nitrogen tri-
chloride, below). Chlorine is led into a solution of ammonium chloride :
2NH4C1 + 3C12 -> Na f + 8HC1.
The most characteristic property of ammonia is its power to form
a base by combination with water :
NH3 (gas) <=> NH3 (diss'd) + H2 0 <=± NH4OH <=> KH4' + OH'.
Probably only a small proportion of the gas is actually combined at
any one time, the greater part being simply dissolved.
The gas unites also with acids, forming salts (cf. p. 337), which, in
solution, are highly ionized *
NH3 + HC1->NH4C1,
NH
Ammonium Compounds. — Since N"H4 plays the part of a metal,
entering into the composition of a base and of a series of salts, it is
named ammonium. It constitutes the positive ion of these compounds.
As this radical forms a univalent ion and gives a distinctly alkaline
base, it is classed with potassium and sodium as one of the metals of
the alkalies -v..
NITROGEN AND ITS COMPOUNDS WITH HYDROGEN 421
Ammonium hydroxide, although much less completely ionized
than potassium hydroxide, affects litmus easily. In a normal solution
about 0.4 per cent of the ammonia is in the form of ammonion,
NH4*. When an acid is added to the solution, the correspondingly
small amount of hydroxidion which exists in it is removed and the
various equilibria are displaced forwards. The final result is the same
as with any other base :
NH8 (diss'd) + H20 ?=* NH4OH <=; NH4' + OH' ) <_ „ Q
2HC1 ±5 2C1' + 2H' ] ~~
When strongly heated, all ammonium salts are decomposed and, usually,
give ammonia and the acid. When the latter is volatile, the whole
material of the salt is thus converted into gas. If the acid is volatile with-
out permanent decomposition, it reunites with the ammonia when the
vapor is cooled :
NH4C1 <=> HC1 + NH8.
This behavior distinguishes ammonium salts from those of the typical
metals, for, with the exception of mercury salts, most other salts are not
easily and completely volatilized. The use of ammonium chloride (sal-
ammoniac) in soldering depends on the dissociation of the salt, by the
heat of the iron, and the action of the liberated hydrochloric acid on
the oxide which covers the surface of the metal to be soldered.
Ammonium salts are recognized by warming them, dry or in solu-
tion, with a base :
2K.
when the odor of ammonia becomes noticeable. When the solution is
used, it is the tendency of the NH4* and OH' to unite to form the slightly
ionized molecular hydroxide that sets the other equilibria in motion.
The principle at the basis of the change is thus the same as in neu-
tralization (p. 354).
In ammonia, nitrogen is trivalent, while in the salts it appears to
be quinquivalent :
H\ H\ H H\ H
H-N H-N/ H-NX
H/ H/ XOH H/ ^Cl
422 INORGANIC CHEMISTRY
HYDRAZINE, HYDRAZOIC ACID, HYDROXYLAMINE.
Hydrazine. — By reduction of a compound of nitric oxide and
potassium sulphite by means of sodium amalgam,* a solution of hydra-
zine hydrate is obtained :
K£03, 2NO + 6H -> N2H4, H20 + K2S04.
The same substance is more easily made from certain organic deriva-
tives. When the hydrate is distilled with barium oxide, under reduced
pressure, hydrazine is liberated :
N2H4, H2O -(- BaO -> N2H4| + Ba(OH)2.
Hydrazine is a white solid, which fumes in moist air, giving the
hydrate once more. It melts at 1.4° and boils at 113.5°.
Hydrazine hydrate freezes at about —40°, boils at 118.5°, and can
be distilled without decomposition. Its aqueous solution is alkaline,
and salts can be formed by neutralization.
Hydrazoic Acid. — When nitrous oxide (q.v.} is led over sodamide
at 200°, water is liberated and sodium hydrazoate remains behind :
NH2Na + N20 -> NaN3 + H20.
A dilute solution of the free acid is best obtained by distilling the lead
salt with dilute sulphuric acid. A similar solution may be made more
directly by adding cold nitrous acid (q.v.) to a cold aqueous solution
of hydrazine hydrate :
NHOH + HN0 -* HN + 3H0.
25
By repeated distillation of the solution the pure acid is obtainable. It
boils at 37°. The operation is a dangerous one, as the pure acid is
violently explosive, resolving itself into nitrogen and hydrogen with
liberation of much heat :
HN8, Aq -> H + 3N + Aq + 61,600 cal.
It is an acid of somewhat greater activity and degree of ionization than
acetic acid. Active metals, like magnesium, displace hydrogen from
its solution. Its silver salt (AgN3) is insoluble in water. It neutral-
* The sodium dissolved in the mercury interacts with the water, giving hydro-
gen (see Active state of hydrogen, below).
NITROGEN AND ITS COMPOUNDS WITH HYDROGEN 423
izes ammonium hydroxide and hydrazine hydrate, giving two salts,
jSTH4N"3 and N2H5N3. These constitute two additional compounds of
nitrogen and hydrogen, but differ from ammonia and hydrazine in
being ionogens.
Hydroxylamine. — Tin displaces hydrogen from dilute hydro-
chloric acid : Sn + 2HC1 — » SnCL, + H2, and this combination forms
a reducing agent (see below). When dilute nitric acid (q.v.} is added
to the mixture, a considerable part of it is reduced to hydroxylamine :
HN08 + 6H -» NH80 + 2H2O.
The hydroxylamine forms a weak base, XH4O.OH, with water which
interacts with the excess of acid, giving hydroxylamine hydrochloride,
NH4OC1. By more complete reduction of part of the nitric acid, some
•ammonium chloride is formed at the same time. To secure the salt
of hydroxylamine, the tin ions are removed by means of hydrogen sul-
phide, which precipitates stannous sulphide. The filtered solution is
then evaporated to dryness, the hydroxylamine hydrochloride is
extracted from the residue with absolute alcohol, and this alcoholic
solution is finally evaporated in turn. The hydrochloride is a white
crystalline salt.
When the hydrochloride is treated with a base, in the absence of
water, and the mixture is distilled under reduced pressure, hydroxyla-
mine passes over. It is a white solid melting at 33° and boiling at
58° at 22 mm. pressure. Even before melting (above 15°), it begins to
decompose, and explodes at or below 130°. In chemical behavior it is
like ammonia. With water it forms a base which combines with acids,
but is less active than ammonium hydroxide. It is a stronger reducing
agent than ammonia, precipitating silver from a solution of silver nitrate.
The union with acids indicates that the molecule of hydroxylamine
is unsaturated, and hence the nitrogen unit is supposed to be trivalent :
H\ H\ H
H-K H-NX
H-0/ H-0/ XC1
An Active State of Hydrogen (Nascent Hydrogen). — Pure
hydrogen gas, whatever its source may have been, shows conspicuous
reducing powers only when heated (cf. p. 109). But when the gas is
absorbed in platinum or palladium (p. 107), the two together reduce
many substances in the cold, although the platinum and palladium
themselves seem to take no part in the change. They are simply
424 INORGANIC CHEMISTRY
catalytic or contact agents, and in theiv presence the hydrogen is more
active. Similarly the hydrogen liberated by the electrolysis of a dilute
acid reduces substances added to the liquid surrounding the cathode
(cf. p. 304), and shows different degrees of activity according to the
material, platinum, carbon, or the like, of which the cathode is made.
Finally, zinc, tin, and other metals with dilute acids give hydrogen,
and this can produce reductions so long as it is in the immediate
neighborhood of the metal (cf. pp. 422, 423, and see Arsine). Thus, as
we have just seen, nitric acid is reduced to hydroxylamine when it is
mixed with the materials which generate the hydrogen. The hydrogen
on the platinum plate forming the cathode of an electrolytic cell is
equally able to reduce nitric acid. The very same hydrogen, however,
if led by a tube from a Kipp's apparatus or electrolytic cell into a
second vessel would be then quite inactive and incapable of affecting
the same nitric acid. Hydrogen, therefore, like other substances, may
show a greatly increased speed of interaction when in immediate con-
tact with a suitable body. This more active state of hydrogen is de-
scribed as the nascent state, because it happens to be a common con-
dition of hydrogen when associated with substances which produce it.
This state has, however, no necessary connection with such an imme-
diately preceding act of liberation, as our first example shows.
Some theorists apply the atomic hypothesis to this phenomenon, and suggest
that it is the freshly liberated atomic hydrogen (H) which possesses this greater
activity. This would require us to believe that nascent hydrogen was a different
substance from free hydrogen, which may or may not be true, but is certainly im-
probable. If the hypothesis of atomic hydrogen were correct, all arrangements that
give free hydrogen under similar circumstances (e.g. electrolytic cells with different
electrode-materials) should show equal reducing powers. But it is notorious that
they do not. Different substances are known, however, to show different efficiencies
as contact agents.
The atomic theory of nascent action is often used to explain the oxidizing
activity of hypochlorous acid (cf. p. 271). But it is less often employed to explain the
oxidizing power of sulphuric acid or other oxidizing agents. Yet, logically, it
ought to be applied even to double decomposition, if it is used at all. If this
were done, salt and sulphuric acid would be assumed to give nascent chlorine and
nascent hydrogen, respectively, since they yield hydrogen chloride in the cold,
while free hydrogen and chlorine do not.
Halogen Compounds of Nitrogen. — When ammonium chloride
solution is treated with excess of chlorine, drops of an oily liquid, nitro-
gen trichloride, are formed :
4HC1.
NITROGEN AND ITS COMPOUNDS WITH HYDROGEN 425
It is extremely explosive, resolving itself into its constituents with
liberation of much heat.
There appear to be several compounds related to ammonia and con-
taining iodine. When a solution of iodine in potassium iodide solu-
tion (p. 156) is added to aqueous ammonia, a brown precipitate is
formed. This seems to have the composition N2H3I8, and is named nitro-
gen iodide. It may be handled while wet, but when dry decomposes into
its constituents with violent explosion if touched with a feather.
Exercises. — 1. When moist air is used as a source of nitrogen, what
advantage is there in using copper rather than the less expensive
metal iron, for removing the oxygen (p. 416) ?
2. How many grams of water at 0° could be frozen (p. 115) by the
removal of the heat required to evaporate 50 g. of liquid ammonia
(p. 419) ?
3. How many grams of ammonia are contained in 1 1. of " concen-
trated ammonia" (p. 418) ?
4. What are the ions of hydrazine hydrate (p. 422) ? Formulate
(p. 454) the neutralization of this base with sulphuric acid.
5. What is the object attained by distilling under reduced pressure
in making hydrazine (p. 422) and hydroxylamine (p. 423) ?
6. Classify (p. 187) the interaction of a nitride with water (p. 417),
and of chlorine and ammonium chloride (p. 420), and the results of
heating ammonium nitrite (p. 416) and ammonium chloride (p. 421).
CHAPTER XXV
THE ATMOSPHERE. THE HELIUM FAMILY
WE have seen that to counterbalance the pressure of the air a
column of mercury of the same diameter averaging 760 mm. in height
is required. Let th'e section of the column be 1 sq. cm. Then the
pressure of a column of air 1 sq. cm. in section, and extending so
far from the earth as any downward tendency of the air exists, is
equal to the weight of a column of mercury containing 76 c.c. of the
metal. The weight of 1 c.c. of mercury being 13.6 g., this volume of
the metal weighs 1033.6 g. This number represents therefore the
pressure which is exerted by the air upon each square centimeter of
the earth's surface. In ordinary units of measure, this is nearly fif-
teen pounds to the square inch.
A more vivid appreciation of the reality of this pressure may be obtained by
noticing one of its effects. By boiling a small quantity of water in a tin can
furnished with a narrow opening, we remove the whole of the air from its interior,
displacing it by steam. While the boiling is in progress, we suddenly close the
opening with a tightly fitting cork and remove the burner. While the steam was
still issuing from the opening, its pressure was practically that of the atmosphere,
and the can was subject to the same pressure iiiside and out. With the removal
of the flame, however, the steam condenses, and the pressure on the interior is
reduced to a minute fraction of its original value, while the pressure on the
exterior is still the same (1 atmosphere). Under this pressure a vessel of ordinary
tin-plate completely collapses.
Components of the Atmosphere. — The first component of the
air to be recognized and studied was oxygen. When a msting metal or
burning body of any kind removes the oxygen by combining with it, a
gas remains which represents about four-fifths of the original volume.
This residual gas is mainly nitrogen. It contains, however, small pro-
portions of several inert gases, of which the most plentiful is argon.
Examination before the removal of the oxygen also reveals the presence
of varying proportions of carbon dioxide, water vapor, and ammonium
nitrate. In the neighborhood of cities the air likewise contains sulphur
dioxide, hydrogen chloride, hydrogen sulphide, and other ingredients
which may be described as accidental. There are thus three classes
426
THE ATMOSPHERE. THE HELIUM FAMILY 427
of substances in the air. Those of the first class, oxygen, nitrogen, and
argon, are present in almost constant quantities ; those of the second
class are very variable in quantity, although found in all samples of
air ; those of the third class are accidental. Finally, one significant
component of the air is the dust which a powerful beam of light reveals
as it passes through the atmosphere of a darkened room.
Components which are Constant in Amount. — The determi-
nation of the oxygen by burning phosphorus in air, and measuring the
residual gas (p. 416), is not capable of application in an exact manner.
It is better to use a large amount of phosphorus in the form of thin
wire. In this way a great surface is obtained, and the absorption of
oxygen from a sample of air may be carried out in a few seconds.
This method gives fairly accurate results, since there is no time for any
appreciable change in the temperature or pressure of the atmosphere
during the experiment. The passage of purified air over heated copper
(p. 416) has also been used for the same purpose. When this method
is employed, the volume of the nitrogen and argon which survives the
action of the copper is measured, while the increase in weight of the
copper, through formation of cupric oxide, gives the weight of the
oxygen which was originally mixed with it.
Still another method, which is in constant employment in the analy-
sis of mixtures of gases, may also be applied to the air. It consists in
mixing a measured volume of air with an excess of hydrogen, reading
off the volume of the whole, and then exploding the mixture by the
passage of an electric spark. A tube like that in Fig. 44 (p. 125)
may be used. After the explosion, the steam which has been formed
condenses. The contraction gives the volume of the gases which have
disappeared through the explosion, and of this one-third is oxygen and
two-thirds hydrogen (cf. p. 125). When, for example, a contraction
of 25 c.c. of gas has occurred, we know that one-third of this volume
(namely, 8.3 c.c.) is the volume of oxygen which the measured sample
of air contained.
In the air taken from mines, from mountain tops, from the surface
of the sea, and from inland regions, the proportion of oxygen to the
residual gas is found to be fairly constant, although easily perceptible
differences are noted. The percentage of oxygen in dried air ranges
between 20.26 and 21.00, the latter being the proportion in normal
air.
When the residual gas is led slowly through a heated tube con-
428 INORGANIC CHEMISTRY
taining magnesium, the nitrogen unites with the metal to form the
solid nitride (p. 416), and only about 10 c.c. out of every liter remains
uncombined. This residuum is argon, mixed with one-hundredth of
its volume of other gases belonging to the same family (see below).
Exact measurement by volume gives 78.06 per cent of nitrogen and
0.94 per cent of argon in dried air.
It is possible that a trace of hydrogen (>/. p. 92) is one of the
regular components of air.
Components which are Variable in Amount. — Pure. country
air contains about 3 parts in 10,000 of carbon dioxide. In city air
there are from 6 to 7 parts in the same volume, while in the air of
audience-rooms, where the ventilation is defective, the proportion may
rise as high as 50 parts.
The simplest way of showing the presence of carbon dioxide in the
air is by exposing a solution of barium hydroxide in a shallow vessel.
After a short time a layer of barium carbonate forms upon the surface,
in consequence of a chemical change represented by the equation,
Ba(OH)2 + C02 — > BaC03 J, + H20. The same action may be utilized
for the purpose of quantitative analysis. A measured volume of air is
bubbled slowly through a measured volume of a solution of barium
hydroxide of known concentration, and the quantity of barium hy-
droxide remaining is determined by titration (p. 352).
The sources of the carbon dioxide in the air are numerous. It
comes from the decay of vegetable and animal matter, in which,
chiefly through the influence of minute vegetable organisms, the
carbon is oxidized to carbon dioxide. It is formed also by the com-
bustion of coal and wood, and is exhaled by animals. The proportion
of this gas in the air would naturally increase continuously, though
slowly, as the result of these processes, were it not that it is removed
just as continuously by the action of growing plants (see Chap, xxviii).
Thus, there is no reason to suppose that the proportion of carbon
dioxide in the air will be altered appreciably, the two kinds of opera-
tions probably balancing one another approximately.
The quantity of water vapor in the air is constantly changing. It
increases locally by evaporation from the soil and from natural waters,
particularly in warm weather. It decreases when local cooling leads
to the precipitation of water in the forms of mist and rain. The
phrase commonly heard, that on a moist day the atmosphere is " laden "
with moisture, is peculiarly inapt. We recognize at once from obser-
THE ATMOSPHERE. THE HELIUM FAMILY 429
vation of the barometer, which is lower in such a state of the atmos-
phere, that the pressure of the air is less. Moist air must be lighter
than dry air, for in it a certain proportion of water molecules, of molec-
ular weight 18, is substituted for an equal number (cf. p. 190) of mole-
cules of nitrogen, and oxygen whose relative weights are 28 and 32
respectively. The result is therefore a diminution in the specific
gravity of the air. The proportion of water in a given volume of air
may be measured most accurately by permitting the air to stream
slowly through tubes filled with calcium chloride or phosphoric anhy-
dride (Fig. 36, p. 100). The increase in weight of the charged tubes
represents the quantity of moisture abstracted from the sample.
The ammonium nitrate arises from the interaction of nitric acid
and ammonia. The latter is formed by the decay of animal matter
(p. 417) ; the former by the union of nitrogen and oxygen during
thunder-storms. The electrical discharges produce nitrogen tetroxide
(see p. 439), which with water gives nitric acid (q-v.~).
The dust varies both in kind and quantity according to the locality.
It is found to be partly inorganic. In the country the nature of the
inorganic material depends upon the prevailing components of the
soil, and in factories the dust may consist of minute particles of glass,
steel, cement, or other substances. The organic dust may be divided
into two kinds. The part which is dead includes coal dust, refuse
from the streets, minute shreds of cotton, linen, hay, etc. The living
dust consists of pollen grains, spores of fungi and other plants, germs
of infusoria, and bacteria. The presence of such germs in the air is
shown by the fact that when nutritive liquids have been exposed to
the air, even for a few minutes, putrefaction very soon sets in. Some
of these germs also produce disease when they gain access to the
body, particularly through wounds, or incisions made in the course of
operations. The object of antiseptic treatment by the use of phenol
(carbolic acid), mercuric chloride, and other substances, is to destroy
such organisms or to hinder their development.
Flasks can be filled with dustless air through the displacement of
that which they contain by air drawn through a wide tube packed with
12-15 inches of cotton. It has been shown by Aitken that air of this
kind behaves differently from ordinary air in respect to the way in
which its moisture condenses.
If a sample of moist air is cooled until it contains more water vapor
than it could take up at the existing temperature, the excess of moist-
ure is deposited. This deposition usually takes place by the forma-
430 INORGANIC CHEMISTRY
tion of a multitude of little particles of liquid water, which together
make up a fog. Now dustless air lacks this property entirely. When
saturated with water and then cooled, it does not give any trace of fog.
The excess of moisture is gradually deposited upon the walls of the
vessel and upon any material objects which it contains, but of fog
there is no trace visible. It seems that the particles of dust are re-
quired as nuclei round which the water may gather. In the absence
of dust, and therefore of proper nuclei, the moisture is not precipi-
tated in the usual way. Thus fogs and rain would be impossible but
for the presence of dust in all ordinary air.
By diluting air with dustless air, generating fog in the mixture,
and, with the help of a microscope, counting the globules when they
settle, an estimate of the number of particles of dust in air may be
made. It is found that rain removes a large proportion of them,
while respiration and combustion greatly increases their number. The
prevalence of fogs in cities is thus accounted for.
NUMBEE OF
DUST PABTICLES
IN 1 c.c.
Outside, raining 32,000
Outside, fair . . . 130,000
A room 1,860,000
A room, near the ceiling 5,420,000
Air above Bunsen flame 30,000,000
Air a Mixture. — Since the main components of air were not defi-
nitely identified until the end of the eighteenth century, we can under-
stand why the substance was for long considered to be an element.
The experiments which we have described, in which the oxygen was
removed from the air and the nitrogen remained, do not prove that
the original constituents were present simply in mechanical mixture.
They might have been combined, and the combustion of phosphorus,
for example, might have represented the removal of oxygen from com-
bination with nitrogen and its appropriation by the phosphorus. It
may be well, therefore, to point out some reasons which lead us to
regard the air as a mixture :
1. When oxygen and nitrogen are mixed in the proper proportions,
we obtain a gas identical with air in all its properties, and there is no
evidence of any production or absorption of heat, such as would occur
in case of chemical combination.
THE ATMOSPHERE. THE HELIUM FAMILY 431
2. The proportion by volume in which the gases are found in the
air is not so simple as the proportions which we observe in cases of
chemical combination. The proportion is close to 4 : 1, but not exactly
4 : 1. Besides, as we have seen, the proportion is not perfectly constant.
3. The composition of air varies, while the composition of definite
chemical substances is always the same. The proportions by weight
also in which the components are contained in air are not integral
multiples 'of the atomic weights.
4. When two substances enter into chemical combination, the new
body invariably has different physical properties from either of the
original ones. Thus, there is no simple relation between the refrac-
tive power- for light which a compound possesses and the refractive
powers of its constituents. In the case of air, however, the refrac-
tive power is exactly that which we should calculate from the refrac-
tive powers of the constituents, taking into account the proportions of
them which air contains. The same relation holds for all ordinary
physical properties of air. For example, the nitrogen and oxygen
dissolve independently in water in proportion to their solubilities and
partial pressures (p. 155). If the air were a compound, it would dis-
solve as a whole, and the relative proportions of the components would
not be changed by the process.
Composition of Air. — Air, when freed from carbon dioxide and
water, contains by volume 78.06 per cent of nitrogen, 21.00 per cent
of oxygen, and 0.94 per cent of argon. When only the water is
removed, the carbon dioxide averages about 0.03 per cent of the
whole.
Graham has suggested an illustration which will make those pro-
portions clearer. He says that if we imagine the air to be divided by
magic into its components, and to remain separated for a time suffi-
ciently long to enable us to note the proportions, and if the substances
arrange themselves in the order of their specific gravities, we should
have the following layers resting upon the surface of the earth and
xipon one another : On the earth, five inches of water ; above that,
thirteen feet of carbon dioxide ; above that, a mile of oxygen ; and on
the top, about four miles of nitrogen. This would be on the assump-
tion thaf these gases were compressed so as to have the same density
throughout. The recent discovery of argon shows that we should add
to this illustration a layer of argon, of about ninety yards thickness,
between the carbon dioxide and oxygen.
432 INORGANIC CHEMISTRY
Air and Health. — As human beings we have an especial interest
in the composition of the air, since our living depends upon the oxygen
which we secure by breathing it. We draw about half a liter of air
into our lungs at each breath, or about half a cubic meter per hour.
The oxygen of this air is partly used, being taken up by the blood,
and part remains in the exhaled air. On the other hand, carbon dioxide
is given off in the lungs and passes out with the unused oxygen. The
nitrogen is unaffected. In 100 c.c. of expired air there are contained
about 15.9 c.c. of oxygen and 3.7 c.c. of carbon dioxide. The total
quantity of oxygen consumed during twenty-four hours is about three-
fourths of a kilogram, or more than half a cubic meter. While the
greater part of this gains access to the body through the lungs, more
or less exchange of gases takes place in all animals through the skin.
The lower limit of oxygen for respirable air is about 10 per cent,
although a candle is extinguished if the proportion falls below 16.5 per
cent. The union of oxygen or carbon dioxide with the haemoglobin
of the red blood-corpuscles is a reversible reaction. In the tissues,
where the concentration of oxygen is small and that of carbon dioxide
large, the former leaves the haemoglobin and the latter combines with
it. Conversely, in the lungs, where the concentration of oxygen is at
a maximum, and that of carbon dioxide at a minimum, the reverse
changes occur.
Liquefaction of Gases. — The earliest experiments of this kind
seem to have been made by Northmore (1805), who liquefied chlorine,
hydrogen chloride, and sulphur dioxide. In 1823 chlorine was again
liquefied by Faraday; and in the same year Davy, whose assistant
Faraday was, liquefied hydrogen chloride. During the following years
Faraday reduced other gases — sulphur dioxide, hydrogen sulphide,
carbon dioxide, nitrous oxide, cyanogen, and ammonia — to the liquid
condition.
The method which he employed was extremely simple. He used a
bent tube shaped like an inverted v (A), into one limb of which materials
for producing the gas were introduced. The other limb was then
sealed up and immersed in a freezing mixture. The gas, usually
liberated by heating, was liquefied by its own pressure in the cold
limb. By means of a more elaborate apparatus, Cailletet and Pictet
simultaneously (December, 1887) obtained, the one, a fog, and the
other, a spray containing droplets of liquid oxygen. In 1883 Wrob-
lewski and Olszewski made larger amounts of the same liquid. About
THE ATMOSPHERE. THE HELIUM FAMILY
433
the same time Dewar devised means of manufacturing large quantities
of liquid air and oxygen.
The principle now used in liquefying gases depends on the fact
that a perfect gas, when expanding into a vacuum, should suffer no
fall in temperature, since it does no work, while ordinary gases
do become cooled very slightly. The work which they do in expand-
ing in such circumstances is done in over-
coming the cohesion between their mole-
cules (p. 131), and occasion for it arises
from the fact that a tearing apart of the
substance, which consumes heat, has to
take place. Since this cohesion becomes
more conspicuous the lower the tempera-
tiire (cf. p. 135), the cooling effect of ex-
pansion becomes greater and greater as the
temperature falls.
By ingenious arrangement of the appa-
ratus, the low temperature produced by the
expansion of the compressed gas is used to
cool fresh portions of the gas which have
not yet been expanded. The tube contain-
ing the compressed gas is placed concen-
trically within another tube, and the
stream of gas after it has expanded is
turned back upon its course and passes
through the annular space along the tube
containing the uiiexpanded gas. This
double tube is many yards long, so that the
most perfect contact and complete interchange of heat can occur. Thus,
by a process of intensification, the temperature of the unexpended gas
is continually depressed, and the temperature after expansion, there-
fore, becomes lower still. Finally, the latter temperature becomes
low enough to produce liquefaction of the compressed gas at the
pressure to which it is subjected, and the jet of expanding gas is filled
with drops of liquid. These are caught in a recess of the vessel, and
may be removed from time to time through a stopcock.
The most successful apparatus for use on a small scale is that de-
vised by Hampson (Fig. 86). The compressed gas enters by the small
tube near the top, and eventually issues by the wider one. The inner
tube is coiled in a small cylinder to secure compactness, and a spiral
FIG. 86.
434
INORGANIC CHEMISTRY
partition between the coils produces the outer tube of which we have
spoken. The gas in the tube A (Fig. 87) is under a pressure of 150-
200 atmospheres. The distance of the nozzle Z> from the plug C is
adjusted so that the pressure of the gas in the
chamber and spiral outer tube is reduced to one
atmosphere.
For handling liquefied gases, a form of appara-
tus devised independently by Weinhold and Dewar
is employed (Fig. 88). It consists of a double
flask in which the space between the inner and
outer bulbs has been exhausted by means of an
air-pump. To prevent by reflection the access of
radiant heat to the interior, the surfaces of the
flasks are often silvered.
FIG. 87.
Liquid Air. — Liquid air varies in composition,
as the nitrogen (b.-p. — 194°) is less condensible
than the oxygen (b.-p. — 182.5°). It boils at about
- 190°, and contains about 54 per cent of oxygen
by weight, while air contains 23.2 per cent. By
allowing evaporation to go on, a liquid containing
75 to 95 per cent of oxygen is easily obtained (c/.
p. 63). The gas secured by
the evaporation of the residue
is pumped into cylinders and sold as compressed
oxygen. Cartridges made of granular charcoal
and cotton waste, when saturated with liquid air,
have been used as an explosive in mining.
•
THE HELIUM FAMILY.
Argon. — Lord Kayleigh was the first to ob-
serve that, while specimens of oxygen and other
gases made purposely from various sources always
had the same density, nitrogen was an exception.
One liter of nitrogen made from air, and supposed to be pure, weighed
1.2572 g. When the gas was manufactured by decomposition of five
different compounds, such as urea and certain oxides of nitrogen, the
results agreed well amongst themselves. The mean weight of a liter
of this nitrogen was only 1.2505 g. The difference, amounting to nearly
pIG.
THE ATMOSPHERE. THE HELIUM FAMILY 435
7 mg., was very much greater than the experimental error. The sus-
picion naturally arose that some heavier gas was present in natural
nitrogen. Soon after (1894), Professor, now Sir William Ramsay
obtained argon by removal of the greatly preponderating nitrogen by
means of magnesium (p. 427). The new gas had a molecular weight of
about 40, and was therefore more than one-third heavier than nitrogen.
In order to make sure that this substance did not have its source in the magne-
sium, a different method was used by Lord Kayleigh to separate it from nitrogen.
He inclosed the nitrogen with a sufficient quantity of oxygen in a flask, through
the sides of which platinum poles had been inserted. A tube entered the flask by
the neck, and through this a constant fountain of potassium hydroxide solution
played upon the interior and kept the surface covered with fresh quantities of the
liquid. Another tube permitted the overflow of the excess of this solution. The
discharge of electricity produced nitrogen tetroxide (q.v.), which was absorbed by
the potassium hydroxide to form potassium nitrate and potassium nitrite. The
volume of gas thus continually diminished, and, by persistent sparking of the mix-
ture with oxygen, the nitrogen was finally all taken out. The excess of oxygen
was then removed, and the gas which remained was found to be identical with that
which Kamsay had obtained.
Lord Rayleigh's method was extremely interesting, since it was a reproduction
of an experiment made by Cavendish towards the end of the eighteenth century.
The latter had remarked that the assumption that the inert atmospheric gas was a
homogeneous single substance had not been confirmed by sufficiently careful ex-
periment. He even endeavored in precisely the above way to remove the nitro-
gen in order to see whether any other body remained. He records the fact that
the residual gas was nothing but a minute bubble, and seems to have dismissed the
subject with the idea that if there was any other constituent of the atmosphere
present in the nitrogen, its amount was exceedingly small. Argon thus narrowly
escaped detection nearly a century before its actual discovery.
The exact density of argon, referred to oxygen = 32, is 39.9.
When liqxiefied it boils at — 186°, and the colorless solid, obtained by
cooling the liquid, melts at — 189.5°. The solubility of the gas in
water (4 volumes in 100) is two and one-half times that of nitrogen. It
has not been found to enter into any sort of chemical combination, and
was named argon on this account (Gk. a/ayos, inactive).
Since the atomic weight of a substance is a quantity showing the
proportion in which it enters into combination, it will be seen that
argon, since it has not yet been found to combine with anything, has,
to speak strictly, no atomic weight (pp. 201, 207). At the same time, it
is manifestly a question of interest to determine whether the physical
properties require the supposition that the molecule of argon contains
one atom, or more than one atom.
If gases were composed of perfectly elastic spheres, the molecules
436 INORGANIC CHEMISTRY
would be altered only in respect to velocity of movement by heating.
Calculation enables us to determine that to raise the temperature of
one G.M.V. of such a gas by one degree would require 3 calories in
every case. Now Regnault found the following values (in calories) for
the heat capacity of gases :
Oxygen (02) 4.96 Carbon dioxide (C02) . 7.56
Hydrogen (H2) .... 4.82 Sulphur dioxide (S02), 7.82
Nitrogen (N2) .... 4.82 Chloroform (CHC13) . 16.55
Hydrogen chloride (HC1), 4.76 Alcohol (C2H6O) . . 18.70
These gases evidently are not constituted as the hypothesis with
which this paragraph opened supposes. Some heat is consumed in
work done inside the polyatomic molecules, and the amounts by which
the numbers exceed 3 calories show that the intramolecular work is
greater as the complexity of the molecules increases. Now, in mercury
vapor the value is exactly 3, and we have already seen that its atomic
and molecular weights are identical (p. 205). According to the molec-
ular hypothesis, its molecules are monatomic, and in them there is no
opportunity for the consumption of heat in intramolecular change.
Hence, when argon was found likewise to give 3 for the value of its
molecular heat-capacity, identity of its atomic and molecular weights
was assumed also.
Helium. — In 1868 Lockyer first detected an orange line in the
spectrum of the sun's prominences which was not given by any terrestrial
substance then known. The line was so conspicuous that it was attrib-
uted to the presence, in considerable quantity, of a new chemical ele-
ment, which was named helium (Grk. 17X105, the sun). Ramsay, in
searching for sources of argon, examined the " nitrogen " which was re-
ported by various mineralogists as being disengaged when certain rare
minerals were heated. These minerals, cleveite, uraninite, and brog-
gerite, were chiefly compounds of uranium, yttrium, and thorium. He
was surprised to find (1895) that the gas was not always nitrogen, nor
was it even argon. It frequently contained a large proportion of a gas,
very much lighter than either, the spectrum of which showed at once that
it was identical with helium. The same gas . has since been obtained
from the water of certain mineral springs, and is found in small amount
in the atmosphere. Helium does not exhibit any tendency to enter
into combination, either with the elements which its parent minerals
contain, or with any others. It is monatomic (cf. p. 435) ; its density
shows that its molecular weight is 4. It has been liquefied by Dewar.
THE ATMOSPHERE. THE HELIUM FAMILY 437
Neon, Krypton, and Xenon. — When the argon obtained from
atmospheric nitrogen is cooled with liquid air ( — 185°), the argon,
krypton, and xenon are liquefied, and the neon and helium are dissolved
by the liquid. When heat is allowed to reach the mixture, the last
two gases escape first, along with much argon. When most of the
argon has escaped, the krypton and xenon still remain liquid. By
repeated liquefaction and fractional evaporation (see under Petroleum),
the krypton and xenon are separated from the argon and from one
another. When the vessel containing the mixture of helium and neon
is immersed in liquid hydrogen ( — 240°), the second freezes to a white
solid, and the helium, which remains gaseous, can be pumped off.
These gases are all entirely inactive chemically, and are all mon-
atomic. Their molecular weights are : Neon, 20; krypton, 81.5; xenon,
128.
Exercises. — 1. A sample of moist air, confined over water at 15°
and 760 mm., occupies 15 c.c. It is mixed with 20 c.c. of hydrogen, and
the mixture is exploded, and suffers a contraction of 9.5 c.c. What
would be the volume of the oxygen it contained if measured dry at 0°
and 760 mm. ?
2. Calculate from the data on pp. 427, 428 and the densities the
percentage by weight of the three principal components of air.
CHAPTER XXVI
OXIDES AND OXYGEN ACIDS OF NITROGEN
THE names and formulas of the oxides and oxygen acids of nitrogen
are as follows :
Nitrous oxide N20 < Hyponitrous acid H^NjOj
Nitric oxide NO
Nitrous anhydride N20, < — > Nitrous acid HN02
Nitrogen tetroxide N204 and N02
Nitric anhydride N205 < — — > Nitric acid HN08.
All the oxides are endothermal compounds, yet, with the exceptions of
the third and the last, they are all relatively stable. The acids, when
deprived of the elements of water, yield the oxides opposite which
they stand. Conversely, excepting in the case of nitrous oxide, the
anhydrides with water give the acids. All of these substances are
obtained directly or indirectly from nitric acid — nitric anhydride by
removal of water, the others by reduction. We turn, therefore, first,
to this acid, its sources and properties.
NITRIC ACID.
Sources. — Sodium nitrate, or Chili saltpeter, is found in a des-
ert region near the boundary of Chili and Peru, and chiefly in the
former country. The deposit is about 5 feet thick, 2 miles wide,
220 miles in length, and contains 20 to 55 per cent of the salt. Puri-
fication is effected by recrystallization. Potassium nitrate, or Bengal
saltpeter, is found in the soil in the neighborhood of cities in India,
Persia, and other oriental countries. It arises from the oxidation of
animal refuse (cf. p. 417) through the mediation of nitrifying bacteria.
The potash and lime in the soil, along with the product of oxidation of
the nitrogen, give nitrates of potassium and calcium. The aqueous
extract of this soil is treated with wood ashes, on account of the potash
(K2C03) contained in them, is poured off from the calcium carbonate
thus precipitated, and is finally evaporated.
438
OXIDES AND OXYGEN ACIDS OF NITROGEN 439
The action of the nitrifying bacteria may be imitated iii a rough way. Air is
caused to pass slowly through concentrated aqueous ammonia, whereby it becomes
mixed with ammonia gas. This mixture is led through a wide tube containing
platinized asbestos and is then discharged into a large flask. When the asbestos is
warmed, it begins to glow, and thereafter the action maintains itself. A part of
the ammonia is oxidized to nitric acid, which combines with the excess of ammonia,
giving ammonium nitrate. This salt forms a cloud which settles in solid form in
the flask.
Preparation. — When any nitrate is treated with any acid, nitric
acid is formed by a reversible double decomposition. As sodium
nitrate is the cheapest salt of nitric acid, it is always employed. For
the same reason, and on account of its activity, and, above all, because
of its relative involatility, sulphuric acid is used to displace it :
NaN08 + KjSO, t> NaHSO4 + HN08f .
The nitric acid is rather volatile (b.-p. 86°), while sulphuric acid (b.-p.
330°) is much less so, and the two salts are not volatile at all. Thus
the interaction proceeds to completion very easily (cf. p. 259). The
materials are heated in cast-iron stills, and the vapor is condensed in
earthenware pipes surrounded by water. In many factories a reduced
pressure is maintained in the stills and condensers, in order that the
distillation may take place at the lowest possible temperature. This
precaution is taken to reduce to a minimum the partial decomposition
of the nitric acid (see below).
Another action by which attempts are being made to manufacture
nitric acid is the direct union of the nitrogen and oxygen of the air
under the influence of an electric discharge. The nitrogen tetroxide
(N02), which is formed in small amounts at a time, is dissolved in
water :
3N02 + H,0 +± 2HN08 + NO.
The nitric oxide gas, on escaping from the water, unites directly with
oxygen to reproduce the tetroxide. The reaction is of interest, inde-
pendently of this one application, because of its reversibility. It pro-
ceeds forward with excess of water, while the reverse action takes
place when nitric oxide (<?.?'.) comes in contact with concentrated nitric
acid in which the quantity of water is at a minimum.
Physical Properties. — Nitric acid is a colorless, mobile liquid,
boiling at 86°, and freezing to a solid which melts at — 47°. It fumes
440 INOKGAN1C CHEMISTRY
strongly when its vapor issues into moist air (of. p. 282). An aqueous
solution containing 68 per cent of the acid boils at 120.5°, While the
pure acid, pure water, and all other mixtures, boil at lower temperatures,
and have, therefore, higher vapor pressures. On this account a more
dilute acid, when heated, loses water until it reaches this strength
(cf. p. 182). The 68 per cent nitric acid forms the " concentrated
nitric acid " of commerce.
Chemical Properties. — 1. Like chloric acid (p. 274), and other
oxygen acids of the halogens, nitric acid is most stable when mixed
with water. The pure (100 per cent) acid decomposes while being
distilled :
4HN03 -> 4N02 + 2H20 + O2,
yet not with explosive violence like chloric acid. The distillate is
colored brown by dissolved nitrogen tetroxide (N02). Repeated dis-
tillation finally leaves 68 per cent of the acid, mixed with 32 per cent
of water formed by the above decomposition. The acid of constant
boiling-point is, therefore, reached, as usual, from more concentrated as
well as from less concentrated specimens.
" Fuming" nitric acid is brown in color, and contains a considerable
amount of dissolved nitrogen tetroxide. It is made by distilling the
acid with a little starch. The latter reduces a part of the nitric acid
and liberates more of the tetroxide than does mere distillation.
2. Nitric acid combines with small amounts of water to form
hydrates HN03,H20 and HNO8,3H20, but they are unstable, and are
decomposed when more water is added. The formula of the former
might be written H3NO4, but no salts corresponding to a tribasic acid
of this constitution are known (see Phosphoric acid).
3. Nitric acid, when dissolved in water, is highly ionized, and gives
a solution containing a relatively large concentration of hydrion. It is
therefore active as an acid. By interaction with hydroxides and oxides
it forms nitrates.
4. When pure nitric acid (b.-p. 86°) is poured upon phosphoric
anhydride, the latter possesses itself of the elements of water, and
distillation of the mixture gives nitric anhydride :
2HN03 + P205 -> N205 1 + 2HP03.
The anhydride is a white solid melting at 30° and boiling at 45°. It
unites vigorously with water to form nitric acid. It cannot be kept, as
OXIDES AND OXYGEN ACIDS OF NITROGEN 441
it decomposes into nitrogen tetroxide and oxygen, 2N205 — » 4N02 -f 02,
with liberation of heat.
5. Like the unstable oxygen acids of the halogens, nitric acid is an
oxidizing agent even when diluted with water. The multiplicity of
the products into which it may be decomposed by reduction, however,
renders separate treatment of this property necessary (see below).
6. Nitric acid interacts energetically with many compounds of
carbon. Thus, when heated with phenol (carbolic acid) it gives picric
acid, which crystallizes in yellow needles in the mixture :
C6H5(OH) + 3HON02 4 C6H2(OH)(N02)S + 3H20.
The presence of water decreases the activity of the molecules. Hence,
in this sort of action, which is not ionic, not only is the most concen-
trated nitric acid employed, but concentrated sulphuric acid is added to
assist in the elimination of the water (cf. p. 388) that arises as one
of the products.
It will be seen that the group N02 has taken the place of hydro-
gen which was formerly attached directly to the carbon of the phenol.
Compounds of this kind are called nitro-derivatives. Picric acid is
trinitrophenol.
7. Organic compounds of another class, the alcohols (q.v.~), interact
with molecular nitric acid in a different way. The latter is mixed with
sulphuric acid with the same object as before. Thus, when glycerine
is added slowly to the cooled mixture, glyceryl nitrate (so-called nitro-
glycerine, see below) is produced :
C3H5(OH)8 + 3HON02 -» C3H5(ON02)3 + 3H20.
Here it is the hydrogen of the hydroxyl groups that is displaced by
N02. The action is not ionic, and the product is not an ionogen.
Gun-cotton is made by this action, prepared cotton (cellulose) being
employed :
(C6H1005)2 + 6HON02 -> C12H14010(N02)6 + 6H20.
8. Nitric acid produces substances of bright-yellow color, known as
xaiithoproteic acids, when it comes in contact with the skin.
The chemical properties of nitric acid are best represented by the
graphic formula :
442 INORGANIC CHEMISTRY
Nitrates. — The nitrates are all more or less easily soluble in
water. When heated they decompose (see below). Sodium nitrate is
used most largely as a fertilizer. Much is employed in sulphuric acid
manufacture, and the rest for conversion into potassium nitrate and in
making nitric acid. Potassium nitrate is used, along with sulphur and
charcoal, in the manufacture of gunpowder. It furnishes the oxygen
with which the charcoal unites ; and potassium sulphide, carbon dioxide,
and nitrogen are amongst the products of the explosion.
NITRIC OXIDE AND NITROGEN TETROXIDE.
•
Preparation of Nitric Oxide. — Pure nitric oxide is obtained
by adding nitric acid to a boiling solution of ferrous sulphate in dilute
sulphuric acid or of ferrous chloride in hydrochloric acid :
2FeS04 + H2S04 -H. Fe2(S04)3 ( + 2H) x 3 (1)
(3H) + HN08 -> NO + 2H20 x 2 (2)
6FeS04 + 3H2S04 + 2HNO8 -» 3Fe2(S04)8 + 2NO + 4HP.
The first partial equation does not take place at all unless an oxidizing
agent like nitric acid is present (p. 307). The multiplication of the
two partial equations by 3 and 2 respectively is required in order that
the hydrogen, which is not a product, may cancel out. This action is
used as a means of determining the quantity of nitric acid in a solution,
or of nitrates in a mixture, by measurement of the volume of nitric
oxide evolved.
As we shall see, this gas may also be obtained when sufficiently
dilute nitric acid (sp. gr. 1.2) acts upon copper. Although some
nitrous oxide and nitrogen are produced in this interaction, it fur-
nishes a convenient method of generating the gas.
Properties of Nitric Oxide. — Nitric oxide is a colorless gas. In
solid form it melts at — 150° and the liquid boils at — 142.4° under 757.2
mm. pressure. Its solubility in water is slight.
The density of the gas shows the formula to be NO ; and there is no
tendency to form a polymer, such as N202, even at low temperatures.
This gas is the most stable of the oxides of nitrogen. At 1700° about
0.5 per cent decomposes into nitrogen and oxygen. At 2000° the
decomposition reaches 1 per cent. These measurements indicate
simply the relative rates at which the resolution into its components
occurs, for the action is not reversible, and no equilibrium is established.
OXIDES AND OXYGEN ACIDS OF NITROGEN 443
Vigorously burning phosphorus continues to burn in the gas, the
heat evolved liberating the oxygen required for the continuation of the
combustion. Burning sulphur and an ignited taper, however, are
extinguished.
Nitric oxide has two characteristic properties. It unites directly
with oxygen in the cold to form the reddish-brown nitrogen tetroxide :
2NO + 02 <=» 2N02.
The same result follows when it is led into warm concentrated nitric
acid (ef. p. 439) :
NO + 2HN03 <=» 3N02 + H20.
It also unites with a number of salts, the compound in the case of
ferrous sulphate being capable of existence in solution and possessing
a brown color. The composition of this compound (a molec-
ular compound, see below) has not been determined, but at
8° the proportion of nitric oxide absorbed by the solution is
about 2NO : 3FeSO4.
Since ferrous sulphate will first reduce nitric acid to nitric
oxide (p. 442), and the excess of the salt will then give a
brown color with the product, a delicate test for nitric acid is
founded upon the above action. The substance supposed to
contain a nitrate is mixed with a strong solution of ferrous
sulphate, and concentrated sulphuric acid is poured down the
side of the tube so as to lie below the lighter mixture (Fig.
89). At the surface of contact the sulphuric acid liberates
the nitric acid, and a brown layer is seen. Even when the
amount of the nitrate is very small, the brown tint may be
distinctly made out by contrast with the nearly colorless
liquids above and below it. FIG. 89.
Molecular Compounds. — When substances formed by union of
two compounds have a prevailing tendency to decompose into the
same two materials, and exhibit the chemical properties of their con-
stituents rather than individual ones of their own, they are often
called molecular compounds. Thus the above substance, 3FeSO4
2NO, gives off the nitric oxide again when warmed, and its solution
has the properties of a mixture of ferrous sulphate and nitric oxide.
Similarly, hydrates (pp. 120, 123) are formed by union of salts or other
substances with water, and are apparently, for the most part, decom-
444 INORGANIC CHEMISTRY
posed by solution. Double salts (p. 360), such as ferrous-ammonium
sulphate FeS04,(NH4)2SO4, 6H20, of which very many are known, are
of the same character. They are stable only in the solid form. There
are also compounds of salts with ammonia (see Compounds of copper
and silver), and with carbon monoxide (CO), one such compound being
formed with cuprous chloride (g.v.~).
The name 'molecular compounds is derived from the supposition
that, in these compounds, the molecules of the components retain their
integrity to some extent and are thus ready to be liberated. This is
an attempt to explain the fact that the behavior is that of the con-
stituents. It distinguishes molecular compounds from substances like
ammonium chloride and phosphorus pentachloride. The former may
be made by union of HC1 and NH3, but usually behaves rather as if
composed of NH4 and Cl. The latter (q.v.~) dissociates into PC13 and
C12, but with water gives phosphoric acid (cf. p. 181), which is derivable
from the pentachloride only. The distinction is of practical rather
than theoretical importance, however, for there are all gradations in
the behavior of molecular compounds. It is useful simply as a rough
means of classifying and remembering certain facts.
Distinguishing molecular compounds from ordinary compounds is
further justified by the fact that the constituents of molecular com-
pounds often seem to be saturated (p. 379), and no ordinary valences
are available for holding the new material. Thus in CaITCl2l the
ordinary valences are all saturated. Yet the salt forms the hydrate
CaCl2,6H20 with water (H^O11) which is likewise a saturated compound.
The conception of molecular compounds implies, therefore, the idea of
a sort of valence of molecules. Thus FeSO4 forms FeS04,7H20 and
FeS04,(NH4)2S04,6H20, and FeS04,K2S04,6H20, in all of which
seven other molecules are combined with it. The sulphates of other
bivalent metals (q-v.), such as copper and magnesium, form molecular
compounds of the same nature. Ammonium chloride, on the other
hand, is not a molecular compound, because, although NH3 unites with
HC1, HBr, HI, and HF, yet nitrogen is quinquivalent, and substances
like N205, NH4C1, etc., may fitly be regarded as ordinary compounds.
Preparation of Nitrogen Tetroxide. — This substance is liber-
ated by heating nitrates, other than those of potassium, sodium, or
ammonium :
2Cu(N03)8 -> 2CuO + 4N02 + O2.
OXIDES AND OXYGEN ACIDS OF NITROGEN 445
Iii most cases the oxide of the metal remains. When the mixed gases
are led through a U-tube immersed in a freezing mixture, the tetroxide
condenses as a pale-yellow liquid, and the oxygen passes on.
The compound may also be made by direct union of nitric oxide
and oxygen, or by oxidation of nitric oxide by concentrated nitric
acid (p. 443). It is likewise almost the sole product of the inter-
action of concentrated nitric acid and copper (see below). If any
nitric oxide were produced by the primary action, it would be oxidized
to nitrogen tetroxide in passing up through the acid.
Properties of Nitrogen Tetroxide. — The most striking pecul-
iarity of this gas is that, when hot, it is deep brown in color, and
when cold, pale yellow. When cooled, it gives a pale-yellow liquid boil-
ing at 22°, and an almost colorless solid melting at about — 12°. The
density of the vapor decreases very rapidly from 27° to 140°, and
increases again as the temperature falls. The molecular weights
calculated from these observations are: at 27°, 76.7; at 70°,
55.6 ; at 135°, 46.3 ; at 154°, 45.7. Now the molecular weights cor-
responding to the formulae N204 and N02 are 92 and 46 respectively,
so that these results mean that the deep-brown gas is N02, and that as
this is cooled it combines to form the colorless N204. Measurement of
the depression the substance causes in the freezing-point (cf. p. 291) of
glacial acetic acid gives the molecular weight 92, so that in solution
and at the temperature of freezing acetic acid (below 17°) the sub-
stance is all N204.
When the temperature is carried above 154°, by passing the
brown gas through a red-hot tube, the brown color disappears once
more, and nitric oxide and oxygen are formed. On cooling, the same
steps through brown gas to pale-yellow gas are retraced :
2NO + 02 <
Colorless Brown Colorless.
Since nitrogen tetroxide yields free oxygen more readily than does
nitric oxide, most ordinary combustibles burn in it. It has powerful
oxidizing properties ; and " fuming nitric acid," which contains it in
solution, is employed when oxidation is the special object in view.
This oxide is intermediate in composition between nitrons and
nitric anhydrides, and, when dissolved in cold water, gives both nitric
and nitrous acids :
N204 + H20 -> HN03 + HNOa.
446 INORGANIC CHEMISTRY
If a base is present, a mixture of the nitrate and nitrite of the metal
is produced (cf. p. 275).
When the water is not cooled, the nitrous acid (q-v.~), being un-
stable, gives nitric oxide and nitric acid, so that the result is :
3N02 + H2O <=> 2HNO8 + NO.
OXIDIZING ACTIONS OF NITRIC ACID.
When nitric acid gives up oxygen to any body, it is itself reduced.
Hence, according to convenience, we shall refer to oxidations by, or
reductions of nitric acid.
Nascent Hydrogen. — The extent to which the acid is reduced
by nascent hydrogen depends on the particular metal with which the
hydrogen is in contact when liberated (cf. p. 424). Thus, with zinc
and very dilute nitric acid, almost the only product, aside from zinc
nitrate, is ammonia :
4Zn + 8HN08 -> 4Zn(N08)2 (+ 8H) (1)
(8H) + HN08 -» NH8 + 3H2O (2)
NH8 + HN08 -* NH4N08 (3)
4Zn + 10HN08 -> 4Zn (N03)2 + NH4N08 + 3H2O
With the excess of nitric acid, ammonium nitrate is formed. Again,
when tin is the metal, the reduction is less complete, and hydroxyla-
inine is a product (p. 423).
When metals more active than zinc, such as magnesium, are used,
some of the hydrogen escapes oxidation and is liberated.
Heavy Metals. — Metals less active than tin, such as copper
and silver, do not displace hydrogen from dilute acids (p. 362), but
reduce nitric acid, nevertheless, and are converted into nitrates.
Platinum and gold (cf. p. 389) alone are not attacked. Thus, copper,
with somewhat diluted nitric acid, gives cupric nitrate and nitric oxide
(NO). In making the equation for this action we may resolve the
formula of nitric acid into those of water and the anhydride H20,N2O6.
This shows that the two molecules of the acid will give 2NO, and 3O
will remain :
2HNO3 -» H20 + 2NO ( + 30) (1)
(30) + 6HN03 + 3Cu -> 3H20 + 3Cu(NO8)2 (2)
8HNO3 + 3Cu -> 4H2O + 2NO + 3Cu(NO8)2
OXIDES AND OXYGEN ACIDS OF NITROGEN 447
The nitric oxide is liberated as a colorless gas, but forms the brown
tetroxide at once on meeting the oxygen of the air (p. 443).
When concentrated nitric acid is used with copper, almost pure
nitrogen tetroxide is obtained :
2HN08 -» H20 + 2N02 (+0) (1)
(0) + 2HN08 + Cu -> H.O + Cu (NO,), _ (2)
4HN03 + Cu -> 2H.X) + 2NO2 + Cu(N08)2
The equations for actions like the above may be built up from partial equations
of various kinds (c/. p. 229). Thus we may begin by forming the nitrate of the
metal, and then use the balance, consisting of hydrogen, along with other mole-
cules of nitric acid to secure the oxide and water :
Cu + 2HNO3 — > Cu(N03)2(+ 2H)
(2H) + 2HNO3-> 2H2O + 2NO2 _
Cu + 4HNO3 -- > Cu(N03)2+2H2
As the subdivision is purely arithmetical (p. 229), this procedure does not involve
the assumption that copper does actually displace hydrogen as a free element.
Yet it would not necessarily be incorrect to make even this supposition. Although
unable to liberate hydrogen in quantity from a dilute acid, copper may be held,
if we choose, to displace a minute amount of it :
Cu + 2HNO3 ±5 Cu(NO3)2 + H2, or Cu + 2H' £3 Cu " -f H2,
and to be restrained by the much more vigorous reverse action (p. 362) from con-
tinuing this operation. In this point of view the oxidation of the trace of free
hydrogen by the excess of nitric aeid continuously annihilates the possibility of
reverse action.
Complexities of Oxidation by Nitric Acid. — The above are
types of the interactions of metals with nitric acid. In actual experi-
ments the behavior is usually more complex. Thus, as a rule, the
action is very slow at first, and gathers speed with the accumulation of
the reduction products, which act catalytically.
Again, different concentrations of nitric acid give different prod-
ucts with the same metal. The most conspicuous effect of this kind
is the production of nitric oxide with diluted acid, and the invariable
formation of nitrogen tetroxide with concentrated acid. This is ex-
plained by the fact that nitrogen tetroxide cannot pass unchanged
through a liquid containing much water, for it gives nitric acid and
nitric oxide with the latter (p. 439). Conversely, where the nitric
acid is concentrated, nitric oxide, even if formed by the interaction
with the metal, must be oxidized to nitrogen tetroxide as it passes up
through the liquid (p. 443).
448 INORGANIC CHEMISTRY
Finally, intermediate concentrations give mixtures of these two
oxides, and, with zinc, even nitrous oxide (N20) and nitrogen may be
found in considerable quantities in the gases evolved.
Non- Metals. — With non-metals the actions are different in so
far that these elements do not give nitrates. Thus, sulphur boiled
in nitric acid gives sulphuric acid, along with nitric oxide, equation (3),
or with nitrogen tetr oxide, equation (6), or with both, according to the
concentration of the acid :
2HN03 -* 2NO + H20 (+ 30) (1)
(30) + H20 + S -> H2S04 (2)
2HNO3 + S -» 2NO + H2S04 (3)
2HN03 -> 2N02 + H20 + 0 x 3 (4)
(30) + H20 + S -» H2S04 (5)
6HN08 + S -> 6N02 + 2H20 + H2S04 (6)
The reader will note (cf. p. 272) that a separate equation, (3) and (6),
must be made for the formation of each reduction product. If NO
and N02 are both formed, they cannot arise from the same molecule of
nitric acid. They result from two actions which are independent,
although proceeding simultaneously in the same vessel (cf. p. 231).
Thus the equation :
2HN03 + C -> H20 + C02 + NO + N02,
is a misrepresentation. It implies that equimolar quantities of the
two oxides of nitrogen are formed. But this could only occur by
chance, and the balance would be destroyed the next moment by the
lowering in the concentration of the acid, giving the advantage to the
nitric oxide.
Compounds. — Compounds like hydrogen sulphide, hydrogen
iodide, and sulphurous acid, which are easily oxidized, interact with
nitric acid. With diluted nitric acid the products are free sulphur,
iodine, and sulphuric acid respectively.
The mixture of nitric acid and hydrochloric acid is known as
aqua regia. The chlorine set free by the oxidation of the hydro-
chloric acid is more active than is the ordinary solution of chlorine in
water, perhaps in consequence of catalytic action of the substances in
this solution, and combines with gold and platinum (q.v.~), converting
them into chlorides. Nitrosyl chloride (NOC1), which, however, does
OXIDES AND OXYGEN ACIDS OF NITROGEN 449
not interact directly with the noble metals, is formed at the same
time :
0
Cl-;H+H-0-i]![| = 6 + 2H: C1-»2H20 + CL, + C1-N=0.
•
The interaction with platinum is, therefore :
8HC1 + 2HM)3 + Pt -* H2Pt016 + 2NOC1 + 4 H20.
Chlorine, in this active state of the element, may be spoken of as
nascent chlorine (cf. p. 423).
NITROUS ACID, HYPONITROUS ACID, AND THEIR ANHYDRIDES.
Nitrites. — When the nitrates of potassium and sodium are heated,
they lose one unit of oxygen, and the nitrites remain :
2NaN08 -> 2NaN02 + 02.
Commonly lead is stirred with the melted nitrate and assists in the
removal of the oxygen. The litharge (PbO) which is formed remains
as a residue when the sodium nitrite is dissolved for recrystallization.
Nitrous Acid. — When an acid is added to a dilute solution of
a nitrite, a pale-blue solution containing nitrous acid is obtained. The
acid is very unstable, however, and, when the solution is warmed,
it decomposes :
3HNO2 -> HN03 + 2NO + H20.
When a concentrated solution of sodium nitrite is acidified, the
nitrous acid decomposes at once, and a brown gas containing the
anhydride escapes :
2H' + 2NO/ t? 2HN02 <=> H20 + N208 1 .
This behavior distinguishes a nitrite from a nitrate.
Keducing agents deprive nitrous acid of part or all of its oxygen :
2HI + 2HN02 ~» 2H20 + 2NO + I2.
Indigo is also converted by it into isatine (cf. p. 269). On the other
hand, oxidizing agents which are sufficiently active, like acidified
potassium permanganate, convert nitrous acid into nitric acid :
450 INORGANIC CHEMISTRY
3H2S04 + 2KMn04 -> K2S04 + 2MnS04 (+ 50) (1)
(50) + 5HN02 -» 5HN03 (2)
3H2S04 + 2KMn04 + 5HN02 -» K2S04 + 2MnSO4 + 5HNO3
Nitrous acid is much used in the making of organic dyes.
Nitrous Anhydride. — A study of the gas arising from the decom-
position of nitrous acid shows that in the gaseous state the anhydride
is almost entirely dissociated :
N203 <z> NO + N02.
When the mixture is led through a U-tube immersed in a freezing
mixture at — 21°, a deep-blue liquid is obtained which appears to be
the anhydride itself. This begins to dissociate before reaching its
boiling-point, and at + 2° gives off nitric oxide.
The same equimolar mixture of the two gases is obtained by the
action of water on nitrosylsulphuric acid (p. 383).
Hyponitrous Acid. — This acid is formed by the interaction of
hydroxylamine and nitrous acid in aqueous solution :
H-0-N=N-0-H.
With nitrate of silver, the yellow, insoluble silver hyponitrite Ag2N202
is precipitated. When this salt is shaken with an ethereal solution of
hydrogen chloride, the acid is liberated, and the insoluble silver
chloride may be separated by nitration. Finally, evaporation of the
ethereal solution leaves hyponitrous acid as a white mass. It explodes
when heated, and its solution in water is an exceedingly feeble acid.
The warm aqueous solution decomposes slowly, giving nitrous oxide :
H2N202 -> H20 + N20,
and this change is not capable of reversal.
Nitrous Oxide. — Nitrous oxide is prepared by heating ammonium
nitrate, or a mixture of a salt 'of ammonium and a nitrate :
NH4N03 -» 2H20 + N20.
The steam condenses, and the nitrous oxide may be collected over
warm water, or dried and compressed into steel cylinders.
Its solubility in cold water is considerable : 130 volumes in 100 at
OXIDES AND OXYGEN ACIDS OF NITROGEN 451
0°. At 25° this falls to 60 volumes in 100. In dissolving, the gas
forms no compound with water. The substance melts at — 102.3°,
and boils at — 89.8°. The vapor tension of the liquid at 0° is 30.75
atmospheres ; at 12°, 41.2 atmospheres ; and at 20°, 49.4 atmospheres.
The critical temperature is 38.8°.
A glowing splinter of wood bursts into flame when introduced into
nitrous oxide, and phosphorus, sulphur, and other combustibles, burn it
with much the same vigor as in oxygen. In all cases oxides are formed,
and nitrogen is set free. The rapidity with which bodies combine with
oxygen obtained from nitrous oxide is doubtless due to the fact that it
is an endothermal compound, and the heat liberated by its decomposi-
tion assists the ensuing combustion :
2N20 -» 2N2 + 02 + 2 x 18,000 cal.
It is to be noted that the -effect of the heat of decomposition will be
partly offset by the dilution of the oxygen with nitrogen. Yet the
proportion of nitrogen to oxygen is only half as great as in air, so
that on the whole the conditions are much more favorable to combus-
tion in this gas.
Nitrous oxide, when cold, does not behave like free oxygen. Nitric
oxide, when mixed with it, gives none of the red nitrogen tetroxide.
Metals do not rust in it, and the haemoglobin of the blood is unable
to use it as a source of oxygen. It was Davy who first observed that
nitrous oxide could be taken into the lungs, and that, since it furnished
no oxygen, insensibility followed its use. By suitable admixture of an
amount of air sufficient to sustain life, it is employed as an anaesthetic
for minor operations. The hysterical symptoms which were observed
to accompany its administration, and caused it to receive the name of
" laughing gas," are such as attend the use of all anaesthetics.
Graphic Formulae of Nitric Acid and its Derivatives : Explo-
sives. — The following equation shows the graphic formulas of nitric
acid and of ammonium nitrate :
H,0.
The structural formula of the latter is intended to explain the fact
that the salt is able to exist at all, by representing the oxygen and
hydrogen as being separated from one another and attached to differ-
452 INORGANIC CHEMISTRY
ent nitrogen units. When the equilibrium of the system is disturbed
by heating, the oxygen and hydrogen unite to form water, an arrange-
ment which is much more stable, and nitrous oxide (see above) escapes
with the steam.
The decomposition of nitroglycerine and gun-cotton (p. 441), as
well as of ammonium nitrite (p. 416), is explained in the same way.
These substances are made by actions which, like the above neutraliza-
tion, take place in the cold, and the groups, containing the oxygen on
the one hand and carbon and hydrogen on the other, become quietly
united without more serious interaction. Thus the formation of nitro-
glycerine (p. 441) appears as follows :
H H
I I
H - C - OH HO - N02 H - C - 0 - N02
& &
I I
H - C - OH + HO - NO2 -* H - C - 0 - N02 + 3HaO
I I
H - C - OH HO - N02 H - C - O - NO2
I I
H H
When the nitroglycerine is heated, or receives a mechanical shock,
the oxygen all unites with the carbon and hydrogen, and the nitrogen
escapes :•
4C8H5(ON02)3-> 12C02 + IOH.,0 + 6N2 + 02.
That nitroglycerine is a more sensitive explosive than gunpowder
is due to the fact that, in the former, the materials required for the
chemical change are already within the same molecule, whereas in
the latter (q.v.) they are contained in the separate molecules of a mix-
ture. Even after the most careful incorporation, the oxygen of the
potassium nitrate can hardly be uniformly so near to the carbon
mechanically mixed with the salt, as are these elements in nitro-
glycerine or gun-cotton. In the latter the oxidation of the hydrogen
and carbon is intramolecular.
Substances like hydrazoic acid (p. 422) and nitrogen iodide (p. 425)
might seem to constitute a third kind -of explosive. Here the change
consists in the resolution of the compound into its constituents. Still, if
we consider the case of hydrazoic acid, for example : 2X3H — » SN2 + H2,
we see that the action consists, after all, in the union of the con-
stituents to form the more stable combinations K2 and H^. It is,
therefore, similar in principle to the explosion of nitroglycerine.
OXIDES AND OXYGEN ACIDS OF NITROGEN 453
The Principle of Transformation by Steps. — It may have
occurred to the reader as strange that it should be possible to make
nitric anhydride by distilling a warm mixture (p. 440) when the product
decomposes spontaneously, even when kept in the cold. How can a
compound be fitted together under certain conditions, when under the
same or, even, under more favorable conditions it proves to be incapable
of continued existence ? We should expect rather that obtaining the
products of its decomposition would have been the only result of the
effort to make it.
Extraordinary as this fact appears to be, it is nevertheless very
commonly encountered. Perchloric acid is made by a distillation
(p. 276) and afterwards breaks up of its own accord. So, also, hypo-
chlorites are formed first and can be isolated. But under the same con-
ditions the further transformation to chlorates will occur (p. 272). A
simple case is that of sulphur made by precipitation (p. 376) at the ordi-
nary temperature. Although it is naturally solid below 119°, yet,
when first thrown down, it is in the form of liquid droplets which, if
undisturbed, may remain fluid for weeks. Similarly, sulphur vapor
condenses on glass in drops which remain liquid until they are touched
or rubbed. Finally, a supersaturated solution (p. 159) is not unlike
cold liquid sulphur.
In all these cases there is a possibility of further change, which,
when it comes, will liberate heat or some other form of energy. Thus,
heat is set free when the liquid sulphur is precipitated. The amount
of heat would have been greater, by the heat of fusion of the sulphur,
if solidification had occurred simultaneously. But, in spite of the
existence of this justification for the final step, this step is not taken.
So, the decomposition of the vapor of the perchloric acid or of the
nitric anhydride would have added to the amount of energy liberated
as heat, but this additional step was postponed. In other words,
transformations -which proceed spontaneously and with evolution of
heat may go f orward by steps, when there are intermediate substances
capable of existence. This is known as the principle of transformation
by steps, and was first described by Ostwald.
Exercises, — 1. Make the equation for the interaction of ferrous
chloride, hydrochloric acid, and nitric acid (p. 442), and for all the ac-
tions concerned when the test for a nitrate (p. 443) is applied to sodium
nitrate. What volume (at 0° and 760 mm.) of NO is obtained from
one formula-weight of nitric acid (p. 216, Ex. 8)?
454 INORGANIC CHEMISTRY
2. Should you classify as molecular compounds (p. 443) : Chlorine
hydrate, ammonium hydroxide, KI3 (p. 235), sulphurous acid, sodium
pentasulphide (p. 376) ? Justify your answer.
3. At 27°, what proportions of the molecules of nitrogen tetroxide
are in the forms of N02 and N204 respectively (p. 445) ? At the same
temperature what fraction of the material, by weight, is in the former
condition ? What are the relative, volumes of the tetroxide, and of the
nitric oxide and oxygen obtained by its decomposition (p. 445) ?
4. Make an equation showing the production of nitrous oxide by the
action of zinc on nitric acid.
5. Make the correct equations showing the formation of nitric oxide
and nitrogen tetroxide by the interaction of carbon and nitric acid
(p. 448).
6. Justify the graphic formula assigned to nitric acid (p. 441).
7. Using the anhydride method (p. 274), make the equations for
the interactions of N2O4 and water (p. 445) and of nitric acid and
sulphur (p. 448).
8. Make a classified list, with examples, of all the kinds of
interactions which, in this and preceding chapters, have been named
oxidations and reductions (e.g. pp. 70, 72, 110, 170, 172, 231, 237, 269,
306, 374, 420, 424, 446. See also Chemistry of copper and tin).
CHAPTER XXVII
PHOSPHORUS
The Chemical Relations of the Element. — There are many
things in the chemistry of phosphorus and its compounds which
remind us of nitrogen. Yet these are largely referable to the fact that
the elements are both non-metals and both have the same valences, viz.
three and five. The behavior of the compounds is often very different.
For the present it is sufficient to say that both give compounds with
hydrogen, NH3 and PH3, and both yield oxides of the forms X208, X204,
and X2O6. The first and last of these oxides are acid-forming, and
phosphorus, therefore, gives acids corresponding to nitrous and nitric
acids, although there is more variety in the proportion of water com-
bined with the anhydride (cf. p. 278). The element is thus a non-metal
(see Comparison with nitrogen and with sulphur, end of this chapter).
Occurrence. — This element is found widely disseminated in na-
ture, usually in the form of phosphates. Calcium phosphate, for ex-
ample, which is derived (p. 278) from phosphoric acid (H8P04) by dis-
placement of its hydrogen by calcium, and has, therefore, the formula
Ca8(P04)2, is found in most soils. It constitutes a large part of the
solid material of the bones and teeth of animals and of the beds of
fossil bones found in Florida and Tunis. A conspicuous mineral
related to this substance is apatite, Ca5F(P04)3. It is found in large
quantities in Canada, and is a component of many rocks. Complex
organic compounds containing phosphorus are essential constituents of
protoplasm and of the materials of the nerves and the brain.
Preparation. — Brand, merchant and alchemist, of Hamburg,
discovered phosphorus (1669) by distilling the residue from evapo-
rated urine, in the course of his search for the Philosopher's stone.
The mode of preparing it from bone-ash was first published by Scheele
(1771). Green bones contain about 58 per cent of calcium phosphate.
After the gelatine has been extracted from them, by means of water
boiling under pressure, they are subjected to destructive distillation, a
455
456 INORGANIC CHEMISTRY
process which yields bone-oil. The residue is a mixture of carbon
(q.vC) and calcium phosphate. It is used by sugar refiners as a decol-
orizer. When its powers in this direction have been exhausted, it is
calcined — that is to say, all the combustible matter is burned out of it,
— and the product is bone-ash. Formerly this was used in making
phosphorus, but now the less expensive calcium phosphate of fossil
origin is employed.
A mixture of powdered bone-ash or calcium phosphate and sulphu-
ric acid (sp. gr. 1.5 to 1.6) is heated with steam and stirred in a
wooden vat :
Ca8(P04)2 + 3H2SO4 -> 2H3PO4 + 3CaS04.
The calcium sulphate is partly precipitated during the heating. The
liquid obtained by filtration is evaporated in leaden pans. During this
process most of the remainder of the calcium sulphate is deposited and
a syrupy, crude phosphoric acid is obtained. This acid is mixed with
sawdust, or carbon in some form, and the mixture is first heated to a
moderate temperature and then distilled in earthenware retorts. Two
actions take place in succession. The phosphoric acid loses water and
turns into metaphosphoric acid, then the latter is reduced by the car-
bon, carbon monoxide and phosphorus vapor passing off :
H3P04->H20
2HP03 + 60 -> H2 + 6CO + 2P.
A white heat is required for the distillation, and a pipe from the tubu-
lar clay retort conducts the vapors into cold water, in which the phos-
phorus collects.
A much simpler process depends on the use of the electric furnace
(Fig. 90). The calcium phosphate is mixed with the proper propor-
tions of carbon and silicon dioxide (sand), and the mixture is intro-
duced continuously into the furnace. The discharge of an alternating
current between carbon poles produces the very high temperature
which the action requires. The calcium silicate which is formed fuses
to a slag, and can be withdrawn at intervals. The gaseous products
pass off through a pipe and the phosphorus is caught under water :
Ca3(P04)2 + 3Si02 + 5C -> 3CaSi03 + SCO -f- 2P.
We may regard the phosphate as being composed of two oxides, 3CaO,
P20ft. It thus appears that the calcium oxide has united with the
457
silica, which is an acid anhydride (cf. p. 381) : CaO + Si02 — •» CaSi03,
while the phosphoric anhydride has been reduced.
The phosphorus, after purification, is cast into sticks in tubes of
tin or glass, standing in cold water.
The Electric Furnace. — By an electric furnace is understood an
electro-thermal arrangement in which the heat produced by some
resistance offered to the current,
such as that of an air-gap between
the carbons, is used to produce chem-
ical change. Electrolysis plays no
part in the phenomena, and an
alternating current, which can pro-
duce no electrolytic decomposition,
is generally employed. The re-
stricted area within which the heat
is developed makes possible the
attainment of a high temperature
(see Calcium carbide).
Physical Properties. — There
are two perfectly distinct kinds of
phosphorus, known as ordinary, or
yellow phosphorus, and red phos-
phorus. Yellow phosphorus, pre-
pared as described above, is at first
transparent and colorless, but after exposure to light acquires a
superficial coating of the red variety. It melts at 44 D and boils at
269° (according to some authorities, at 287°). Its molecular weight is,
at 313°, 128, and at a red heat 119.8. As the atomic weight is 31, the
formula, within this range, is P4. At 1700° the value 91.2 is held to
indicate partial dissociation into P2. In solution the formula is P4.
Yellow phosphorus is very soluble in carbon disulphide, less soluble in
ether and other organic solvents, and insoluble in water. It is ex-
ceedingly poisonous, less than 0.15 g. being a fatal dose. Continued
exposure to its vapor causes necrosis, a disease from which match-
makers are liable to suffer. The jawbones and teeth are particularly
liable to attack.
Red phosphorus is a dull red powder consisting of small tabular
crystals. It is obtained by heating yellow phosphorus to about 250°
in a vessel fvoin which air is excluded, The change is ranch iwn-«
FIG. 90.
458 INORGANIC CHEMISTRY
rapid at slightly higher temperatures. Since a great amount of heat is
evolved in the transformation, the action is apt to become violent and
to cause volatilization of a large part of the phosphorus. In presence
of a trace of iodine the transformation is greatly accelerated, and takes
place even in the cold.
Red phosphorus does not melt, but passes directly into vapor. Its
vapor is identical with that of yellow phosphorus. It is insoluble in
carbon bisulphide and other solvents. It is not poisonous, and,
unlike yellow phosphorus, does not require to be kept under water to
avoid spontaneous combustion.
Chemical Properties. — Yellow phosphorus unites directly with
the halogens with great vigor. It unites slowly with oxygen in the
cold, and with sulphur and many metals when the materials are heated
together.
The slow union of phosphorus with atmospheric oxygen is accom-
panied by the evolution of light, although the temperature is not such
as we usually associate with incandescence. The name of the element
(Gk. <£ws, light ; <}>fp(a, to bear) records this property. Apparently the
chemical energy, transformed in connection with the oxidation, is
converted, in part at least, into radiant energy instead of completely
into heat.* A curious fact in connection with the luminosity and
concomitant oxidation of phosphorus is that these occurrences depend
upon the concentration of the oxygen gas as well as upon the tempera-
ture. Thus, phosphorus does not shine or oxidize in pure oxygen
below 27°. If the concentration of the oxygen is reduced to 200 mm.
or less by means of a pump, or by mixing with an indifferent gas such
as nitrogen, phosphorescence becomes perceptible at the ordinary
temperature. This explains the luminosity shown in the air. At
lower temperatures, lower pressures have to be used. The phosphores-
cence may be destroyed by the vapor of turpentine and other sub-
stances, which perhaps act as catalytic agents and restrain the
oxidation.
* The same production of light from chemical action in a cold body is seen in
the luminosity of certain parts of some animals, such as fireflies and some species
of fish. In many violent chemical changes the light given out is conspicuously
greater than that proper to the temperature produced (cf. p. 73), and must come,
therefore, in part, directly from the chemical energy. Thus, burning magnesium
has a temperature of about 1350°, while the production of light of the same char-
acter, by mere incandescence, would require a temperature of about 5000°.
PHOSPHORUS 4f>9
The difference in behavior of pure and diluted oxygen may be shown by pour-
ing a solution of phosphorus in carbon disulphide on to two strips of filter paper.
One of the strips, hung in the air, catches fire as soon as the evaporation of the
solvent has exposed a large area of finely divided phosphorus. The othep, hung
in a jar of oxygen, remains unaffected, but becomes ignited instantly upon removal
from the jar.
The slow oxidation of phosphorus is accompanied by the produc-
tion of ozone, but the nature of the action is still unknown.
Chemical Properties of Red Phosphorus. — This variety of the
element, since it is formed with evolution of heat, contains less energy
than yellow phosphorus and is much less active. It does not catch
fire in the air below 240°, while ordinary phosphorus ignites at 35-45°.
Indeed, it is the vapor that begins to combine with oxygen, and
this behavior is only an independent proof of the low vapor tension of
the red variety. When the vapor tension of yellow phosphorus has
reached 760 mm. (at 269°, the b.-p.), that of red phosphorus is
almost imperceptible. Even when the former has become two atmos-
pheres, the latter is still very small.
Ked phosphorus is to 'be regarded as the normal, stable form of
phosphorus. The fact that yellow phosphorus can be kept a great
length of time, and is changed but slightly on exposure to light, only
shows that the transformation into a stabler condition is retarded by
the lowness of the temperature (of. p. 72). The relation between the
two varieties of phosphorus is quite distinct from that between rhom-
bic and monoclinic sulphur (p. 368). In the latter case there is a
definite temperature of transformation (96°) above which one form
completely disappears, and below which the other form is incapable of
permanent existence. With the varieties of phosphorus no such point
of transformation exists. The red phosphorus is the more stable at
all temperatures at which both forms are known. The yellow turns
into red, but the red never into the yellow. It is only by condensing
the vapor that the yellow kind is obtained. It seems probable that
red phosphorus is a polymer (p. 242) of yellow phosphorus (P4).
The production from the vapor (also P4) of the yellow solid, instead
of the red solid whose formation would be accompanied by a larger
liberation of heat, is simply an illustration of the principle of transfor-
mation by steps (p. 453).
When two or more forms of an element, or even of a compound (see Isomers),
occur, they are commonly spoken of as allotropic modifications. The term is
applied to oxygen and ozone (q.v.} which are certainly, and to red and yellow
460 INORGANIC CHEMISTRY
phosphorus which are, probably, chemically distinct substances. It is used also of
rhombic and monoclinic sulphur, where the difference is purely physical. In short,
it has at present no scientific value, for it covers a heterogeneous mass of phe-
nomena which, in part, still await elucidation. If allotropic modifications
were to be denned as substances (p. 32) composed of the same materials,
but possessing different proportions of free or available energy, and,
therefore, different physical properties and different degrees of chemical activity,
which is apparently the sense in which the expression is commonly employed, then
ice, water, and steam would be examples of such substances. In each of the
above four illustrations, the second (in the case of water, the third) is the more
active form.
Uses of Phosphorus. — The greater part of the phosphorus of
commerce is employed in the manufacture of matches. The first arti-
cles of this sort (1812) were sticks coated with sulphur and tipped
with a mixture of potassium chlorate and sugar. For ignition they
were dipped into a bottle containing asbestos moistened with concen-
trated sulphuric acid. Matches involving the use of phosphorus (1827)
have now displaced all others. In making common matches which
strike on any rough surface, the sticks are first dipped in melted
sulphur or paraffin to the extent of about half an inch. The head is
often composed of manganese dioxide or red lead, and a little potas-
sium chlorate, which supply oxygen, a small proportion of free phos-
phorus and antimony trisulphide, which are both combustible, and
dextrine or glue. A paste made of these materials is spread evenly
upon a slab, and the prepared sticks fixed in a frame are dipped once
or twice in the mixture.
In the case of " safety " matches, the mixture upon the head is not
easily ignited by itself. It is composed of potassium chlorate or di-
chromate, some sulphur or antimony trisulphide, and a little powdered
glass to increase the friction, all held together with glue. Upon the
rubbing surface on the box is a thin layer of antimony trisulphide
mixed with red phosphorus and glue. The friction converts a little
of the red phosphorus into vapor. To prevent smoldering of the
burned matches, the upper ends of the sticks are sometimes soaked
in a solution of alum or sodium phosphate.
A small amount of yellow phosphorus is employed in making rat
poison, which is a mixture of phosphorus, lard (as solvent), and flour,
made into dough.
Phosphine. — Three hydrides of phosphorus are known. These
are, phosphine PH3 (a gas), a liquid hydride P2H4, which is presumably
PHOSPHORUS 401
the analogue of hydrazine (N2H4), and a solid hydride P4H2. Phosphine
does not seem to be produced under ordinary circumstances by the
direct union of the elements. It is formed slowly, however, by the
action of nascent hydrogen, from zinc and hydrochloric acid at 70°, upon
yellow phosphorus. The gas may be made by boiling yellow phos-
phorus with potassium hydroxide solution in an apparatus similar to
that used for generating hydrogen. Potassium hypophosphite is
formed at the same time :
3KOH -f- 4P + 3H2O -> 3KH2P02 + PH3.
The gas made in this fashion contains a small proportion of the vapor
of the liquid hydride, which is spontaneously inflammable, and conse-
quently the mixture catches fire on emerging into the air from the
delivery tube. To avoid explosions, the air in the flask must be dis-
placed by hydrogen or illuminating-gas before heat is applied. This
product contains also free hydrogen, in increasing quantities as the
action goes on, in consequence of the reduction of the water and potas-
sium hydroxide by the potassium hypophosphite : KH2P02 + KOH +
H20 —> K2HPO4 + 2H2. Potassium phosphate is formed.
The simplest method of preparing the gas is by the action of water
upon calcium phosphide :
Ca3P2'+ 6H20 _» 3Ca(OH)2 + 2PH
This action is analogous to that of water upon magnesium nitride
(p. 417) by which ammonia is produced. In consequence of the fact
that calcium phosphide is a substance of irregular composition, a mix-
ture of all three hydrides is generally obtained. By passing the gas
through a strongly cooled delivery tube, however, the liquid compound
is condensed and fairly pure phosphine passes on.
Phosphine is a colorless gas, which is easily decomposed by heat
into its elements. When burned, it forms phosphoric acid. It is exceed-
ingly poisonous and, unlike ammonia, it is insoluble in water, and pro-
duces no basic compound corresponding to ammonium hydroxide when
brought in contact with this substance. It resembles ammonia, for-
mally at least, in uniting with the hydrogen halides (see below). It
differs from ammonia, however, inasmuch as it does not unite with the
oxygen acids. Phosphine acts upon solutions of some salts, precipitat-
ing phosphides of the metals :
3CuS04 + 2PH8 -* Cu3P2 + 3H2S04.
462 INORGANIC CHEMISTRY
The liquid hydrogen phosphide boils at 57°. The molecular weight,
as determined by the density of its vapor, shows the formula to be P2H4.
It forms no salts, and is therefore quite unlike hydrazine. When
exposed to light, it decomposes, giving phosphine and the solid hydride.
Phosphonium, Compounds. — Hydrogen iodide unites with phos-
phine to form a colorless solid crystallizing in beautiful, highly refract-
ing, square prisms : PH3 -f- HI — > PH4I. Hydrogen chloride combines
similarly with phosphine, but only when the gases are cooled by a freez-
ing mixture, or are brought together under a total pressure of 18 atmos-
pheres at 14°. When the pressure is released, rapid dissociation occurs.
This dissociation is one of the many cases where an action which
absorbs heat, nevertheless goes on spontaneously (cf. p. 27). The
indispensable fall in the energy of the system takes place by virtue of
the diffusion of the constituents, and in amount this more than offsets
the heat acquired.
' In imitation of the ammonia nomenclature, these substances are
called phosphonium iodide and phosphonium chloride. They are
entirely different, however, from the corresponding ammonium deriva-
tives, for the PH4* ion is unstable. When brought in contact with
water they decompose into their constituents, the hydrogen halide
going into solution, and the phosphine being liberated as a gas.
Halides of Phosphorus. — The existence of the following halides
has been proved conclusively :
P2I4 (solid)
PF8 (gas) PC13 (liquid) PBr3 (liquid) P3I (solid)
PF5 (gas) PC15 (solid) PBr5 (solid)
These substances may all be formed by direct union of the elements.
They are incomparably more stable than are the similar compounds of
nitrogen. They are all decomposed by contact Math water, and give an
oxygen acid of phosphorus and the hydrogen halide (see below). This
action was used in the preparation of hydrogen bromide (p. 231) and
hydrogen iodide (p. 238).
Phosphorus trichloride is made by passing chlorine gas over
melted phosphorus in a flask until the proper gain in weight has
occurred. It is a liquid, boiling at 76°. When excess of chlorine is
employed, phosphorus pentachloride, which is a white solid body, is
PHOSPHORUS 463
formed. When moist air is blown over any of these substances, the
water is condensed to a fog by the hydrogen halide. In the case of the
interaction of phosphorus pentachloride and water, phosphoric acid is
formed:
PC16 + 4H20 -» H3PO4 + 5HC1.
With a limited supply of water the hydrolysis is not so complete,
and phosphoryl chloride * (phosphorus oxychloride), a liquid boiling at
107°, is produced :
PC15 + HOH -> POC1., + 2HC1.
This interaction of phosphorus pentachloride and water' is a per-
fectly general one, and takes place with most compounds containing
hydroxyl. Thus, when alcohol (which differs from water in having
ethyl (C2H5), instead of hydrogen, combined with hydroxyl) is poured
upon it, ethyl chloride, phosphoryl chloride, and hydrogen chloride are
formed :
C2H5OH + PC15 -> C2H5C1 + POC13 + HC1.
The same action takes place with all carbon compounds containing hy-
droxyl, and is used as a means of showing the presence of this group
in their structure. The reaction is shown by inorganic compounds
also. Thus, anhydrous sulphuric acid gives sulphuryl chloride, which
may be separated from the phosphoryl chloride by fractional distilla-
tion (see Petroleum) :
S02(OH)2 + 2PC15 -> SOijCLj + 2POC13 + 2HC1.
Phosphorus pentachloride, when heated, reaches a vapor tension
of 760 mm. at 140°, and while still solid. It therefore passes freely
into vapor (boils, so to speak) at this temperature, and condenses di-
rectly to the solid form. This sort of distillation is called sublimation.
At a pressure above that of the atmosphere it melts at 148°. This is
simply a case in which the vapor tension of the solid, increasing with
rise in temperature, happens to pass the arbitrary value of the opposing
pressure (one atmosphere) peculiar to experiments carried on in open
vessels, before the melting-point is reached. The same phenomenon is
shown by sulphur trioxide (p. 381).
* This substance is a mixed anhydride (p. 397) of phosphoric acid and hydro-
gen chloride.
464 INORGANIC CHEMISTRY
Phosphorus pentachloride (cf. p. 255) and pentabromide, when
vaporized, are partially dissociated :
Since the first two members of this equilibrium are colorless, while the
bromine is brown, this action may be used to illustrate the effect upon
a system, of increasing the concentration of one of the interacting
substances (p. 249). Two tubes of equal volume and containing
equal amounts of the pentabromide are prepared. A small amount of
the tribromide is added to the second, and both are sealed up. When
the tubes are now heated to the same temperature, the contents of the
second will be less strongly colored by bromine in consequence of the
greater activity in it of the reversing action.
Oxides of Phosphorus. — The oxides of phosphorus are the so-
called trioxide P406, the pentoxide P205, and a tetroxide P204.
The pentoxide is a white powder formed when phosphorus is burned
with a free supply of oxygen. It unites with water with great violence
to form metaphosphoric acid (see below), and hence is known as
phosphoric anhydride : P205 + H2O — » 2HP08. In the laboratory
this action is frequently utilized for drying gases (p. 101) and for re-
moving water from combination (p. 276). The vapor density of
the pentoxide indicates that its formula is P4010, but its whole chemical
behavior is equally well represented by the simpler formula.
The trioxide is obtained by burning phosphorus in a tube with a
restricted supply of air. It is a white solid, melting at 22.5° and boil-
ing at 173°. On account of the ease with which it may be volatilized,
it can be separated by distillation from any pentoxide formed at the
same time. The operation must be carried out in an apparatus from
which the air is excluded, as the trioxide unites spontaneously with
oxygen. The vapor density of the substance shows that its formula is
P406. This formula is preferred to the simpler one because, although
the oxide is the anhydride of phosphorous acid, it nevertheless unites
exceedingly slowly with cold water to form this substance. It inter-
acts vigorously with hot water, but phosphine, red phosphorus, hypo-
phosphoric acid, and phosphoric acid are amongst the products, and
very little phosphorous acid escapes decomposition. When this oxide
is heated to 440° it decomposes, giving the compound P2O4 and red
phosphorus.
PHOSPHORUS 465
Acids of Phosphorus. — There are four different acids of
phosphorus in which, probably, four distinct stages of oxidation are
shown. The highest stage is represented by three phosphoric acids,
where the degree of hydration of the anhydride varies. These are
orthophosphoric acid H8P04, pyrophosphoric acid H4P207, and metaphos-
phoric acid HP08. The other acids are less important. They are
phosphorous acid (H3P08), hypophosphoric acid (H4P206), and hypo-
phosphorous acid (H3P02).
The Phosphoric Acids. — The relation between the three different
phosphoric acids may be seen by considering them as being formed
from phosphorus pentoxide and water. It will be remembered that in
the majority of cases already considered, this sort of action takes place
for the most part in but one way. Thus, nitric acid is known in but one
form, which is produced by the union of one molecule each of nitrogen
pentoxide and water : N205 + H20 — > 2HN08. Similarly the chief sul-
phuric acid is the one formed from one molecule of sulphur trioxide
and one molecule of water : S03 + H20 — > H2S04, although here we
have both the hydrate H2S04, H20, which might be written H4S06 and
disulphuric acid H2S207. Referred to the anhydride these three acids
are H2O,S08, 2H20,SOS, and H20,2S08. Periodic acid (p. 278) has
a set of even more complexly related acids or salts.
Now, when phosphoric anhydride acts upon water we obtain a solu-
tion which, on immediate evaporation, leaves a glassy solid, HPO8,
known as metaphosphoric acid. This is H20, P20B. When, however,
the solution is allowed to stand for some days, or is boiled with a little
dilute nitric acid whose hydrion acts catalytically, the residue from
evaporation is H8P04, orthophosphoric acid :
P205 + 3HP -> 2H8P04 or HP08 + H20 -» H8P04.
This acid is 3H20, P206, and no further addition of water can be
effected.
Conversely, when orthophosphoric acid is kept at about 255° for a
time, it slowly loses water, and H4P207, pyrophosphoric acid, is ob-
tained :
This acid is 2H20, P205. Further desiccation leaves metaphosphoric
acid, which cannot be further resolved into phosphorus pentoxide and
water. When dissolved in water, pyrophosphoric acid slowly resiimes
466 INORGANIC CHEMISTRY
the water which it has lost and gives the ortho-acid again. The pyro-
acid does not seem to be formed by hydration of the meta-acid, but
only by dehydration of the ortho-acid.
The relations of all these substances are more clearly seen in the
graphic formulae :
=0
=0
0
= 0
= 0
The addition or removal of water leaves the valence of the phosphorus
unchanged.
Pyrosulphuric acid and its salts when dissolved in water give sul-
phuric acid and acid sulphates respectively. That is to say, the ion S207
is not capable of existence. But the very slow rate at which the less
hydrated phosphoric acids change into the more hydrated ones shows
that ions like PO/", P08', and P207"" may be comparatively stable.
The behavior of solutions of the salts shows this even more clearly.
Orthophosphoric Acid and Its Salts. — As we have seen, ordi-
nary calcium phosphate is the source of the impure, commercial acid.
Pure orthophosphoric acid may be made by boiling red phosphorus
with slightly diluted nitric acid and evaporating the water and
excess of nitric acid. The product is a white, crystalline, deliques-
cent solid.
This acid is much weaker than sulphuric acid, and is dissociated
chiefly into the ions H* and H2PO/. The dihydrophosphanion is
broken up to some extent into H" and HPO4", as we learn from the
fact that the solution of the sodium salt NaH2P04 is acid. The ion
HP04" is hardly dissociated -at all, for a solution of the salt NagHPC^
is alkaline in reaction.
As a tribasic acid, it forms salts of three kinds, such as NaH2P04,
Na.jHPO4, and Na8P04. These are known respectively as primary,
secondary, and tertiary sodium orthophosphate. The primary sodium
phosphate is faintly acid in reaction. The secondary one is slightly
alkaline, because of hydrolysis arising from the tendency of the
PHOSPHORUS 467
hydrion of the water to combine with the HP04" to form H2PO/ (cf.
p. 344). The tertiary phosphate is stable only in solid form, and can
be made by evaporating to dryness a mixture of the secondary
phosphate and sodium Irydroxide :
Na2HP04 + NaOH <=> NagPO, + H2O f.
When the product is dissolved in water, the action is reversed (cf.
p. 375). Mixed phosphates are also known, particularly sodium-am-
monium phosphate (microcosmic salt), NaNH4HP04, and the insoluble
magnesium-ammonium phosphate, MgNH4P04.
Primary calcium phosphate, known in commerce as " superphos-
phate," is used as a fertilizer. On account of its insolubility, and
since plants can take up soluble substances only, calcium phosphate is
of relatively little service to plants. It is therefore converted into the
" superphosphate," which is soluble, by treatment with dilute sulphuric
acid:
+ H2S04 <=» 2CaS04 '+ CaH4(P04)2.
The tertiary phosphates are unchanged by heating. The primary
and secondary phosphates, however, retaining, as they do, some of the
original hydrogen of the phosphoric acid, are capable of losing water
like phosphoric acid itself, when heated. The actions are slowly re-
versed when the products are dissolved in water :
KaH2P04 <=± NaPO3 + H20,
± Na4P2O7 + H20.
It will be seen that the meta- and pyrophosphates of sodium are
formed by these actions ; and this is indeed the simplest way of form-
ing these substances, since the acids themselves are not permanent in
solution, and are too feeble to lend themselves to exact neutralization.
Ammonium salts of phosphoric acid lose ammonia, as well as water,
when heated (cf. p. 421). Thus, microcosmic salt gives primary
sodium phosphate :
NaNH4HP04-*NaH2P04 + NH8,
and this in turn is converted into the metaphosphate by loss of water.
Pyrophosphoric Acid. — This acid, obtained by heating ortho-
phosphoric acid, may be prepared in pure form by making the spar-
ingly soluble lead salt from sodium pyrophosphate, and precipitating
468 INORGANIC CHEMISTRY
the plumbion by addition of sulphuric acid. In solution it gradually
reunites with water. Although tetrabasic, having four hydrogen
atoms which may be displaced by metals, only two kinds of salts
of this acid are known. These are the normal salts, such as Na4P207,
and those in which one-half of the hydrogen has been displaced by
a metal such as
Metaphosphoric Acid. — This is the " glacial phosphoric acid "
of commerce, and is usually sold in the form of transparent sticks.
It is obtained by heating orthophosphoric acid, or by direct union of
phosphorus pentoxide with a small amount of cold water. It passes
into vapor at a high temperature, and its vapor density corresponds to
the formula (HP08)2. The existence of certain complex salts confirms
our belief in the existence of a tendency to association (p. 242).
Sodium metaphosphate NaP08, in the form of a small globule
obtained by heating microcosmic salt on a platinum wire, is used in
analysis. When minute traces of oxides of certain metals are
placed upon such a globule, known as a bead, and heated in the Bun-
sen flame, the mass is colored in various tints according to the oxide
used (bead test). This action may be understood when we consider
that sodium metaphosphate takes up water to form primary sodium
orthophosphate : NaP08 + H20 — > NaH2P04. In the same way, but at
higher temperatures, it is able to take up oxides of elements other than
hydrogen, giving mixed orthophosphates. Thus with oxide of copper
a part of the metaphosphate unites according to the equation :
NaP08 + CuO -» NaCuP04,
and the product confers a green tinge on the bead.
Distinguishing Tests. — When a solution of nitrate of silver is
added to a solution of orthophosphoric acid or any soluble orthophos-
phate, a yellow precipitate of silver orthophosphate Ag3P04 is produced.
This is a test for phosphanion. With pyrophosphoric acid or any pyro-
phosphate the product is white Ag4P207. With metaphosphoric acid a
white precipitate, AgP03, is obtained also. Metaphosphoric acid coagu-
lates a clear solution of albumen, while pyrophosphoric acid has no
visible effect upon it.
A test for orthophosphoric acid, or rather the ion P04"', consists in adding a
drop of the solution containing this ion to a solution of ammonium molybdate (q.v.)
in dilute nitric acid. A copious yellow precipitate of an ammonium phospho-
PHOSPHORUS 469
molybdate, (NH4).,PO4,llMoO3, 6H2O, appears on warming. In presence of excess
of ammonia, the formation of the white insoluble ammonium-magnesium phosphate
(p. 467) serves as a test also. Arsenic acid (g.t>.) gives precipitates of appearance
and composition similar to these two.
Phosphorous Acid. — With cold water phosphorus trioxide (P406)
yields phosphorous acid very slowly. With hot water the action is
exceedingly violent and complex (p. 464). This acid may be obtained
also by the action of water upon phosphorus trichloride, tribromide,
or tri-iodide and evaporation of the solution :
PCL, + 3H20 -> P(OH)8 + 3HC1.
A certain amount of this acid, along with phosphoric acid and hypo-
phosphoric acid, is formed when moist phosphorus oxidizes in the air.
In spite of the presence of three hydrogen atoms, this acid is dibasic,
and two only are replaceable by metals. To express this fact, the first
of the following formulae is preferred :
since the symmetrical formula would indicate no difference between the
three hydrogen atoms. H united directly to P, as here and in PH8, is
not acidic. Phosphorous acid is a powerful reducing agent, precipitat-
ing silver, for example, in the metallic form from solutions of its salts.
When heated, it decomposes, giving the most stable acid of phosphorus
(cf. pp. 268, 274, 395), namely, metaphosphoric acid, and phosphine :
4H8P08 _> 3HP08 + 3H20 + PH8.
Hypophosphorous Acid. — The potassium salt of this acid is ob-
tained, as we have seen, when phosphorus is heated with potassium
hydroxide solution (p. 461). It may be prepared in the free form by
substituting barium hydroxide for potassium hydroxide :
3Ba(OH)2 + 8P + 61^0 -> 3Ba(H2P02)2 + 2PH8.
By careful addition of dilute sulphuric acid to the resulting liquid,
barium sulphate is precipitated. On evaporation of the water the
white crystalline acid, H8P02, is obtained. This acid is monobasic ;
two of its hydrogen atoms cannot be displaced by metals. To ex-
470 INORGANIC CHEMISTRY
r-H
press this fact the graphic formula 0 = P -< — H is used. This sub-
( -OH
stance is also a powerful reducing agent, tending, by the acquisition of
oxygen, to pass into phosphoric acid.
Structural Formulae of Salts of Hydrogen. — As a rule, the
formulae of acids have thus far been written with the ionizable hydro-
gen in front : HC1, H2SO4, HC2H802. This is only one illustration of
the method by which chemists have constantly sought to utilize form-
ulae for the purpose of expressing, not merely the composition of a sub-
stance, but some of its properties as well. By another typographical
device we have attempted to indicate the behavior of dilute solutions
by putting the radicals in brackets : Cu(N08)2, Ba(OH)2. These are
called structural or constitutional formulae, and their object is not to
show actual structure, but to exhibit the modes of action of the sub-
stance, by means of a supposed structure. Now the modes of action of
a single substance are often rather various, and one and the same
structural formula cannot represent all of these at once. We have
observed this, particularly, with the oxygen acids. Thus, H2.S04 ex-
presses the mode of activity in dilute solution and often when no solvent
is present, as in the action on chlorides (p. 178) and nitrates (p. 439).
But when all the hydrogen of an acid is not ionizable, we regard that
which is so as part of an hydroxyl group in the parent molecules, and
the rest as being attached to the characteristic non-metal of the acid,
for example, the phosphorus (cf. pp. 391, 392). Thus, we should write
phosphorous acid POH(OH)2, instead of H2P03H, to chronicle this fact.
So also the formula POH2(OH) is used for hypophosphorous acid.
Molecular actions, such as those of sulphuric acid S02(OH)2 (p. 389),
are well shown by these formulae :
SO2(OH)2 + 2HBr -> S02 + 2H20 + Br2.
It must be noted particularly that this sort of formula, when the sub-
stance for which it stands is an acid, represents only some features
in the behavior of the anhydrous substance and of the molecules, and
not the ionic action in solution. A formula like Ba(OH)2, where the
material is a base, on the other hand, represents both the ionic and the
molecular behavior. The graphic formula is more general (cf. p. 224).
It shows all these relations, and often still others, but none of them so
specifically.
PHOSPHORUS 471
Sulphides of Phosphorus. — Yellow phosphorus when heated with
sulphur unites with explosive violence. By using red phosphorus the
action can be controlled. By employing the proper proportions the
pentasulphide P2S5 is secured. It is purified by distillation from a
retort in which a current of carbon dioxide is maintained (see below).
The distillate solidifies to a yellow crystalline mass, which melts at
274° and boils at 530°. Materials undergoing chemical change, which
are to be kept at a constant, high temperature, are often placed in
tubes suspended in the vapor of the pentasulphide. When a lower
temperature is required, boiling sulphur (448°) is used.
Distillation in a stream of some inactive gas is a common means
of distilling under reduced pressure (cf. p. 276). The dilution of the
vapor lowers its partial pressure, just as would evacuation. This plan
has the advantage, however, of sweeping the vapor away from the
heated region into the condenser, and so diminishing the amount of
decomposition. In dealing with compounds of carbon, a current of
steam is often used for the above purposes. It enables us also to sepa-
rate a slightly volatile substance from one which is almost involatile.
Phosphorus pentasulphide acts upon water in the cold, and upon
substances containing hydroxyl when heated with them, the actions
being similar to those of the pentachloride (p. 463) :
P2S5 + 8H2O u» 2H3P04 + 5H2S.
Other sulphides, P4S8, P2S3, and P3S6, may be prepared by using the
constituents in the proportions represented by these formulae.
Comparison of Phosphorus with Nitrogen and with Sulphur.
— Although phosphorus and nitrogen are regarded as belonging to one
family, the differences between them are more conspicuous than the
resemblances. The latter are confined almost wholly to matters con-
cerned with valence. The differences are seen in the facts that nitro-
gen is a gas and exists in but one form, while phosphorus is a solid
occurring in two varieties, and that the former is inactive and the latter
active. The contrasts between phosphine and ammonia (pp. 461-2)
and between the halides of the two elements (p. 462) have been noted
already. The pentoxide of nitrogen decomposes spontaneously ; that
of phosphorus is one of the most stable of compounds. Nitric acid
is very active ; the phosphoric acids are quite the reverse.
On the other hand, the resemblance of phosphorus to sulphur is
472 INORGANIC CHEMISTRY
marked. Both are solids, existing in several forms. Both yield stable
compounds with oxygen and chlorine. The hydrogen compounds
interact with salts to give phosphides of metals and sulphides of
metals, respectively. Against these must be set the facts, that hydro-
gen sulphide does not unite with the hydrogen halides at all, and that
phosphoric acid is hard to reduce, while sulphuric acid is reduced with
comparative ease.
Exercises. — 1. Explain the effect of sulphuric acid in setting
fire to the earliest matches (p. 460).
2. Make a brief definition of a substance capable of sublimation
(p. 463).
3. Why would a mixture of potassium dichromate and hydrochloric
acid (p. 374) be less suitable than nitric acid for making phosphoric
acid from red phosphorus ?
4. Why is not the tertiary phosphate of sodium (p. 467) decom-
posed by heating ? What tertiary phosphates would be decomposed
by this means ?
5. Formulate the hydrolyses of the secondary and tertiary sodium
orthophosphates as was done for sodium sulphide (p. 375).
6. How should you prepare Ca2P4O7 and Ca(P08)2 ?
7. What product should you confidently expect to find after
heating (a) sodium phosphite, Na2HP08, (6) barium hypophosphite
(p. 469) ?
8. Compare the elements chlorine and phosphorus after the man-
ner of the comparisons on p. 471.
CHAPTER XXVIII
CARBON AND THE OXIDES OF CARBON
The Chemical Relations of the Element. — The elements of the
carbon family are carbon, silicon, germanium, tin, and lead. Of these
the first two are entirely non-metallic, while the others are metals
showing more or less strong resemblances to the non-metals. All these
elements are quadrivalent as regards the maximum valence which they
exhibit. With the exception of silicon, however, they all form com-
pounds in which they are bivalent.
The chemistry of the compounds of carbon is an exceedingly ex-
tensive and complex subject. It is commonly known as organic
chemistry, on account of the fact that the majority of the substances
composing, and produced by, living organisms are compounds of carbon,
and that it was at first supposed that their artificial production, e.g. with-
out the intervention of life, was impossible. But many natural organic
products have now been made from simpler ones or from the elements,
a process called synthesis, and the preparation of the others is de-
layed only in consequence of difficulties caused by their instability and
complexity. On the other hand, hundreds of compounds unknown to
animal or vegetable life, including many valuable drugs and dyes, have
now been added to the catalogue of chemical compounds.
The elements entering into carbon compounds are chiefly hydrogen
and oxygen. After these, nitrogen, the halogens, and sulphur may be
named.
CARBON.
Occurrence. — Large quantities of carbon are found in the free
condition in nature. The diamond is the purest natural carbon, and
at the same time the least plentiful. Graphite, or plumbago, which is
the next purest, is found in limited amounts, and is a valuable mineral.
Coal occurs in numerous forms containing greatly varying proportions
of free carbon. Small quantities of the free element have been found
in meteorites.
In combination, carbon is found in marsh-gas, or methane CH4,
which is the chief component of natural gas. The mineral oils consist
473
474 INORGANIC CHEMISTRY
almost entirely of mixtures of various compounds of carbon and hydro-
gen. Whole geological formations are composed of carbonates of com-
mon metals, particularly calcium carbonate or limestone, and a double
carbonate of calcium and magnesium, known as dolomite.
The Diamond. — The varieties of carbon differ very markedly in
their physical properties, and to some extent also in their chemical
behavior. Diamonds, which are found in India, Borneo, Brazil, and
South Africa, are scattered sparsely through metamorphic and volcanic
rocks which seem to have undergone secondary changes. They are
covered with a crust which entirely obscures their luster, and possess
natural crystalline forms belonging to the regular system. A form re-
lated to the octahedron is frequently observed. It should be noted that
this natural form bears no relation whatever to the pseudo-crystal-
line shape which is conferred upon
the stone by the diamond-cutter.
Thus, a "brilliant" possesses one
rather large, flat face, which forms
the base of a many sided pyra-
mid (Fig. 91, showing two views).
This form is given to the stone, in
FIG. 91. order that the maximum reflection
of light from its interior may be
produced. The diamond is harder than any other variety of matter,
with the exception, perhaps, of one compound of boron, while only one
or iwo other materials, like carborundum, approach it. It may be re-
marked in passing, that the hardness of a substance is measured, on an
arbitrary scale, by the way in which it is able to scratch smooth sur-
faces of other bodies. The corner of one of the natural diamond crys-
tals will scratch the surface of almost every other substance, while its
surfaces in turn are scratched by carbide of boron alone. Its specific
gravity is about 3.5, and it is the densest form of carbon. Few materi-
als are capable of dissolving any of the forms of carbon. Molten iron
(y.v.) dissolves five or six per cent, part of which goes into combina-
tion, and a few other substances at high temperatures dissolve much
smaller quantities. The diamond is a nonconductor of electricity.
The largest diamond known, the Jubilee, was exhibited at the Paris
Exposition of 1900, and weighed 49 g. The Kohinoor weighs 22 g.
The diamond, although its origin in nature is still a matter of uncer-
tainty, has been made artificially in several ways. Moissan (1887)
ro-
CARBON AND THE OXIDES OF CARBON 475
dissolved carbon in molten iron and, after chilling the mass so as to
produce a solid crust, which by its shrinkage severely compressed the
interior, allowed the whole to cool very slowly. Portions of the
interior of the ingot were treated with acid to dissolve the iron, and
amongst the insoluble particles were recognized a few microscopic
fragments (none larger than 0.5 mm.) which exhibited the form and
hardness of the diamond. The greater part of the carbon, however,
appeared as graphite.
That the diamond contains nothing but carbon, is shown by the
fact that when burned it produces nothing but carbon dioxide.
Graphite. — Graphite (Gk. ypa<£o>, I write) is found in Cumber-
land, Siberia, Ceylon, and elsewhere. Good crystals are seldom
found, but the form appears to belong to the hexagonal system. The
mineral is extremely soft, in utter contrast to the diamond, and has a
smaller specific gravity (about 2.3). It conducts electricity. It is now
made artificially by an electro-thermal process (cf. p. 457). A power-
ful alternating current is passed through a mass of granular anthracite,
and the latter, although not melted by the high temperature, is largely
converted into graphite.
Graphite is now used exclusively for making the anodes in the
electrolytic manufacture of chlorine and in related processes. Mixed
with fine clay it forms the " lead " of lead pencils. As " black-lead ",
it is employed to protect ironware from rusting. It takes the place of
oil as a lubricant in cases where the former would be decomposed by
heat. The fine, slippery scales, which it forms when pulverized, fill
the inequalities in the bearings, and glide over one another with little
friction.
Amorphous Carbon. — This is the name given to the varieties of
the element which have no crystalline form. They vary in specific
gravity up to 1.9. Coke, which is now manufactured in immense
quantities by heating coal until all the volatile matter has been dis-
tilled off, is a very dense variety of amorphous carbon. It is used in
the reduction of iron ores. The value of coking lies largely in the
removal of the sulphur, originally contained in the coal in the form of
pyrite.
By the imperfect combustion of heavy oils and resins, in which
the flame plays upon a cooled surface, a finely divided form of carbon
476 INORGANIC CHEMISTRY
known as lampblack is produced. This is much used in the manu-
facture of printer's ink.
Charcoal is chiefly made by the heating of •wood out of contact
with air. In the more refined forms of the process the charring is con-
ducted in retorts, and the materials which distil off are used in vari-
ous ways. Wood consists largely of cellulose (C6H1006)n, incrusted
with lignin and holding much moisture and resinous material. The
products of its distillation are partly gaseous and partly fluid. The
gases, consisting mainly of hydrogen, methane CH4,- ethane C2H6, ethy-
lene C2H4, and carbon monoxide CO, are employed, on account of their
combustibility, as fuel in the distillation itself. The fluids form a
complex mixture containing large quantities of water, wood spirit, or
methyl alcohol CH8OH, acetic acid, acetone (CH8)2CO, and tar. Wood
charcoal exhibits the cellular structure of the material from which it
was made, and is therefore exceedingly porous. The original mineral
constituents of the wood appear in the ash of the charcoal when the
latter is burned.
For certain purposes, charcoals made in the same fashion as the
above from bones and from blood, find wide application. The former,
called bone black (p. 456), contains much calcium phosphate. In the
chemical laboratory, pure carbon is made, as a rule, by the charring of
sugar (cane-sugar, C^H^OjJ. The sugar is purified from mineral
matter, before use, by crystallization from water.
The tendency of almost all carbon compounds to char, when
heated, is used as a means of recognizing their presence.
Properties of Charcoal. — Charcoal exhibits certain properties
which are not shared to any extent by other forms of carbon. For
example, it can take up large quantities of many gases. Boxwood
charcoal will in this way absorb ninety times its own volume of am-
monia, fifty-five volumes of hydrogen sulphide, or nine volumes of
oxygen. Freshly made dogwood charcoal (used in making the
best gunpowder), when pulverized immediately after its preparation,
often catches fire spontaneously on account of the heat liberated by the
condensation of oxygen. It is therefore set aside for two weeks, to
permit the slow absorption of moisture and air, before being ground up.
The absorbed gases may be removed unchanged by heating the charcoal
in a vacuum. The disappearance of these immense quantities of gas
into small pieces of charcoal is described as adsorption, and is caused
by the adhesion of the gases to the very extensive internal surface
CARBON AND THE OXIDES OF CARBON
477
which the charcoal possesses. Solid and liquid bodies are also in
many cases taken up by charcoal in a similar fashion. Thus, strychnine
may be removed from an aqueous solution by agitation of the latter
with charcoal. In the manufacture of whiskey (q.v.), the fusel oil, which
is extremely harmfiil, is in many cases removed by nitration of the
diluted spirit through charcoal, before rectification. Organic coloring
matters, such as litmus and indigo, belong to the class of bodies thus
extracted from solution by charcoal. In the refining of sugar the
syrup is boiled with charcoal for the purpose of removing a brown
resin, in order that the product may be perfectly white. It is, in part,
upon this property that we rely, also, in the employment of charcoal
filters. The organic materials dissolved in the drinking water under-
go adsorption in the charcoal. In this connection, however, it must
be remembered that the quantity which a given mass of charcoal may
take up is limited, and that careful cleansing is required in order that
the efficiency of the filter may be maintained.
Coal. — Peat, brown coal, soft coal, and anthracite .represent, in a
general way, different stages in the decomposition of vegetable matter
in absence of air. Water and compounds of carbon and hydrogen are
given off in the process. The ratio which the carbon combined with
oxygen and hydrogen bears to the free carbon decreases in the order in
which the substances stand above. The following table shows this
change in composition and the relations of the substances to fresh
wood on the one hand and charcoal and coke on the other :
Percentage, excluding Ash
and Moisture, of :
Percentage
Ash.
Per-
centage of
Calorific
Value
C
H
0
N
(Air Dried).
Calories.
Wood . .
45
6
48
1
1.5
18-20
2700
Peat . .
60
6
32
2
6—20
20-30
3500
Brown coal
70
5
24
1
3—30
16
30-6000
Soft coal .
82
5
12
1
1—16
4
66-8000
Anthracite
94
3
3
. B
1.6
2
70-8000
Charcoal .
95
1.7
3.4
4
6.5
8080
Coke . .
96
0.7
2.5
i
3.4—11
2
7700
Chemical Properties of Carbon. — Diamond, graphite, and
amorphous carbon probably differ from one another, not merely in
physical properties, but also chemically. Certainly the stability of
478 INORGANIC CHEMISTRY
compounds containing many units of carbon in their molecules indi-
cates a great tendency of carbon to combine with itself, and gives
plausibility to the belief that the molecule of free carbon may itself be
complex. Differences in the size or structure of these complex mole-
cules would account for the variety in the forms of the element.
Amorphous carbon is the least stable of the three, for it liberates most
heat in entering into combination. Since graphite is formed at high
temperatures, and diamonds turn into a black mass under the same
conditions, we may presume that graphite is the most stable, at least
at 3000°.
The most common uses of carbon depend upon its great tendency
to unite with oxygen, forming carbon dioxide. Under some circum-
stances carbon monoxide is produced. Aside from the direct employ-
ment of this action for the sake of the heat which is liberated, it is
used also in the reduction of ores of iron, copper, zinc, and many other
metals. When, for example, finely powdered cupric oxide and carbon
are heated, copper is obtained. The gas given off is either carbon
dioxide, or a mixture of this with carbon monoxide, according to the
proportion of carbon used :
CuO + C -> Cu + CO,
2CuO + C -> 2Cu + C02.
Carbon unites directly with hydrogen very reluctantly. When an
electric arc is produced between carbon poles in a tube through which
a stream of hydrogen passes, acetylene C2H2 is formed. The presence
of this gas may be shown by the luminosity its combustion confers
on the hydrogen flame. This substance can form the starting-point
for the artificial preparation of many carbon compounds, and its syn-
thesis possesses therefore a certain interest.
At the high temperatures produced in the electric furnace, carbon
unites with many metals and some non-metals. Compounds formed
in this way are known as carbides, such as aluminium carbide A14C3,
calcium carbide CaC2, and carborundum CSi.
Calcium Carbide. — This compound, which is colorless whe:
pure, is manufactured in an electric furnace, by the interaction of finely
pulverized limestone or quicklime with coke :
CaO + 3C -> CaC2 + CO.
The operation is a continuous one, the materials being thrown into the
left side of the drum (Fig. 92, diagrammatic), and the product re-
CARBON AND THE OXIDES OF CARBON
479
moved on the right. The carbon poles are fixed. The arc having
been established, the drum is rotated slowly as the carbide accumu-
lates. The current enters by one carbon, passes through the carbide,
and leaves by the other. The high resistance of the partially trans-
formed material causes the production of the heat. When the action
in one layer approaches completion, the resistance falls, the current
increases, and an armature round which the wire passes (not shown in
Fig. 92) comes into operation and turns the drum. In this way the
carbide just formed is continuously moved away from the carbons, and
new material, introduced on the left, falls into the path of the current.
The iron plates which form the circumference of the drum are added
on the left and removed on the
right, where also the carbide is
broken out with a chisel. The
drum revolves once in about
three days. The product is
used for making acetylene (q-v.).
CARBON DIOXIDE AND
CARBONIC ACID.
Occurrence. — Carbon di-
oxide is present in the atmos-
phere, and issues from the
ground in large quantities in
certain neighborhoods, as, for
example, near the Lake of
Laach, in the so-called Valley of
Death in Java, and in the Grotta del Cane near Naples. Effervescent
mineral waters contain it in solution, and their effervescence is caused
by the escape of the gas when the pressure is reduced. Well-known
waters of this kind are those of Selters (whence, by a singular per-
version, the English word seltzer is derived) and of the Geyser
Spring at Saratoga.
Modes of Formation. — Carbon dioxide is produced by combus-
tion of carbon in the presence of an excess of oxygen :
C + 02 -> C02.
The combustion of all compounds of carbon, as well as the slow oxida-
FlG. 92.
480 INORGANIC CHEMISTRY
tion in the tissues of plants and animals, leads to the formation of the
same product.
It was Joseph Black (1757) who first recognized the gas as a
distinct substance. He observed its formation when marble or mag-
nesium carbonate was heated :
CaC03 «=» CaO + C02,
and named the gas " fixed air " from the fact that it was contained in
these solids. The above action had been used for centuries in making
quicklime (calcium oxide). All common carbonates, excepting the
normal carbonates of potassium and sodium, decompose in this way,
leaving the oxide of the metal.
Black found that the gas was also produced when acids acted upon
carbonates, and this method is commonly employed in the laboratory :
CaC08 (solid) ±? CaC03 (diss'd) *=? Ca"+C0," ) rn
2HC1 (diss'd) *=> 2CP + 2H' J ^ **'()**=* ^°
Since the carbonic acid is very slightly ionized, the action is like
that of acids on sulphites (p. 398). The carbonate of calcium, how-
ever, is very slightly soluble, so that an additional equilibrium controls
its solution. In this respect the action is like that of acids on ferrous
sulphide (p. 371).
Carbon dioxide is also a product of the fermentation of sugar (q.v.),
as Black had the credit of showing.
Physical Properties. — Carbon dioxide is a colorless, odorless
gas. It is much heavier than air. The G.M.V. weighs 44 g. The
critical temperature is 31.1°. Liquefied carbon dioxide boils at —79°.
The sp. gr. of the liquid at 0° is 0.95. At 0° its vapor tension is
35.4 atmospheres and at 20° 59 atmospheres. It must be preserved,
therefore, in very strong, wrought-iron cylinders. Large quantities of it,
often collected from fermentation vats, are sold in such cylinders, and
used in operating beei-pumps and in making aerated waters. When
the liquid is allowed to flow out into an open vessel it cools itself by
its own evaporation and forms a white, snowlike mass. Solid carbon
dioxide evaporates without melting (cf. p. 463). A mixture of solid
carbon dioxide with ether is frequently used as a freezing mixture
(—80°). The ether is employed to secure better contact with the body
to be cooled.
CARBON AND THE OXIDES OF CARBON 481
The great contrast in the speeds of a chemical change at two
cemperatures (cf. p. 72) may be illustrated by putting a minute piece
D£ sodium in some 30 per cent hydrochloric acid which has been
cooled in the above mixture. Hardly any interaction can be observed.
But if the temperature of the acid is allowed to rise, the action becomes
more and more rapid, and ends by being explosively violent.
Carbon dioxide gas under a pressure of 760 mm. and at a tempera-
ture of 15° dissolves in its own volume of water. Up to pressures
of four or five atmospheres Henry's law describes its solubility accu-
rately. An aqueous solution, prepared under a pressure of 2-3 atmos-
pheres, is familiarly known as soda water.
Chemical Properties. — Carbon dioxide is a stable compound,
and shows a little dissociation only at the highest temperatures :
2C02 +± 2CO + 02.
The most active metals, such as potassium, sodium, and magnesium,
burn brilliantly when heated in carbon dioxide, producing the oxide
of the metal and free carbon. Less active metals, such as zinc and
iron, give an oxide of the metal and carbon monoxide (q.v.}.
Carbon dioxide unites directly with many oxides, particularly those
of the more active metals, such as the oxides of potassium, sodium,
calcium, etc. Hence the decomposition of calcium carbonate by heat-
ing (p. 480) is a reversible action, which proceeds in the opposite direc-
tion when a sufficient pressure of carbon dioxide is employed (cf. p.
256 and Chap. xxxv).
Carbon dioxide, when dissolved in water, forms an unstable acid :
C0 <=± HC0.
The name carbonic acid is frequently, though improperly, given to
the gas itself, which is really the anhydride of the acid and has no
acid properties.
Chemical Properties of Carbonic Acid. — The solution of car-
bon dioxide in water exhibits the properties of a weak acid. It con-
ducts electricity, although not well. It turns litmus red, though not
so decidedly as do strong acids. Its feebleness is due, however, not
exclusively to the small degree of ionization, but also to the fact that
ordinary solutions of carbon dioxide are necessarily very dilute. The
iouization takes place chiefly according to the equation :
482 INORGANIC CHEMISTRY
In a deci-normal solution, less than two molecules of the acid in a
thousand are ionized. The conditions of equilibrium between the gas
and the solution are precisely similar to those described under sul-
phurous acid (p. 393).
Carbonates. — When excess of an aqueous solution of carbonic
acid is mixed with a solution of a base like sodium hydroxide, or, as
the operation is more usually performed, when carbon dioxide is
passed directly into a solution of the alkali, water is formed and the
carbonate remains dissolved:
H2C03 + NaOH «=> H2O + NaHC03, or H* + OH' -> H20.
The product is sodium hydrogen carbonate (sodium bicarbonate).
Although technically an acid salt, its solution is neutral on account of
the exceedingly slight dissociation of the HCOS' ion. By addition of
an equivalent of sodium hydroxide to the solution of the bicarbonate
the normal carbonate is obtained :
NaOH + NaHC08 <=» H20 + Na^COg.
This solution is alkaline in reaction, for the same reason that a solution
of secondary sodium orthophosphate is so (cf. p. 466).
The carbonates, with the exception of those of potassium, sodium,
and ammonium, are insoluble in water, and may be obtained by precipi-
tation when the proper ions are employed. For example :
MgS04 + Na2C03 +± MgC03 [ + Na^SO, or Mg" + C03" ^± MgC08|.
The aqueous solution of carbon dioxide interacts with solutions of
barium and calcium hydroxides in a similar manner :
Ca(OH)2 + H2COS fcj CaC03 | + H2O.
The formation of precipitates with these solutions is used as a test for
carbon dioxide and a means of estimating its amount in a sample of
air (q.v.~).
Excess of carbon dioxide converts calcium carbonate into the more
soluble bicarbonate, and hence considerable quantities of " lime " are
frequently held in solution by natural waters :
HsCOg + CaC03 ^± H2Ca(C03)2.
A considerable excess of carbon dioxide is required to convert the
whole of the carbonate into the soluble bicarbonate, since the action
CARBON AND THE OXIDES OF CARBON 483
is markedly reversible. In the same fashion, the carbonates of iron
(FeCO3), magnesium, and zinc are somewhat soluble in water con-
taining free carbonic acid. In fact, the solution, transportation, and
deposition of all these carbonates take place in nature on a large scale
by the alternate progress and reversal of this action. Water contain-
ing calcium carbonate in solution is known as hard water (q.v.), that
containing ferrous carbonate as chalybeate, or iron water.
Hole of Chlorophyl-bearing Plants in Storing Energy. —
While in plants the same consumption of oxygen and production of
carbon dioxide goes on as in animals, only with less rapidity, an action
which is in a general way a reversal of this takes place at the same
time. The chlorophyl and protoplasm in the leaves of the plant have
the power of taking up carbon dioxide. Part of the oxygen is restored
to the air, and the rest of the substance, including all the carbon, is
used by the plant as food. This operation goes on only in sunlight
(see below). The details of the chemical changes are not thoroughly
understood, but the various chemical compounds which plants con-
struct in large quantities, of which, sugar, starch, and cellulose are
prominent examples, are built up as the result of this action. In a
rough fashion, and disregarding the steps by which the process takes
place, we may represent the chemical change by means of the thermo-
chemical equation :
6C02 + 5H20 -+ C6H10O5 + 602 - 671,000 calories.
Since the production of carbon dioxide by the combustion of any
organic compound gives out heat, the partial reversal of this combus-
tion, of which the green parts of the plant are the scene, requires the
expenditure of energy, and the source of this energy is to be found in the
sunlight. The difference in total energy between water and carbon
dioxide, on the one hand, and the cellulose, starch, or sugar and free
oxygen on the other, is very considerable. The above figures indicate
roughly (p. 79) this difference (= the amount of energy stored) for
cellulose, and the values for the other compounds are of the same
order.
The importance of this remarkable endothermal action, involving
the storing of the energy of sunlight, is very great. Aside from a
little work done by water-power, the whole energy used by man and
by animals comes from the reversal of it. The compounds forming
the structure of the plant are employed in several different ways for
484 INORGANIC CHEMISTRY
this purpose. When consumed by herbivora as food, the chemical
changes which they undergo furnish the energy necessary for the con-
tinued life of the organism. The whole of the material is not at once
reduced to the State of carbon dioxide, but passes into other forms of
combination, which in turn become the food of the carnivora, or flesh-
eating animals. By the oxidation to carbon dioxide which takes place
in the bodies of these animals, the process of exhausting the possible
energy of the carbon compounds is completed.
In another fashion we secure energy from the materials of plants
by burning wood and employing the heat thus produced. In still
another fashion, after the wood has undergone partial decay and
conversion into coal, we secure the remaining energy which this con-
tains by its final combustion in the furnace of the steam-engine.
It should be noted that the energy in the last case is not stored exclusively in the
coal, but is shared between carbon and the oxygen of the air. If our atmosphere
consisted of compounds of carbon, then the material corresponding to stores of coal
would have to be oxygen or compounds of oxygen, and we should be likely then
to speak of the energy as being stored for us and sold in the form of oxygen. That
we are in the habit of speaking of it, at present, as going with the carbon is because
the oxygen of the air is supplied free of charge, while the coal and wood have
to be purchased.
Photochemical Action. — We have seen that light may simply
act catalytically, as on a mixture of hydrogen and chlorine (p. 174) or
an aqueous solution of hypochlorous acid (p. 268). These actions in-
volve the liberation of energy and go on spontaneously (cf. p. 271) under
proper conditions. On the other hand, light may actually be consumed
in large amount in producing a chemical change, as in decomposing
silver chloride (p. 14), or in the above instance. All wave lengths
of light, which is the same as to say all colors of light, are not equally
active in any one case. But there is no particular set of wave lengths
which is of special chemical activity. In the action on silver chloride,
green and bhie light is very active, while red is almost without effect.
Here, in the actions in which chlorophyl is concerned, it is the red
and yellow light that produces the chemical change, and a plant ex-
posed to blue light (e.g., by shading with blue glass) will assimilate
none of the carbon dioxide in the air surrounding it. The chemical
substances in the retina of the eye seem to resemble those in the leaves
of plants, for they are most affected by red and yellow light. To put
this another way, a spectrum of uniform intensity throughout, when.
CARBON AND THE OXIDES OF CARBON 485
viewed by a plant or a human eye, would appear to be brightest in
the red and yellow portions, while a considerable stretch towards the
blue extremity would actually be invisible. On the other hand, to an
eye in which the active substance was silver chloride, if such an eye
could be imagined as existing, the red end would be invisible and the
blue and ultra-violet would be the most brilliant parts.
CARBON MONOXIDE.
Preparation. — Carbon monoxide is formed in many industrial
operations. We commonly observe the blue flame of burning carbon
monoxide playing on the surface of a coal fire. The gas is produced
by the passage of the carbon dioxide, which is first formed/through the
upper layers of heated coal :
C02 + C -» 2CO.
A similar reduction of carbon dioxide is produced by metals such as
zinc, when a moderate heat is applied :
C02 -f Zn -> ZnO + CO.
On a large scale, a mixture of carbon monoxide and hydrogen is
prepared as the basis of -water gas. Steam is turned into an iron
cylinder lined with fire clay and filled with vigorously burning coke :
C + H20 -> CO + H2.
The products are both combustible, and, by the addition of substances
which burn with a luminous flame, the mixture is used for the manu-
facture of illuminating-gas (g.v.*).
In the laboratory, carbon monoxide is frequently obtained by heat-
ing oxalic acid, a solid, white, crystalline substance, with concentrated
sulphuric acid. The latter is here employed simply as a dehydrating
agent (p. 388), so that it need not be included in the equation :
H2C204 -> C02 + CO + H20.
To obtain pure carbon monoxide from this mixture it is necessary to
remove the carbon dioxide and to dry the gas. The carbon dioxide
may be absorbed by passing the gas through a concentrated solution
of potassium hydroxide. By treatment of formic acid, or sodium for-
mate, with sulphuric acid, the .presence of the carbon dioxide may be
avoided :
HCHO2 -> CO + H20.
486 INORGANIC CHEMISTRY
Physical Properties. — Carbon monoxide is a colorless, tasteless,
odorless gas. It is very slightly soluble in water. Its density is
almost the same as that of air, for the G.M.V. weighs 28 g. When
liquefied it boils at — 190°.
Chemical Properties. — All the chemical properties of carbon
monoxide are referable to the fact that in it the element of carbon
appears to be bivalent : C = 0. The compound is in fact unsaturated,
and combines with oxygen, chlorine, and other substances directly.
Thus the gas burns in the air, uniting with oxygen to form carbon
dioxide. Again, iron (q.v.) is manufactured by the reduction of the
oxide of iron by gaseous carbon monoxide in the blast furnace :
Fe2O8 + SCO <=> 2Fe + 3C02.
In sunlight carbon monoxide unites directly with chlorine to form
car bony 1 chloride COClj. It is absorbed by a solution of cuprous
chloride in hydrochloric acid or ammonium hydroxide, forming a com-
pound whose composition is probably represented by the formula
Cu2Cl2,CO,2H20. It unites directly with certain metals, notably nickel
and iron, with which it forms the so-called nickel carbonyl and iron car-
bony 1, respectively. The former is a colorless, volatile liquid Ni(CO)4.
The gas is an active poison. When inhaled it unites with the
haemoglobin of the blood to the exclusion of the oxygen or carbon
dioxide, each of which forms less stable compounds (cf. p. 432). A
quantity equivalent to about 10 c.c. of the gas per kilo, weight of the
animal is sufficient to produce death, about one-third of the whole
haemoglobin having entered permanently into combination with carbon
monoxide.
The quantities of heat given out by the successive unions of two units of oxy-
gen with one unit of carbon are worth recording :
C + O — > CO + 2630 calories,
CO + O — » C02+ 6830 calories.
It will be seen that the addition of the second atom of oxygen appears to cause the
evolution of a very much larger amount of heat than does that of the first. It
must be remembered, however, that the carbon monoxide is gaseous, while the
carbon in the first equation is solid, and probably in a condition of complex molec-
ular aggregation. The heats produced by the unions of the two units are probably
not very different, but in the first case a large amount of the heat is used up in
disintegrating the carbon and bringing it iuto the gaseous condition.
CARBON AND THE OXIDES OF CARBON 487
CARBONYL CHLORIDE AND UREA.
Carbonyl Chloride. — This substance is also named phosgene
(Gk. <£ws, light ; yewSv, to produce), on account of its formation by the
catalytic influence of sunlight (p. 484). On a commercial scale it is
obtained by passing the mixed carbon monoxide and chlorine over
animal charcoal, which assists the union catalytically. It is a liquid
which boils at 8°, possesses a suffocating odor, and is very soluble in
benzene and some other hydrocarbons. When brought into contact
with water it is hydrolyzed at once, forming carbonic acid and hydro-
chloric acid :
COCL, + 2H20 -> H2C08 + 2HC1.
Urea. — When ammonia and carbonyl chloride are mixed in the
proper proportions, urea, a most interesting chemical substance, is
produced :
Cl H-NH2 NH2
0 = CX -(- -* 0 = C/ +2HC1
NC1 H-NH2 XNH2
Excess of ammonia has to be used to combine with the hydrogen
chloride thus set free, so that the final equation is :
COCL, + 4NH8 --> CO(NH2)2 + 2NH4C1.
The urea, a white, crystalline solid, is soluble in alcohol, while ammo-
nium chloride is not, so that the former may be washed out by means of
this solvent and recovered by evaporation. A little reflection will
show that, using the above action as the final stage, urea can be built
up from the simple substances composing it.
Urea was known long before any method for its synthesis had
been discovered. It is the chief product of the decomposition of com-
pounds of nitrogen in the animal body, and is found in the liquid
excrements of animals. It was regarded as a typical organic substance,
in the old sense of the word (p. 473). In 1828 Wb'hler succeeded in
preparing it artificially (see below). This was the first synthesis by a
chemist, of a true " organic " substance, and its preparation proved to
be the precursor of many discoveries of a similar nature. From a
later year, about 1840, we may date the transition of organic chem-
istry, a science in which the mystery of life was supposed to be
supreme, into the chemistry of the compounds of carbon, which is a
branch of inorganic chemistry.
. INORGANIC CHEMISTRY
WOhler used ammonium cyanate (q.v.), a substance in whose preparation we
are independent of all products of life processes. When ammonium cyanate, or a
mixture of any ammonium salt with potassium cyanate in solution in water, is
warmed for some time, an intramolecular change (cf. p. 16) takes place, and long
prisms of urea are deposited as the liquid cools :
NH4.CNO <=» CO(NH2)r
Since the action is reversible, about four or five per cent of the ammonium cyanate
remains unchanged.
The two substances just mentioned are entirely different in chemical proper-
ties. Ammonium cyanate is a highly ionized salt, while urea is not a salt at all,
but a substance like ammonia which unites with acids to form salts. Materials
which, like these, have the same convpnsjt'011 and the same numbers of units in
their molecules, and yet possess different properties, are spoken of in chemistry as
isomers. The formulae we have employed attempt to explain the differences in
their properties by suggesting a difference in their molecular structure (cf. p. 224).
Assisted by the catalytic action of certain ferments, urea, when dis-
solved in water, can take up two molecules of the solvent to form
ammonium carbonate :
2H.O -+ (NH4)2C08 <=» 2NH8 + H20 + C02.
Ammonium carbonate (g'.v.) is a somewhat unstable compound, and in
turn, gives off ammonia and carbon dioxide. To this action is due in
part the pronounced odor of ammonia arising from the decomposition
of sewage.
Carbon Disulphide. — This compound is used in inorganic chem-
istry chiefly as a solvent. It is made by direct union of sulphur
vapor and glowing charcoal. An electro-thermal method of carrying
this out employs a furnace like that in Fig. 90 (p. 457). The sub-
stance comes off as a vapor and is condensed.
Carbon. disulphide is a colorless, highly refracting liquid. When
pure it possesses a pleasant odor, but traces of other compounds give
the commercial article a disagreeable smell. It boils at 46° and burns
in air, forming carbon dioxide and sulphur dioxide. Iodine, phos-
phorus, sulphur, rubber, and other substances dissolve freely in it.
Exercises. — 1. To which of the factors in the interaction of cal-
cium carbonate and hydrochloric acid (p. 480) is due the forward dis-
placement of all the equilibria ?
2. What will be the excess of pressure inside a bottle of soda-
CARBON AND THE OXIDES OF CARBON 489
water when four volumes of carbon dioxide are dissolved in one
volume of water ?
3. What volume of liquid carbon dioxide, measured at 0°, will be
required to give 75 liters of the gas at 0° and 760 mm. pressure ?
4. What will be the effect of increase in pressure on the dissocia-
tion of carbon dioxide (p. 481) ?
5. Prepare a diagram showing the whole scheme of equilibria in-
volved in the hydrolysis of sodium carbonate (p. 482).
CHAPTER XXIX
SOME CARBON COMPOUNDS
THE compounds of carbon with hydrogen are called hydrocarbons.
Those containing oxygen as well are divided into numerous and
extensive groups according to their behavior. Thus there are acids
like acetic acid, carbohydrates like sugar and starch, alcohols like
common (ethyl) alcohol, esters like ethyl acetate and fat, ethers like
common (ethyl) ether, and bodies related to cyanogen like prussic
acid. We can discuss only one or two examples from each of the
groups named.
THE HYDROCARBONS.
More than two hundred and fifty compounds of carbon and hydro-
gen have been described. They fall into several distinct series, the
chief one of which contains methane CH4 as its simplest member.
On account of the fact that certain members of this set are found in
paraffin, it is commonly known as the paraffin series. For the reason
that in this series the carbon has all its four valences employed, the
members are also called the saturated hydrocarbons.
Paraffin Series of Hydrocarbons. — The following list gives
the formulae of a few members of this series, with their names and
their boiling-points or melting-points :
CH4 .
C2H(i.
CH.
C6H14
C7H16
methane b
ethane
propane
-P-
- 164° ^
- 89.5° 1
- 37° \
butane
i
1° J
pentane ... . .
i
35° ")
hexane
4
71°
heptane
4
99° I
Gases
Liquids
490
SOME CARBON COMPOUNDS 491
^H.,, . . hexadecane m.-p. 18° 1 b.-p. 287.6°
, . . pentatriacontane ..." 74.7° ^Solids
2 . . hexacontane (dimyricyl) . " 102° J
It will be noted that the hydrocarbons up to butane are gases. From
pentane to hexadecane they are liquids. The remainder are solids.
In composition each is related to the preceding one by containing the
additional units CH2. The formula of any member of the series is
therefore representable by the expression CnH2B+2. Substances re-
lated in this way form an homologous series. Their relations will be
more clearly perceived if we employ the graphic formulae. Since
hydrogen appears uniformly to be univalent, the carbon must form
the backbone of each of the compounds. The formulae of the first
members are therefore as follows :
H H H H H H
I II III
H-C-H H-C-C-H H-C-C-C-H
I II III
H H H H H H
Transferences of H one step to the right and interpositions of CH2
constitute the successive differences.
A large number of these substances occur in nature. Methane is
present in large, and ethane in small, proportion in the natural gas of
Pennsylvania and Ohio. Many of the others occur in petroleum.
Petroleum. — This oil consists of a mixture of the liquid and
solid members of the series in varying proportions, and is found in
many parts of the United States, in Ontario, at Baku on the Caspian,
in India, and in Japan. In oil-refining, advantage is taken of the
differences in the boiling-points to make a partial separation of the
components by fractional distillation (see below). The compounds
containing sulphur which are often present, and would give the ob-
noxious sulphur dioxide when the oil was burned, are deprived of
this constituent by heating the oil with powdered cupric oxide. The
unsaturated hydrocarbons (q. v.} are removed by agitation with con-
centrated sulphuric acid. The following are some of the products of
the oil refinery, with their components and uses.
492
INORGANIC CHEMISTRY
Name.
Components.
B.-P.
Uses
Petroleum ether
Gasolene
Naphtha
Benzine
Pentane-hexane
Hexane-heptaue
Hep tail e-octane
Octaiie-nonane
40°- 70°
703- 90°
80°-120D
120°-150°
Solvent, gas-making
1( 11
" fuel
11
Kerosene
Decaue-hexadecane
150°-300D
Illuminating-oil
The portions of still higher boiling-point are employed as lubricating
oils.
The vapor of these products is more inflammable the more volatile
the components. The sale of kerosene is controlled legally by the re-
quirement that the vapor it gives when heated shall not catch fire
from a naked flame until the oil has reached a certain minimum tem-
perature, the " flash point." This varies from 44° to 68.5° in different
states and countries.
By cooling the residues from the retorts with a freezing mixture
(cf. p. 164), some of the solid members of the series, C^H^ to C28H68,
are obtained as white flakes, which are separated by nitration in
presses. This material forms the paraffin used in waterproofing paper,
in laundry work, and as an ingredient in candles. In some cases
vaseline, consisting of substances melting at 40°-5CP, C^H^ to C^H^,
is obtained also.
From ozocerite, which is a sort of natural paraffin, ceresin, a sub-
stitute for beeswax, is made. Asphalt is another natural mixture of
hydrocarbons.
The formation of these hydrocarbons in nature is not yet thor-
oughly explained. According to one theory, they are formed by the
action of water upon carbides of metals ; while according to another,
they result from the decomposition of vegetable or animal matter.
Possibly both of these sources have contributed to their formation.
Certain differences between the natural oils of different localities point,
at all events, to some difference in their origin.
Fractional Distillation. — When the boiling-points of two com-
ponents of a liquid are very far apart, the vapor pressure of the one
may be very low when that of the other, by heating, has reached 760
mm. In this case the first distillate will contain little of the high-
boiling component. When, as in the case of petroleum, the differ-
ences in boiling-points are not great, complete separation of the
SOME CARBON COMPOUNDS 493
components is impossible. Yet by distillation in which the distillate
is caught, not in one vessel, but in several successively, " fractions "
are obtained such that the earlier ones contain more of the low-boiling
and the later ones more of the high-boiling materials. The vessels
are changed when the thermometer immersed in the vapor (Fig. 16,
p. 38) reaches certain temperatures. When these fractions are then
distilled one at a time, beginning with the lowest, and the several dis-
tillates are divided from one another by the same temperatures as
before, a more complete separation is effected. This process is called
fractional distillation, and may be repeated as often as we please with
constantly increasing differentiation of the fractions.
An experimental illustration may be given by mixing 0.4 c.c. of
benzene (b.-p. 80.4°) with 8 c.c. formic acid (b.-p. 100°) and 2 c.c.
benzyl alcohol (b.-p. 206.5°) and boiling a part of the mixture in a
test-tube with a small flame. The components come off in succession,
and are recognized by the fact that the first and last burn with a
luminous flame, while the flame of the second is non-luminous. By
passing the vapors into a condenser, and using the method described
above, a more or less complete separation can be made.
General Properties of Hydrocarbons. — All these substances
are extremely indifferent in their chemical behavior. They have
none of the properties of acids, bases, or salts. The halogens, notably
chlorine and bromine, however, interact with them (see below). When
burned they all produce carbon dioxide and water. Wrhen their
vapors are passed through a white-hot tube they suffer decomposition
into a mixture of hydrogen and hydrocarbons of smaller or larger
(see Benzene) molecular weight.
Methane. — Methane, otherwise known as marsh-gas, is the
chief component of natural gas. It rises to the Surface when the
bottoms of marshy pools are disturbed, and issues from seams in coal
beds. In these two cases it results from the decomposition of vege-
table matter in absence of air. When methane enters mines from
a coal seam it is called " fire-damp " (Ger. Dampf, vapor), on ac-
count of the explosive nature of the mixture it forms with the air.
The carbon dioxide formed by the explosion is called by the miners
" choke-damp."
Methane may be made from inorganic materials by the action of
494 INORGANIC CHEMISTRY
water upon aluminium carbide, prepared by the interaction of alumi-
nium oxide and carbon in the electric furnace (cf. p. 478) :
A14C3 + 12H20 -> 4A1(OH)3 + 3CH4.
In the laboratory the gas is commonly obtained by the distillation
of a dry mixture of sodium acetate and sodium hydroxide :
NaC2H302 + NaOH -» Na,C08 + CH4.
When a mixture of methane and chlorine is exposed to sunlight
several changes occur in succession (cf. pp. 176, 214) :
CH4 + CL, -> CH3C1 + HC1,
CH3C1 + C12 -» CH2C12 + HC1,
CHjCl, + 01, -» CHC13 + HC1,
CHC13 + Cij-> CC14 +HC1.
This kind of interaction with the halogens is characteristic of com-
pounds of hydrogen and carbon. It takes place slowly, and is there-
fore entirely different from ionic chemical change. It consists in a
progressive substitution of chlorine for hydrogen, unit by unit. The
various groups which, in the first three of these products, are asso-
ciated with chlorine, occur in many organic compounds, and receive
the names methyl (CH3— ), methylene (CH2=), andmethenyl (CH = ).
The compounds are known, therefore, as methyl chloride, methylene
chloride, methenyl chloride (chloroform), and carbon tetrachloride.
The last two are volatile liquids, chloroform being the only one of the
four which is a familiar substance. The corresponding iodine deriva-
tive iodoform CHI3 is a common antiseptic. These substances are not
salts, and are not ionized in solution. They are very slowly hydro-
lyzed by water, — carbon tetrachloride, for example, giving carbonic acid
and hydrochloric acid. Although carbon is a non-metal (cf. p. 405),
this action requires a high temperature.
Organic Radicals. — In carbon chemistry there are groups of
units which pass unaltered from compound to compound and receive
the name organic radicals. They usually lack a property which inor-
ganic radicals generally possess, namely, the power to form ions (p. 312).
Methyl is such a radical, being found in methane CHS.H, methyl chloride
CH3.C1, methyl alcohol CH3.OH, and acetic acid CH3.C02H. Similarly
we have ethyl C2H6 in ethane C2H5.H and in ethyl alcohol C
SOME CARBON COMPOUNDS 495
Methyl, ethyl, and propyl (C3H7 — ) are univalent radicals. We have
also ethylene C2H4 = , propylene C3H6 = , and so forth, which are biva-
lent. Groups like NO.,1 (p. 441), NH2X (p. 487), CHgCO1, and many
more, are other non-ionized radicals found in organic compounds (see
Acetic acid, below).
Ethylene. — In addition to the paraffin series there are several
other homologous series (p. 491) of hydrocarbons. Ethylene C2H4 is
the first member of the second series. It corresponds to ethane, but
contains in each molecule two hydrogen units less than does this
substance. The general formula for this series is CBH2B. As we shall
see, ethylene and the members of the ethylene series are thus all un-
saturated, possessing two free valences.
Ethylene is most easily made by heating common alcohol, which
is ethyl alcohol, with concentrated sulphuric acid :
C2H5OH -> H.,0 + C2H4.
The action really takes place in two distinct stages, and the interme-
diate product can be isolated. First, ethyl sulphate (cf. p. 441) is
formed, C2H5OH + H2S04 & C2H5HS04 + 1^0. Above 150°, how-
ever, this substance, which is a thick syrup, is dissociated, giving
ethylene and sulphuric acid, C2H5HS04 — > C2H4 + IL^04. A compari-
son of the structural formulae of the alcohol and ethylene shows that
this loss of water must leave the carbon partly unsaturated :
H H H H H H
II II II
H-C-C-O-H H-C-C-H or H-C = C-H
II -II
H H
The water may also be removed by allowing alcohol to fall drop by
drop on heated phosphoric anhydride. The solid phosphoric acid re-
mains behind and ethylene escapes.
Ethylene is formed along with acetylene and other substances,
when any saturated hydrocarbon is heated strongly. Even methane
gives it:
2CH4 -> C2H4 + 2H2.
Ethylene is a gas, which, when liquefied, boils at —105°. Its critical
temperature is 130°. At 0° it may be liquefied by a pressure of 42
atmospheres. It burns in the air with a flame which, on account of
496 INORGANIC CHEMISTRY
the great separation of free carbon which takes place temporarily dur-
ing the combustion (cf. Flame), is highly luminous. It will be seen
that in the formula but three of the valences of each carbon unit are
occupied. As carbon is either bivalent or quadrivalent, we should
expect that in this compound the combining capacity of the carbon
would not be completely satisfied. We find this to be the case. Ethy-
lene is easily reduced by nascent hydrogen (p. 423) to ethane, taking
up two xinits of hydrogen in the process. When ethylene is passed
through liquid bromine it is rapidly absorbed, and the bromine seems to
increase in volume and finally loses all its color, leaving a transparent
liquid having the composition C2H4Br2, ethylene bromide. The second
of the above graphic formulae for ethylene is the one generally used.
In spite of appearances, it is not intended to indicate that the two
units of carbon are more forcibly held together than in other com-
pounds (cf. p. 106). It simply chronicles the fact that one valence
of each carbon unit is unoccupied.
Acetylene. — This substance, likewise a gas, is the first member of
still another unsaturated, homologous series. Its formula C2H2 shows
that its molecule lacks four of the hydrogen units necessary to the com-
plete saturation which we find in ethane. Graphically its structure is
usually represented thus : H— C = C — H. This gas is formed in
small quantities by direct union of carbon and hydrogen in the electric
arc (p. 478). It is also produced when ethylene is passed through
a heated tube : C2H4— >C2H2 + H2 (cf. Flame). When calcium carbide
(p. 478) is thrown into water, violent effervescence occurs, the calcium
carbide is disintegrated, a precipitate of calcium hydroxide is formed,
and acetylene passes off as a gas :
CaC2 + 211,0 -» Ca(OH)2 + C2H2.
This action is like that of water on calcium phosphide (p. 461), calcium
sulphide (p. 376), and magnesium nitride (p. 417).
Acetylene burns with a flame which is still more luminous than
that of ethylene. Its most characteristic property is that when passed
through an ammoniacal solution of a cuprous salt, it yields a red pre-
cipitate of a carbide of copper known as copper acetylene. The
equation : Cu2(OH)2 -f C2H2 — > CugC,; + 2H20, partially represents the
change. This red precipitate, when dried, is extremely explosive, on
account of the great amount of energy set free when it breaks up into
SOME CARBON COMPOUNDS 497
its constituents. Its formation is used as a test for acetylene in mix-
tures of gases.
Acetylene may be handled safely as a gas at the ordinary pressure,
but when contained in cylinders at more than two atmospheres pressure
it is readily exploded by any shock. This is due to the fact that it is an
endothermal compound :
C2H2 -> 20 + H2 + 53,200 calories.
When used as an illuminant, it is developed in a suitable generator as
it is needed. It begins visibly to decompose at 780°.
The unsaturated nature of this substance is shown by the avidity
with which it unites with hydrogen and the halogens, forming satu-
rated compounds.
Benzene. — Limits of space forbid the discussion of any of the other
series of hydrocarbons. One of the most important has not been men-
tioned, however. It? is that of which the first member is benzene, C6H6.
More than half of the known compounds of carbon are derived from
this substance. Phenol (cf. p. 441) C6H5OH is the fundamental alcohol
of this set. Benzene is obtained from the products of the dry distilla-
tion of coal (cf. Coal gas), being formed, probably, from the acetylene
which the decomposition of other hydrocarbons yields. At all events,
when acetylene is passed through a heated tube some benzene is pro-
duced, 3C2H2 — » C6H6, along with free carbon and hydrogen.
THE ACIDS OF CARBON.
Formic Acid. — The removal of water from formic acid produces
carbon monoxide (p. 485). Although we cannot reverse the process
and cause carbon monoxide to combine with water, we can make it
unite with bases. By passing carbon monoxide over hot sodium
hydroxide, we obtain sodium formate, from which formic acid may be
liberated by double decomposition with another acid :
CO + NaOH -» CHO(ONa).
This acid is secreted by red ants, and is found in stinging nettles. It
is a liquid boiling at 100.1° and freezing at 8.6°. Although one of the
weaker acids, it is much more active than acetic acid. The molecule
contains two atoms of hydrogen, but the acid is, in fact, monobasic.
498 INORGANIC CHEMISTRY
The structural formula of the acid must take account of this fact.
Three possibilities present themselves :
0-H 0 H 0
Cx || ^C/ \
^0-H H-C-O-H H/ ^ O
In the first and last the hydrogen units should behave alike. The
second formula is the only one which expresses the replaceability of
one unit and not of the other by a metal. Since the hydrogen in
methane is not replaceable by metals (p. 493), we infer that the unit
directly combined with carbon is the non-replaceable one. Sodium
formate is therefore
0
II
H - C - 0-Na.
Acetic Acid. — This acid is produced in the dry distillation of
wood* (p. 476). Large quantities of it are manufactured from dilute
alcohol. The liquid is allowed to flow in a slow stream through a
barrel filled with shavings. Holes in the barrel provide for the access
of air, and a bacterium with which the shavings are infected promotes
that oxidation of the alcohol in which the change essentially consists :
C2H5OH + 02 -> C2H30(OH) + H,O.
Oxygen alone does not affect alcohol in the cold. The bacterium (B.
aceti) " mother-of- vinegar " ) assists this action, as lower organisms are
found to assist many chemical actions, in a way which is not as yet
thoroughly understood, and which may be described roughly as
catalytic (see Fermentation, p. 501).
The dilute solution of acetic acid produced in this manner contains
from five to thirteen per cent of acetic acid, and is known as vinegar.
By fractional distillation the solution may be concentrated until a little
water only remains, and finally, by freezing (cf. p. 294), the acetic acid
may be crystallized out. Pure acetic acid, in consequence of its
freezing readily in cold weather, is known as " glacial" acetic acid. It
melts at 16.7° and boils at 118°.
Although four atoms of hydrogen are contained in its molecxile,
but one of these is replaceable by metals. This fact is recognized in
the constitutional formula (p. 391) of the acid, CH3CO(OH). In this
* The dry distillation of bones (p. 455), on the other hand, and of animal mat-
ter (p. 417) in general, gives alkaline liquids, because of the ammonia that is formed.
SOME CARBON COMPOUNDS 499
acid a radical, methyl (CH3— ), takes the place of H in formic acid.
The other organic acids similarly are related to formic acid and con-
tain other organic radicals (see Palmitic acid). Thus the group
0
II
— C — 0 — H, called carboxyl, is contained in all carbon acids, and in
each of them, as in formic acid, bears the replaceable hydrogen unit.
The other three hydrogen units in acetic acid, however, are replaceable
by chlorine, as is the case with the hydrogen units in hydrocarbons.
The above brief statements in regard to the mode of expressing the chemical
properties of a substance by an elaborated formula bring out a tendency which
prevails in the behavior of organic substances and is almost entirely lacking in
inorganic chemistry. The units may be removed from the molecule of an organic
substance one by one, and other units or groups may be substituted for them with-
out disturbing the rest of the molecule. The changes take place, not as in the case
of ionized substances, by the splitting of the molecule into two or more groups which
act as wholes, but by the displacement of the units piecemeal and the introduction
of new properties according to the nature of the groups introduced. Thus, if by any
O
II
means we replace an atom of hydrogen by the organic radical— C — O — H the prod-
uct is an acid. If we replace it simply by the group OH the product is an alcohol.
Each substitution may take place repeatedly in a given molecule, so that di-basic
or tri-basic acids, di-hydric or tri-hydiic alcohols (see Glycerine), or substances
which contain both OH and — COOH in the same molecule (like lactic acid and tar-
taric acid), are formed. Other groups which may be introduced or removed are
— NH2, —NO,, — CN, etc., each of which confers upon a substance the properties
which go with the group, irrespective of the other features which the structure of
the substance may already present.
Oxalic Acid. — This acid has the composition H2C204, and is di-
basic. Its calcium salt is the least soluble of the salts of calcium, and
is found in many plants in the form of bundles of needle-shaped crys-
tals. Potassium hydrogen oxalate is found in the juices of various
species of oxalis. The acid may be made by oxidation of sugar with
nitric acid.
Oxalic acid is commonly used in the form of the white crystalline
hydrate, H2C204,2H20. When heated carefully it sublimes unchanged.
Stronger heating decomposes it into carbon dioxide and formic acid,
and the latter breaks up, in part, into water and carbon monoxide. In
the presence of dehydrating agents like sulphuric acid, water and the
two oxides of carbon alone are formed (p. 485).
500 INORGANIC CHEMISTRY
CARBOHYDRATES AND FERMENTATION.
Carbohydrates, — The various kinds of sugar, starch, and cellu-
lose form a closely related group of substances. As it happens that
the proportion of hydrogen to oxygen in the composition of most of
them is the same as that of these elements in water, they are known
by the name of carbohydrates. None of these substances show any
distinct evidence of ionization.
Dextrose, otherwise known as glucose or grape-sugar, is a white
crystalline substance having the composition C6H1,06. It is found
dissolved in the juices of sweet fruits, such as grapes.
The most familiar sugar is cane-sugar C12H22On, which may form
as much as 18 per cent by weight of the juices pressed from the
sugar-cane, and sometimes reaches 15 per cent of the fluid material in
the sugar-beet. It is prepared by boiling the juices with animal char-
coal (p. 476), to remove the coloring matter which would otherwise
give the sugar a brown tint. The liquid is then concentrated until
crystals appear. The mother-liquor which no longer deposits crystals
is known as molasses.
When a solution of cane-sugar is boiled with water containing a
small amount of any acid, a slow hydrolysis of the sugar takes place,
whereby two other sugars, namely, dextrose and levulose, are produced :
C12H22011 + H2O-> C6H1206 -f C6H1206. The action of the acid is cata-
lytic, and the rate of this more or less leisurely chemical change
depends upon the concentration of the hydrogen ions. It therefore
furnishes one means of comparing acids as regards their chemical
activity, and has the special advantage that the acid is not consumed
during the process (cf. p. 356), but remains of constant concentration
throughout the whole time. This process, by which cane-sugar is
decomposed, is spoken of as inversion, and the mixed product is called
invert-sugar. The change can be produced, not alone by acids, but
also by certain complex chemical compounds secreted by yeast (see
below).
The relation of starch to the sugars is seen, not only in the formula
(C6H1006)n, but in the fact that by boiling starch* with dilute acids,
dextrose is formed along with other products of the hydrolysis. Com-
mercial "glucose" is made by this process. Starch is an insoluble
white substance which is found in the form of fine particles in the
fruit and other parts of plants.
Cellulose has the same composition as starch. It forms the frame-
SOME CARBON COMPOUNDS 501
work of the cells of plants. In many cases it is overlaid with a con-
siderable thickness of liguin, which in paper-making is removed by
boiling the wood with sodium hydroxide or calcium bisulphite solu-
tion (p. 476). When the product has been washed thoroughly with
water, almost pure cellulose remains. Matted cellulose in thin sheets
forms the basis of paper, and filter paper contains nothing else (see
under Aluminium sulphate). Other forms of pure cellulose are known
as cotton, linen, and jute, according to their sources. Although the
chemical composition of these varieties is identical, the physical prop-
erties vary considerably.
Fermentation. — This is the name given to a number of different
chemical changes, brought about by catalytic action of complex chemi-
cal compounds secreted by living organisms. These compounds are
called enzymes, and, in many cases, have been separated from the
organisms by means of solvents. Their action must be regarded as
catalytic, since small quantities of the active organisms or of the
enzymes can produce very extensive chemical changes without them-
selves suffering alteration in the process.
The organisms may be divided into three classes, each secreting
different enzymes which confine themselves for the most part to special
kinds of chemical change. (1) The molds, when grown in sugar solu-
tion or beef extract, or other nutritive solutions, produce decompositions
known collectively as putrefaction. (2) Certain bacteria promote the
oxidation of alcohol to acetic acid (p. 498). Some also decompose
sugar, furnishing butyric or lactic acid as one of the products. (3)
The yeasts (saccharomycetes) flourish in solutions of some sugars, and
decompose them into alcohol and carbon dioxide. This decomposition
is known as alcoholic fermentation. These changes are usually brought
about by actual introduction of the organism. In brewing, however,
the enzyme itself, diastase, is employed to hydrolyze starch.
The juice of grapes when set aside at a suitable temperature soon
begins to ferment, owing to the propagation in it of a yeast (S. ellipsoi-
deus} which is found upon the skins of the grapes, and decomposes the
sugar. While small quantities of a number of different compounds are
formed, by far the greater part of the sugar is resolved quantitatively
into alcohol and carbon dioxide : C6H1206 -+ 2C2HBOH + 2C02. The
liquid effervesces, and the carbon dioxide escapes into the air. The
wine is allowed to stand, after fermentation, until it has deposited a
considerable crust of material known as argol, which consists mainly of
502 INORGANIC CHEMISTRY
potassium-hydrogen tartrate (KHC4H406, cream of tartar). The con-
centration of the sugar in the grape-juice being small, the quantity of
alcohol contained in the product is not very great. By distillation of
wine, a liquid containing a much larger proportion of alcohol is made,
and is known as brandy. The special flavors of wines and brandies
depend upon materials, other than sugar, originally contained in the
fermented liquid, upon by-products of the fermentation, and upon
materials which arise by slow chemical changes while the liquor is
stored.
The preparation of beer involves a preliminary step, for the sugar
needed for the fermentation is made from the starch contained in various
kinds of grain, particularly barley. The conversion of starch into
sugar is effected by the use of diastase, which is formed in considerable
amounts in sprouting barley.* When the quantity of this unorganized
ferment has reached a maximum, the barley is dried and crushed. The
product, known as malt, is mixed in water with other grain which has
not gone through this process, and the whole is slowly heated. During
the heating, the diastase hydrolyzes the molecules of the starch, giving
a sugar called maltose (C^H^On). The solution, after nitration and
boiling, is cooled. It is then placed in the fermentation vats, and com-
mon yeast (S. cerevisice) is added. During the growth of this plant the
maltose is decomposed, each molecule producing two molecules of
dextrose, and the latter is broken up into alcohol and carbon dioxide
according to the equation already given. Aside from the alcohol and
carbon dioxide, considerable quantities of other substances extracted
from the grain remain in the solution and form the so-called " extract.''
which varies in kind and quantity in different varieties of beer.
Whiskey is prepared in a somewhat similar fashion, although other
sources of starch, such as rye and corn, may be used. The fermented
liquid contains but a small proportion of alcohol, and is distilled
(rectified}. The alcohol, being more volatile than water, tends to pass
over first, and a product containing any desired proportion of it (see,
however, below) can be made. Alcohol for technical or chemical use is
made in the same way, potatoes sometimes being used as a source of
the starch.
* In the digestion of bread, potatoes, and other food containing starch, the
same office is performed by the saliva. Starch, being insoluble, could not be
absorbed through the walls of the alimentary tract, while sugar is soluble and can be
so absorbed. In a roughly similar way the albuminous parts of food are rendered
soluble and capable of assimilation by the pepsin of the stomach, and the pan-
creatin of the large intestine, which convert them into peptones.
SOME CARBON COMPOUNDS 503
ALCOHOLS, ESTERS, AND ETHERS.
Alcohols. — We have already seen that when wood is distilled,
methyl alcohol is found in the fluid product. When purified this is a
colorless liquid boiling at 66°. Its solution in water shows no evidence
of ionization, although it would probably be safer to say that the
ionization is so slight as to be imperceptible, than to say that the com-
pound is npt ionized at all. The formula (CH3OH) makes it impossible
to represent the structure of the substance in more than one way :
•
H
I
H-C-O-H
I
H
All alcohols contain the group =C— 0— H (cf. p. 499).
Common alcohol, ethyl alcohol C2H5OH, is formed in the fermenta-
tion of solutions of sugar by yeast (p. 501), and is separated from the
water and the other products of fermentation by distillation. The
product contains 95 per cent of alcohol and 5 per cent of water, and
is applicable to most commercial uses. Absolute alcohol, entirely free
from water, cannot be made by distillation alone (see below). The
95 per cent spirit is placed in vessels filled with quicklime, the latter
interacts with water producing calcium hydroxide, and the clear liquid
which is poured off is distilled once more. Pure alcohol boils at 78.3°.
Mixtures of two liquids, when distilled, behave in one of three ways.
Two of these have been described already (p. 183), and alcohol (b.-p.
78.30°) and water (b.-p. 100°) illustrate the third. In this case the
vapor tension of a certain mixture is higher than that of any other
mixture and higher than that of either component separately. This
special mixture has, therefore, a lower boiling-point than any other.
In the present instance this mixture contains 95.57 per cent of alcohol
and 4.43 per cent of water and boils at 78.15°. When the fermented
liquid, with its large percentage of water, is distilled, the alcohol all
tends to pass off first, in association with that part of the water required
to constitute the mixture of minimum boiling-point. Repeated distil-
lation simply eliminates more completely the excess of water beyond
this amount (viz., 4.43 per cent), by leaving it in the residues.
Glycerine (p. 452) is an alcohol containing three hydroxyl groups,
a trihydric alcohol,
504 INORGANIC CHEMISTRY
Esters. — When an organic acid and an alcohol are mixed, a very
slow chemical action takes place, which, being reversible, in no case
reaches completion. With the simplest members of these groups, for-
mic acid and methyl alcohol, for example, the change is :
HCOOH + HOCH8 +± HCOOCH3 + H2O.
The product is known as methyl formate. The corresponding action
between acetic acid and ethyl alcohol : CH3COOH + HOC2H5 +± CH3CO
OC2H5 + H20, results in the formation of ethyl acetate. In this case,
when equivalent quantities of the initial substances have been used
without any solvent, and a condition of equilibrium has been reached,
two-thirds of the material is found to have been transformed into
ethyl acetate and water. If we start with the latter materials in pure
form, the same equilibrium point is reached, and one-third of the mate-
rial is converted into acetic acid and alcohol.
This action is a general one, and occurs between all alcohols and
acids. The products are sometimes known as ethereal salts, because
they result from the displacement of the hydrogen of an acid by a radi-
cal. This designation, however, is not very happy, since the products
are not ionized and possess none of the properties of salts. The
special name esters, therefore, has been given to them. The action is
always extremely slow and never complete, but it may be hastened and
carried to completion by the introduction of some substance capable of
absorbing the water and so preventing the reversal. Concentrated sul-
phuric acid, for example, or anhydrous cupric sulphate, may be used.
Inorganic acids also interact with alcohol, giving esters. Thus,
nitroglycerine (p. 452) is an ester, and should be called glyceryl trini-
trate. The use of sulphuric acid to assist in the removal of the water
is illustrated in the preparation of this substance. Gun-cotton (p. 441)
is an ester of nitric acid also, for cellulose is a complex alcohol. Ethyl
sulphate (p. 495) is an ester of sulphuric acid. In this case the action
may be made complete by using sulphuric acid containing an amount
of sulphur trioxide sufficient to combine with the water to be pro-
duced.
The above actions, in which an ester, like ethyl acetate, is formed,
may be almost completely reversed if a sufficient amount of water is
added (cf. p. 250). The hydrolysis of the ester is hastened by the
presence of free acids in the water. This is owing to the catalytic
action of the hydrogen ions, and the acceleration is proportional to the
activity of the acid used. The acid, however, although it hastens the
SOME CARBON COMPOUNDS 505
action, does not carry it beyond the condition of equilibrium which it
would eventually have reached with the same amount of water alone
(cf. p. 237).
When esters are boiled with strong bases, such as sodium hydroxide
solution, the salt of the acid and an alcohol are formed :
CH8COOC2H5 + NaOH -> CH3COONa + HOC2H5.
With more complex esters the sodium salts of the acids thus produced
are known as soaps, and this general kind of action is called, therefore,
saponification (Lat. sapo, soap). The speed with which it proceeds
may be used as a means of measuring the activity of bases.
Soap. — Soap is prepared by the decomposition of fat. The latter
substance is a mixture of several rather complex esters. In beef fat
the chief esters present are tripalmitin, tristearin, and triolein.* It
will be sufficient to illustrate the chemistry of the change by discussing
the case of one of these substances. Tripalmitin is the glyceryl ester of
palmitic acid. When fat is mixed with hot sodium hydroxide solution
it first forms an emulsion f in which the fat is disseminated in minute
droplets through the liquid. This is a result of surface tension.
When the emulsion is boiled, the fat is slowly decomposed into sodium
palmitate and glycerine. The change is precisely similar in plan to
the simpler one jiist discussed :
C15H31CpO-C-H2 HOCH2
I I
C15H31COO - C - H -f- 3NaOH -» SC^l^CC-ONa + HOCH
I I
C15H31COO - C - H2 HOCH2
Changes similar to this occur with the other two substances. The
only difference is that the organic radical in the case of tristearin is
C17H35, and in the case of triolein C^H^. Both products in each case are
soluble in water, but when common salt is added to the solution the
sodium salts of the organic acids are separated (" salted out," see
Chap, xxxv) as a solid mass, which is known as soap. When potas-
sium hydroxide takes the place of sodium hydroxide the mass is semi-
fluid, and is known as soft soap.
* Butter fat contains in addition to the above a certain amount of tributyrin,
in which the organic radical is C3H7. Olive oil consists mainly of tripalmitin and
triolein.
t In the intestines the same office is performed by the gall, secreted by the liver,
and so the fat is prepared for absorption into the system.
506 INORGANIC CHEMISTRY
These complex esters, like the simpler ones, are decomposed also
by water without the aid of a base, although much more slowly. By
the use of superheated steam, however, rapid hydrolysis can be pro-
duced, and the products are the free organic acids and glycerine. The
mixture of acids set free from fat by the action of steam is a solid,
waxy mass, known as " stearin," and used in the manufacture of
candles. The oleic acid, which is a liquid, is pressed out.
These acids, not being soluble in water, have no effect upon litmus ;
but the fact that they are acids may be recognized when it is found
that they are converted into soluble salts by bases, such as sodium
hydroxide :
+ NaOH <=?. H20 -1- C17H36COONa.
The cleansing power of soap solution seems to depend on the surface
tension of the liquid rather than on any chemical action.
The components of soap, like other salts, are highly ionized in solu-
tion, and show all the properties of ionogens. When soap is dissolved
in hard water (cf. p. 113), a white, flocculent precipitate is formed,which
coagulates upon the sides of the vessel. This is a mixture of the cal-
cium salts formed by union of the proper ions. For example, the
sodium palmitate is changed as follows :
2C15H31COONa + CaS04 -» (C16H31COO)2Ca|+ ^SO,.
Most of the salts of these acids, with the exception of those of potas-
sium and sodium, are insoluble in water.
Drying Oils. — The oils commonly used as " dryers " for mixing
with varnish and paint and in making linoleum, such as linseed oil,
hemp oil, poppy oil, and nut oil, contain esters of acids with unsaturated
radicals. One of the constituents, for example, is the glyceryl ester
of linoleic acid. The formula of this acid is C17H31COOH. It con-
tains four hydrogen atoms less than the corresponding saturated acid
(stearic acid). These oils, especially after having been recently
heated, alone or with catalytic agents like lead oxide and manganese
dioxide, absorb oxygen rapidly from the air, and become solid. They
do not dry, in the ordinary sense, by evaporation.
Ether. — When two molecules of an alcohol lose one molecule of
water, an ether is produced :
2CH,OH -rt (CH8),O + H20,
SOME CARBON COMPOUNDS 507
Thus, methyl alcohol gives methyl ether, and ethyl alcohol, ethyl or
common ether. The action is most easily carried out by two steps.
In making common ether, ethyl alcohol acts upon sulphuric acid, giv-
ing ethyl sulphate (p. 495) ; and the latter, when warmed gently with
excess of alcohol, gives ethyl ether :
C2H5HS04 + C2H5OH -» (C2H5)20 1 + H2S04.
The ether escapes as vapor and is condensed.
Ethyl ether is a volatile liquid boiling at 34.6°. It is largely used
as a solvent for iodine, fats, and other substances not readily soluble
in water, and as an anaesthetic.
CYANOGEN.
Cyanogen. — This compound is formed in small amount when a
discharge of electricity takes place between carbon poles in an atmos-
phere of nitrogen (cf. p. 416). Cyanogen, being an endothermal sub-
stance, is more easily made as one product in an exothermal action
(p. 301). It is prepared by allowing a solution of cupric sulphate to
trickle into a warm solution of potassium cyanide. The cupric cyanide,
at first precipitated, quickly decomposes, giving cuprous cyanide and
cyanogen :
2KNC + CuSO4 -> Cu(NC)2 J. + K2S04,
2Cu(NC)2 -* 2CuNC + C2N2 f .
Cyanogen is a very poisonous gas with a characteristic, faint odor.
Hydrocyanic Acid. — This acid, called also prussic acid, has the
formula H — N = C, #nd is most easily made by the action of an acid
upon a cyanide (see Potassium cyanide) followed by distillation. It
is a colorless liquid boiling at 26.5°. It has an odor like that of bitter
almonds, and is highly poisonous. In aqueous solution it is an ex-
tremely feeble acid, and is hardly ionized at all. In consequence of
this, potassium cyanide is markedly hydrolyzed by water, and its aque-
ous solution is strongly alkaline. The behavior of hydrocyanic acid
shows it to be an unsaturated body, a fact which is taken account of
in the above formula, and illustrated in the two following paragraphs.
Cyanates. — When potassium cyanide is fused and stirred with
an easily reducible oxide, like lead oxide (PbO), the metal (for example,
508 INORGANIC CHEMISTRY
the lead) collects at the bottom of the iron crucible in molten form, and
potassium cyanate is produced ;
KNC + PbO -> KNCO + Pb.
Cyanic acid is itself very unstable. Ammonium cyanate is chiefly
remarkable for its transformation into urea (p. 488).
Thiocyanates. — When potassium cyanide in aqueous solution is
boiled with sulphur or with a polysulphide (p. 376), it is converted into
potassium thiocyanate, KCNS. This salt is used in testing for ferric
ions on account of the deep-red color of ferric thiocyanate (cf. p. 250).
The ammonium salt undergoes at 170° a transformation parallel to
that of ammonium cyanate, thiocarbamide (sulpho-urea) being formed.
Exercises. — 1. Make the graphic formulae of hexane (p. 491),
methyl acetate (p. 499), ethyl formate, ethylene bromide (p. 496),
oxalic acid (p. 499), ethyl ether (p. 507).
2. Make equations for the formation of aluminium carbide
(p. 478), the hydrolysis of starch to maltose (p. 502), the saponifi-
cation of triolein (p. 505).
3. Name the radicals : C5HU, C5H10, C5H9, C16H13, and the sub-
stances C6HUC1, C5HUOH, (C5Hn)20, C6HUHS04.
4. Is the hydrocarbon radical of oleic acid (p. 505) saturated,
or not ?
5. Prepare a summary of the various statements that have been
made in the text about catalysis (e.g. pp. 75, 109, 111, 170, 175, 189,
237, 380, 396, 399, 423, 484), and illustrate fully.
CHAPTEE XXX
FLAME
«
Meaning of the Term. — In the combustion of charcoal there is
hardly any flame, for the light emanates almost entirely from the incan-
descent, massive solid. When two gases are mixed and set on tire, a sort
of flame passes through the mixture, but this can hardly be accounted a
flame, in the ordinary sense, either. The rapid movement of the flash,
and the explosion which accompanies it, are in a manner the precise
opposite of the quiet combustion which is characteristic of flames.
. With illuminating-gas the production of its very characteristic flame
is due to the chemical union of a stream of one kind of gas in an
atmosphere of another. The flame is made up of the heated matter
where the two gases meet. In the case of a burning candle, one of the
bodies appears to be a solid, but a closer scrutiny of the phenomenon
shows that the solid does not burn directly. A combustible gas
is manufactured continuously by the heat of the combustion and
rises from the wick. The introduction of a narrow tube into the
interior of the flame enables us to draw off a stream of this gas and to
ignite it at a remote point. Thus, a flame is a phenomenon produced
at the surface where two gases meet and undergo combination with
the evolution of heat and, sometimes, light.
In the chemical point of view, it is a matter of indifference whether
the gas outside the flame contains oxygen, and the gas inside consists
of substances ordinarily known as combustibles, or whether this order
is reversed. In an atmosphere of ordinary illuminating-gas, the flame
must be fed with air. This condition is easily realized (Fig. 93).
The lamp-chimney is closed at the top until it has become filled with
illuminating-gas. After the lapse of a few minutes this can be ignited
as it issues from the bottom of the wide, straight tube which projects
from the interior. When the hole in the cover of the lamp-chimney is
then opened, the upward draft causes the flame of the burning gas to
recede up the tube, and there results a flame fed by air and burning in
coal-gas. In an atmosphere of this kind, materials playing the part of
a candle burning in air would have to be substances which, under the
609
510
INORGANIC CHEMISTRY
influence of the heat of combustion, give off oxygen. Strongly heated
potassium chlorate, for example, appears to burn continuously in such
an atmosphere when lowered into it in a deflagrating spoon.
Luminous Flames. — The flame of hydrogen, under ordinary cir-
cumstances, is almost invisible, nearly all the energy of the combustion
being devoted to the production of heat. A part of this, however,
may be transformed into light by the sus-
pension of a suitable solid body, such as a
platinum wire, in the flame. The holding of
a piece of quicklime in an oxy hydrogen flame
(cf. p. 109) is a practical illustration of this
method of securing luminosity. In general,
luminosity may be produced by the presence
of some incandescent solid.
In the Welsbach lamp the flame itself is
non-luminous, and, but for the mantle, would
be identical with the ordinary Bunsen flame.
The mantle which hangs in the flame, how-
ever, by its incandescence, furnishes the light.
This mantle is composed of a mixture of 99
per cent thorium dioxide (ThO,) and one per
cent cerium dioxide (Ce02). While many
oxides would give out a white light and could
be obtained much more cheaply than these,
they have not sufficient coherence to make
their use practicable. It is worth noting that
any appreciable variation from the above pro-
portions, by the introduction of either more
or less cerium oxide, produces a marked
diminution in the intensity and whiteness of
the light (see, also, under Thorium).
In cases of brilliant combustion, as of magnesium ribbon or phos-
phorus, a solid body is formed whose incandescence accounts for the
light. The flame of ordinary illuminating-gas does not at first sight
appear to give evidence of the presence of any solid body. But if a
cold evaporating dish is held in the flame for a moment, a thick deposit
of finely divided carbon (soot) is formed, and we at once realize that
the light is due to the glow of these particles in a mass of intensely
hot gas. Carbon is, indeed, an extremely combustible substance, and
FIG. 93.
FLAME 511
is eventually entirely consumed. But a fresh supply is continually
being generated in the interior of the flame, while the oxygen with
which it is to unite is outside the flame altogether. Thus the carbon
particles persist until, drifting with the spreading gas, they reach the
periphery of the flame.
It cannot be said that no flames are luminous unless a solid body is
contained in them. When compressed hydrogen is burned in an atmos-
phere of oxygen under pressure, the light given out by the flame is
much greater, and, in general, illuminating power seems to be height-
ened by increase in the density of the gas. In special cases, also, such
as the explosion of a mixture of nitric oxide with carbon bisulphide
vapor, a flame which has considerable illuminating power is produced,
although the density of the gases is low and solids are lacking.
The Bunsen Flame. — In the burner devised by Robert Bunsen, a
jet of ordinary illuminating-gas is projected from a narrow opening
into a wider tube. In this tube it becomes mixed with air, drawn in
through openings whose dimensions can be altered by means of a per-
forated ring. When the supply of air is sufficient, the flame becomes
non-luminous. With a somewhat different construction, and the use of
a bellows to force a larger proportion of air into the gas, a still hotter
flame can be produced. The instrument in this case is known as a
blast-lamp. The high temperature of the Bunsen flame is not difficult
to account for. It will be seen at once, on handling the burner, that
the flame diminishes to one-half or one-third of its previous size when
air is admitted. Since the same amount of gas is burning in both
cases, and the products in both cases are ultimately alike, the total
amounts of heat produced must in both cases be the same. This state-
ment may safely be made, since experiment shows that when chemical
substances pass from one condition to another, the amount of heat
evolved is the same whatever intermediate steps may be taken or
omitted (cf. p. 78). The production of an equal amount of heat in a
smaller flame necessarily causes this to have a higher average temper-
ature. It will be noted that it does not follow from this that the tem-
perature is higher in every part of the non-luminous flame than in the
corresponding locality in the luminous flame. We shall see later, in-
deed, that this is not the case.
It is instructive to note the effect of forcing in larger and larger
proportions of air into the Bunsen flame. The flame at the top of the
tube continually diminishes in size, even after it has become non-
512
INOKGANIC CHEMISTRY
luminous. Finally, a point is reached at which the flame is so unstable
that the smallest further increase in the supply of air causes it to
descend into the tube. The mixture of illuminating-gas and air in the
tube of a Bunsen burner is an explosive one, and the explosion-flame
will proceed through it with greater rapidity, the more nearly the
quantity of air approaches that required for complete combustion.
When the speed with which the explosion-flame would move, equals
that with which the stream of the mixed gases is coming upward
through the tube, the flame reaches the unstable condition just men-
tioned. Any increase in the proportion of air raises the speed with
which the explosion can travel, and the flame is thus able to proceed
down the tube against the stream of gas. This phenomenon is fre-
quently noticed in the Bunsen burner, when the holes admitting the
air are too large, or a draft momentarily causes an increase in the
supply of air. The flame "strikes back," and thereafter continues
to burn at the bottom of the tube.
Structure of the Bunsen Flame. — When an ex-
ceedingly small luminous flame is examined, the various
parts of which it consists may easily be made out. In the
interior there is a dark cone which is composed of illumi-
nating-gas and air, and in it no combustion is taking place.
A match-head may be held here for some time without
being set on fire. This is therefore not properly a part of
the flame. Outside this is a vivid blue layer ( C, Fig. 94)
which is best seen in the lower part of the flame, but ex-
tends beneath the luminous sheath, and covers the dark
inner cone completely. Outside the blue flame, and cover-
ing the greater part of it, is the cone-shaped luminous
portion (B). Over all is an invisible mantle of non-lumi-
nous flame (A), which becomes visible when the light from
the luminous part is purposely obstructed. In the lumi-
nous gas-flame, therefore, there are four regions, if we
count the inner cone of gas. The difference between this
and the non-luminous Bunsen flame is that in the latter
the luminous region is omitted, and the inner, dark cone, the blue
sheath, and the outer mantle, are the only parts which can be dis-
tinguished. We shall see that in these different regions the chemical
changes taking place are different.
FIG. 94.
FLAME
513
Composition of Illuminating -Gas. — Before considering the
chemistry of the gas-flame, it is necessary to know what substances are
burning. The illuminating-gas in Europe, and in many of the smaller
cities of the United States, is usually coal-gas, while in the larger cities
of America it is almost always made from water-gas. The former is
obtained by the destructive distillation of soft coal, and is freed from
ammonia (cf. p. 418) and tar by washing and cooling, and from hydrogen
sulphide and carbon dioxide by passage through layers of slaked lime.
The water-gas, made by the action of steam upon anthracite or coke,
being composed of carbon monoxide and hydrogen (cf. p. 485), has no
illuminating power. It is therefore " carburetted," that is, mixed with
hydrocarbons, by passage through a cylindrical structure filled with
white-hot firebrick, upon which falls a small stream of high-boiling
petroleum. The relatively involatile hydrocarbons of which the oil is
composed are thus decomposed ("cracked "), and gaseous substances
of high illuminating power are produced. The following table shows
the composition of each of these kinds of gas, together with that of
oil-gas (Pintsch's), which is composed entirely of the products from
" cracking " oil :
Components.
Coal-Gas.
Water-Gas.
Oil-Gas.
Illuminants ....
5.0
16.6
45.0
Heating gases:
Methane ....
34.5
19.8
38.8
Hydrogen
Carbon monoxide
49.0
7.2
32.1
26.1
14.6
Impurities:
Nitrogen ....
Carbon dioxide .
3.2
1.1
2.4
3.0
1.1
Candle power .
17.5
25.0
65.0
These are average numbers, and considerable variations from these pro-
portions are often met with. The illuminants are unsaturated hydro-
carbons, such as ethylene, acetylene, and benzene, and the value of the
gas for illuminating purposes depends on the amount of these particular
components.
Chemical Changes Taking Place in the Bunsen Flame.—
The study of the chemical changes taking place in the Bunsen flame,
514 INORGANIC CHEMISTRY
particularly with the object of explaining (1) the luminosity of
the flame of the pure gas, and (2) the non-luminosity of that
produced by the same gas when it is mixed with air, has been the sub-
ject of many elaborate investigations. The questions are : Why is
carbon liberated in the former case, and why is it not liberated in the
latter ? Let us consider these questions in order.
1. The first suggestion of an explanation of the presence of free
carbon in the luminous flame was that the hydrocarbons contained in
the gas underwent a selective combustion. This was supposed to take
place in such a way that the hydrogen was first burned out of the
molecules by the oxygen, which, by diffusion, had penetrated farthest
into the flame. The carbon had, therefore, to remain free until all the
hydrogen was satisfied, and until the former had drifted into a region
in which a more liberal supply of oxygen was obtainable. Thus the
carbon was rendered incandescent by the heat of combustion of the
hydrogen. This hastily formed theory is founded upon ill-considered
premises, and is not supported by experiment. It is based upon the
familiar fact that hydrogen gas burns more easily than solid charcoal.
Such a comparison, however, is not applicable to the case of the com-
bustion of a gaseous hydrocarbon, for in such a substance both of the
elements are in a gaseous condition, and the inferior combustibility of
charcoal is due, not to an inferior power of uniting with oxygen, but
to the fact that, since it is a solid, the oxygen finds access to it only
with difficulty.
The question could manifestly be placed beyond dispute if it were
found possible to determine exactly the nature of the change which
passage through the inner, blue zone of the flame produced. If, for
example, gases could be extracted by means of a tube from the region
just outside this zone, and it were found that all the hydrogen had been
converted into water while the carbon bad not yet reached the stage of
carbon dioxide, this theory would be confirmed. It was found impos-
sible, however, to perform the experiment in this fashion, on account
of the disturbing influence of the currents produced by the withdrawal
of the stream of gas. An apparatus devised by Smithells, however,
furnished an ingenious means of securing the interconal gas, alone and
in any quantity. The air and combustible gas are admitted in propor-
tions which can be varied, and the mixture burns at the top of the
wider tube (Fig. 95). As the quantity of air is increased, however, the
speed with which an explosion-flame would pass through it becomes
greater, and finally the inner cone passes down and rests upon the
FLAME
515
mouth of the narrow tube through which the mixture of gases is issuing
more rapidly. A preliminary combustion takes place in the blue cone,
while the final conversion of the whole material into carbon dioxide and
water is completed in the outer mantle. This
remains at the top of the wider tube, where
alone the necessary air can be obtained. By
means of a side-tube (not shown) fused into
the wider tube, the gases which have traversed
the inner cone and undergone chemical change
in it can be extracted and subjected to exam-
ination.
Now it was found that hardly any of the
hydrogen had been burned by the blue cone,
while all of the carbon had reached the stage
of carbon monoxide. This was the case, not
only with illuminating-gas, which initially
contains much free hydrogen, but also when
the flame was fed with pure methane. The
greater part of the hydrogen of the latter
was set free, and was found uncombined in
the interconal gas. It is evident, therefore,
that when a hydrocarbon undergoes partial
combustion the carbon is first attacked and
the burning of the hydrogen is postponed
until a sufficient supply of oxygen becomes
available.
We may safely infer, therefore, that since
in the luminous flame a portion at least of
the carbon remains for a time uucombined,
the hydrogen must a fortiori have survived combustion also. The
luminous cone must contain a mixture of free hydrogen and free carbon.
The production of this mixture can be explained in only one way,
namely, by the dissociation of the hydrocarbons concerned, of which
ethylene is the most important. Now we have seen (pp. 496 and 497)
that, when heated, ethylene dissociates, giving acetylene, and acetylene
gives carbon and hydrogen. Hence this stochastic hypothesis is highly
probable. It only remains to be shown that acetylene is actually formed
as an intermediate substance. Evidence of this is obtained in two ways.
It is found that when the Bunsen flame " strikes back," and the combus-
tion of the gases is rendered incomplete by the contact of the flame with
FIG. <j5.
516
INORGANIC CHEMISTRY
the cold tube, a large amount of acetylene is formed. Again, when the
gases surrounding the flame of air burning in illuminating-gas (p. 509)
are withdrawn by means of a pump and caused to pass through an
ammoniacal solution of cuprous chloride (Fig. 96), large quantities
of copper acetylene are formed.
The conception that when hydrocarbons burn, they first undergo dis-
sociation, and then union with oxygen, is in harmony with what we
have observed also in the
case of the combustion of
hydrogen sulphide, where
the presence of free sul-
phur and free hydrogen
in the interior of the flame
was demonstrated (p.
378).
2. The influence of
the air admitted to the
Bun sen burner, in inter-
fering with this dissocia-
tion in such a way as to
destroy all luminosity, is
FIG. 06. the most difficult point to
explain. The effect is
frequently attributed to the oxygen which the air contains. This view,
however, is seriously weakened by a consideration of the undoubted
fact that oxygen is not required. Carbon dioxide and steam are equally
efficient when introduced instead of air. Even nitrogen, which cannot
possibly be suspected of furnishing any oxygen, likewise destroys the
luminosity. Lewes has shown that 0.5 volumes of oxygen in 1 volume
of coal-gas destroy the luminosity. But 2.30 volumes of nitrogen or
2.27 volumes of air accomplish the same result. Thus the efficiency
of air is not much greater than that of nitrogen, in spite of the fact
that one-fifth of the former is oxygen.
It is evident that the effect is due, in part at least, to the dilution
with cold gas. This is confirmed by the observation that a cold plati-
num dish held in a small luminous flame is similarly destructive of the
luminosity. Comparison of the temperatures of the inner sheaths of
the luminous and non-luminous flames shows that the temperature of
the latter is markedly lower. If the tube of the Bunsen burner is
heated so that the mixed gases are considerably raised in temperature
FLAME 517
before reaching the non-luminous flame, the latter becomes luminous.
It is probable, therefore, that the cold gas lowers the temperature of the
inner flame, and at the same time the dilution diminishes the speed of
dissociation (Lewes). Even if the temperature is not reduced below
that at which dissociation of the ethylene can occur, yet the dilution
and cooling together prevent that sharp dissociation at this particular
point which is necessary for the production of the great excess of free
carbon needed to furnish the light.
Exercises. — 1. In what way will calcium hydroxide remove
hydrogen sulphide from coal-gas (p. 513) ?
2. Make a section showing the shape of the flame produced by
burning hydrogen gas when the latter issues from a circular opening.
CHAPTER XXXI
SILICON AND BORON
IN respect to chemical relations there is a close resemblance between
silicon and carbon. The former element gives no monoxide, however,
and is quadrivalent in all its compounds. In chemical character it is
strictly non-metallic.
Occurrence. — Silicon, unlike carbon, is not found in the free con-
dition. In combination it is the most plentiful element after oxygen,
and constitutes more than one-quarter of the crust of the earth. The
oxide is silica or sand (Si02), and this oxide and its compounds are
components of many rocks. In the inorganic world silicon is the char-
acteristic element to almost as great an extent as is carbon in the
organic realm.
Preparation. — When finely powdered magnesium and sand are
mixed, and one part of the mass is heated, a violent action spreads
rapidly through the whole : *
2Mg + Si02 -> Si + 2MgO.
At the same time, and especially if excess of the metal is used, some
magnesium silicide Mg2Si is formed also. The mixture is treated with
a dilute acid which decomposes the magnesium oxide and the silicide,
and leaves the silicon alone undissolved. A simpler method, although
one using less common materials, is to pass the vapor of silicon tetra-
chloride over heated sodium. The mixture of common salt and silicon
which remains in the tube can be separated by washing with water.
Both of these methods give amorphous silicon in the state of fine powder.
When amorphous silicon is dissolved in molten zinc, the mass, after
solidification, is found to contain crystalline silicon. This form may
be made in one operation, by heating three parts of potassium fluosili-
cate K2SiF6 with one part of sodium and four parts of zinc in a closed
crucible. The sodium displaces the silicon and combines with the
fluorine, while the zinc acts as a solvent as before. The zinc is removed
by the action of a dilute acid, the silicon remaining unaffected.
518 •
SILICON AND BORON 519
Properties. — Amorphous silicon is a brown powder. It unites
with fluorine at the ordinary temperature, with chlorine at 430°, with
bromine at 500°, with oxygen at 400°, with sulphur at 600°, with nitro-
gen at about 1000°, and with carbon and boron at temperatures attain-
able only in the electric furnace. It is slowly oxidized by aqua regia
to silicic acid, and is dissolved by a mixture of hydrofluoric acid and
nitric acid, giving the fluoride.
Crystallized silicon consists of black needles belonging to the hex-
agonal system, and is less active than the other variety. It oxidizes
superficially at 400°, and the dioxide formed on the surface hinders
further action. With chlorine and fluorine it unites easily when
heated. Gaseous hydrogen fluoride interacts violently with it at a high
temperature, heat and light are given out, and silicon tetrafluoride and
hydrogen are formed. It is slowly attacked by hydrofluoric acid
mixed with nitric acid, but ntfb'by any others of the oxygen acids.
Both kinds of silicon act with fair speed upon boiling solutions of
potassium or sodium hydroxide (cf. p. 99), the metasilicate being
formed :
Si + 2KOH + -H20 -» K2Si03 + 2H2.
; *
Silicon seems to be more active than carbon, for, when it is heated with
fused potassium carbonate, potassium ^iljgate is formed and carbon is
liberated. Both kinds of silicon differ from carbon in being fusible at
a very high temperature. The product is crystallized silicon.
Silicon Hydride. — Silicon differs from carbon in giving only two
well-defined compounds with hydrogen. The chief one may be liberated
as a gas by the action of hydrochloric acid upon magnesium silicide :
Mg2Si + 4HC1 -> 2MgCl2 + SiH4.
The action is similar to that by which hydrogen sulphide is made.
Since the magnesium silicide always contains free magnesium, hydro-
gen is liberated at the same time. The gases may be separated by
passage through a tube surrounded by liquid air. The silicon hydride
is reduced to liquid form, while the hydrogen passes on. The mixture
with hydrogen is spontaneously inflammable. The pure gas becomes
so only when its pressure is reduced. In the air, however, it is easily
inflammable, by contact with a warm body. When heated, it decom-
poses into its constituents,
520 INORGANIC CHEMISTRY
Carbide of Silicon. — This compound is manufactured for use as
an abrasive, and is sold under the name of carborundum. A mixture
of quartz-sand, coke, and common salt is heated to about 3500° in the
electric furnace :
SiO2 + 3C -> SiC + 2CO.
The materials are piled upon an oblong base provided with vertical
walls at the ends, through which the terminals project. A long ridge
of loosely packed coke furnishes an imperfect conductor of high resist-
ance between the terminals. The above mixture is heaped on each side
of and above this, as well as below it. The coke becomes white-hot
when the current is turned on, and the heat radiating from this core
brings about a change, in the above sense, which affects the whole mass
more or less completely.
Silicon carbide when pure is composed of transparent, colorless,
hexagonal plates. Ordinarily the crystals are brown or black. It
stands next to the diamond and carbide of boron in hardness. It is
not oxidized by the air even at a white heat, being protected by the
layer of silica first formed. Carborundum is not affected by water or
acids, but is decomposed by alkalies. It is used in making machinery
for polishing hard rock, such as granite, and is employed also for pro-
tecting the walls of puddling ^furnaces (y.v.), and in other ways in the
steel industry.
Silicon Tetrachloride. — This compound is made by direct union
of the free elements. It is more conveniently prepared by passing
chlorine over a strongly heated mixture of silicon dioxide and carbon.
The gaseous products enter a condenser in which the tetrachloride
assumes the liquid form :
2CL, + Si02 + 2C -» SiCl4 + 2CO.
Chlorine is unable to displace oxygen from combination with silicon,
and has, therefore, when alone, no effect upon sand. In the above
action, therefore, the carbon is used to secure the oxygen while the
chlorine combines with the silicon. This kind of interaction is in
some degree different from any which we have hitherto encountered.
It bears no special name, but the principle underlying it is very com-
monly employed (see, e.g., Chlorides of boron and aluminium).
Silicon tetrachloride is a colorless liquid (b.-p. 59°) which fumes
SILICON AND BORON 521
strongly in moist air, giving silicic acid. It acts violently upon cold
water, and in this respect differs from carbon tetrachloride :
SiCl4 + 4H20 -> 4HC1 + Si(OH)4.
The silicic acid (<?.*'.) which is formed soon appears as a gelatinous pre-
cipitate.
When silicon is heated in a stream of dry hydrogen chloride, a
mixture of silicon tetrachloride and silico-chloroform SiHCl8 is pro-
duced. The latter is a volatile liquid boiling at 34°.
Silicon Tetrafluoride. — When strong hydrofluoric acid acts upon
sand, this gas is liberated :
SiO2 + 4HF -» 2H20 + SiF4.
Since the water interacts with the tetrafluoride (see below), the latter
is usually made by the use of powdered calcium fluoride and excess
of sulphuric acid. In this way the hydrogen fluoride is' generated in
contact with the sand, and at the same time the sulphuric acid takes
possession of the water. Hydrofluoric acid acts in a corresponding
way upon all silicates (q.v.), whether these are minerals or are artificial
silicates like glass (cf. p. 243).
Silicon tetrafluoride is a gas which becomes solid, without liquefy-
ing (cf. p. 463), when cooled to —102°. It fumes strongly in moist
air, and acts vigorously upon water. This interaction is different from
that of the tetrachloride, because the excess of the tetrafluoride forms
a complex compound with the hydrofluoric acid :
SiF4 + 4H20 -> Si(OH)4 ( + 4HF) (1)
(4HF) + 2SiF4 -> 2H2SiF6 (2)
_
3SiF4 + 4H2O -> Si(OH)4 + 2H2SiF
The silicic acid is precipitated in the water, and may be separated by
filtration, leaving a solution of hydrofluosilicic acid.
Hydrofluosilicic Acid. — This acid is stable only in solution.
When the water is removed by evaporation, silicon tetrafluoride is given
off, while most of the hydrogen fluoride remains to the last. Its salts
are decomposed in a corresponding way when they are heated. This
acid is used in analysis chiefly because its potassium and sodium salts
are amongst the few salts of these metals which are relatively
522 INORGANIC CHEMISTRY
insoluble in water. The barium salt is also insoluble, but most of the
salts of the heavy metals are soluble.
Silicon Dioxide. — This substance may be made in the form of a
fine white powder by heating precipitated silicic acid. It is found in
many different forms in nature. In large, transparent, six-sided prisms
with pyramidal ends it is known as quartz or rock crystal. When col-
ored by manganese and iron it is called amethyst, when by organic
matter, smoky quartz. A special arrangement of the structure gives
cat's eye. Amorphous forms of the same material, often colored brown
or red with ferric oxide, are agate, jasper, and onyx, the last much used
in making cameos. Infusorial or diatomaceous earth is composed of the
tests of minute organisms. Slightly hydrated forms of silica are the
opal and flint. A crystalline variety belonging to the hexagonal sys-
tem, but showing entirely different crystalline forms, occurs occasion-
ally in minute crystals, and is called tridymite.
Silicon dioxide, although differing profoundly from carbon dioxide
in its physical nature, nevertheless behaves like the latter chemically.
Thus, when boiled with potassium hydroxide solution (cf. p. 482), it
forms potassium orthosilicate :
Si02 + 4KOH -* K4Si04 + 2H20.
The salt is left as a gelatinous solid ("soluble glass ") when the water
is evaporated. The silicates of potassium and sodium may also be
obtained by boiling sand with the carbonates of these metals, carbon
dioxide being displaced. They are produced more rapidly, however,
as metasilicates (see below), by fusing the sand with the alkali
carbonates :
Si0 + KC0 -> KSi0 + C0.
Silica is found in the hard parts of straw, of some species of horse-
tail (equisetum], and of bamboo. In the form of whetstones it is used
for grinding. The clear crystals are employed in making spectacles
and optical instruments. Pure sand is used in glass manufacture
(gvtf.). Infusorial earth, on account of the tubular form of many of the
minute structures of which it is composed, can absorb three times its
own weight of nitroglycerine, and is therefore employed in making
dynamite. Kecently, small pieces of chemical apparatus have been
manufactured by fusing quartz in the oxyhydrogen flame. The mate-
rial is very difficult to work, on account of its inf visibility ; but owing
SILICON AND BORON 523
to the low coefficient of expansion of silica, the vessels fashioned out
of it can be heated or cooled as suddenly as we choose, without risk of
fracture.
Silicic *lcid* — When acids are added to a solution of sodium sili-
cate, silicic acid is set free. After a little delay it usually appears as a
gelatinous precipitate. When, however, the silicate is poured into
strong hydrochloric acid, no precipitation occurs. The silicic acid
remains in colloidal solution (see below). The acid before precipi-
tation is supposed to be orthosilicic acid :
Na4SiO4 + 4HC1 -> 4NaCl + Si(OH)4,
2HC1 + H2O -* 2NaCl + Si(OH)4,
but the gelatinous precipitate when it has been dried contains a
smaller proportion of the elements of water. There seem to be no
definite stages, indicating the existence of various acids, such as we
observe with phosphoric acid. We should expect the vapor tension of
the water to decrease by steps, each of which should correspond to some
acid of a particular degree of hydration (cf. p. 122), but nothing of the
sort is observed. The final product of drying is the dioxide. A jelly
of the kind described above is called a hydrogele. A solution such as
that containing dissolved silicic acid is called a hydrosole.
Silicic acid is a very feeble acid, and, therefore, gives no salt with
ammonium hydroxide. For the same reason silicic acid can be com-
pletely displaced from its salts, even by so weak an acid as carbonic acid.
Since the water in some geyser regions contains alkali silicates, the action
of the carbon dioxide in the air causes deposition of silica round the
places where the water issues from the ground. This form is known
as silicious sinter. Striking scenic effects, as in the white and pink
terraces of New Zealand, are sometimes produced by this method of
deposition.
Colloidal Solution. — Silicic acid in colloidal solution can be
separated from the other dissolved substances by a method dis-
covered by Graham and called dialysis. The solution is placed in
a vessel whose bottom is composed of vegetable parchment, or some
animal membrane, and the whole is hung in a vessel of water. Salts
which may be present pass slowly through the membrane into"the
water outside, and, when the latter is frequently changed, all of the
material of this class in the solution may finally be removed.
524 INORGANIC CHEMISTRY
The silicic acid, however, remains confined in the cell. Substances
possessing the ability to pass through a membrane are usually crystal-
line, and in the present connection are therefore called crystalloids.
On the other hand, materials of a gluey nature are unable to traverse
the partition,' and are named colloids. This behavior must not
confused with that in which a semi-permeable membrane (p. 284) is
question. In the latter case all dissolved substances are equally
unable to penetrate the membrane.
A pure aqueous solution of a colloid lacks the properties charac-
teristic of solutions. It boils and freezes at the same temperature as
the pure solvent, and independent evidence shows that, although in
too fine a state of division to be retained by a filter paper, the colloid
is nevertheless simply suspended in the liquid.
Colloidal solutions of other substances which are insoluble in water,
such as ferric hydroxide, aluminium hydroxide, and starch, and even
of metals like gold, silver, and platinum, can likewise be made.
Silicates. — While, in the absence of definite knowledge, silicic
acid is presumed to be the ortho-acid Si(OH)4, and no other silicic acids
have been made, the salts are most easily classified by imagining them
to be derived from various acids representing different degrees of hy-
dration of the dioxide, or, to put it the other way, different degrees of
dehydration of the ortho-acid. The following equations show the rela-
tion of the ortho-acid to some of the silicic acids whose salts are most
commonly found amongst minerals :
H4Si04 - H20 -> H2Si03 Metasilicic acid.
2H4Si04 - H20-»H6S2O7 ] _. ...
2H4Si04-3H:0-,H2Si206 |DlSlllC1C aClds"
3H4Si04- 4H20 -> H4Si3O8 Trisilicic acid.
Di- andtrisilicates are those derived from acids containing two and three
units of silicon, respectively, in the formula.
The composition of minerals is often exceedingly complex. This is
due to the fact that amongst them mixed salts (p. 359) are exceedingly
common, in which the hydrogen of the imaginary acid is displaced by
two or more metals in such a way that the total quantity of the metals
is equivalent to the hydrogen. The following list presents in tabular
form some typical or common minerals arranged according to the
above mentioned classification :
SILICON AND BORON 525
f Zircon, ZrSi04
Orthosilicates (H4Si04) J Garnet> Ca3Fe2"I(Si04)8
j Mica, KH2Al3(Si04)3
[Kaolin, H2AL,(Si04)8,H2O
f Wollastonite, CaSi03
Metasilicates (H2Si08) EeWl> Gl3Al2(Si08)6
1 Enstatite, MgSi03
Hornblende and augite
Disilicate (H6S,07) Serpentine, Mg3Si2O7,
Trisilicate (H4Si808) Orthoclase (feldspar), KAlSi308
It will be seen that the total valence of the metal units is equal to that
of the acid radicals. Thus, in beryl there are six equivalents of
glucinum (beryllium) and six of aluminium, taking the place of twelve
units of hydrogen in (H2Si03)6.
Mica, which is obtained in large sheets from Farther India, is used
in making lamp-chimneys and as an insulator in electrical apparatus.
Kaolin, or clay, like mica, is an acid orthosilicate. It is formed in
nature as the result of the action of water and carbonic acid upon
minerals like feldspar. In such cases, those elements which can form
carbonates, like potassium, magnesium, and calcium, are usually dis-
placed from combination with the silicic acid. Aluminium, however, is
too feeble a base to form a carbonate, and is thus left in combination
as silicate. A clay containing lime is called a marl, and one contain-
ing sand, a loam.
Some of these minerals frequently occur mixed together as regular
components of certain igneous rocks. Thus, granite is a more or less
coarse mixture of quartz, mica, and feldspar. Frequently the oblong,
flesh-colored or white crystals of the last are large and very conspicu-
ous. In basalt and porphyry the components are usually more inti-
mately mixed and less easily visible to the eye. Lava is the name for
any rock recently ejected from a volcano. Pumice-stone is the name
given to the parts of the lava which are porous, having acquired this
texture from the expansion of bubbles of gas consequent upon release
of pressure. Sandstone is composed of sand cemented together by clay
or by calcium carbonate, and colored brown or yellow by ferric oxide.
Since many silicates are not affected by acids, or at least are affected
with extreme slowness, special means have to be taken to get the con-
stituents of such minerals into soluble form for the purpose of analysis.
526 INORGANIC CHEMISTRY
Two methods are in use. Sometimes the finely powdered mineral is
heated in a platinum dish with hydrofluoric acid until all the silicon
has passed off in the form of the tetrafluoride. Since the use of the
fluorides would lead to difficulties in the course of the analysis, the
resulting mixture of fluorides of the metals is next heated strongly
with concentrated sulphuric acid, and the mixture of sulphates thus
produced is treated according to the usual routine. In other cases the
finely powdered mineral is fused at a white heat with a mixture of
potassium and sodium carbonates. In this way carbonates of the
metals are formed, along with potassium and sodium silicate. Treat-
ment with water dissolves the latter, and leaves the carbonates to be
handled in the usual way.
BOKON.
As regards chemical relations, boron, being a uniformly trivalent
element, is a member of the aluminium family. Yet it is a pronounced
non-metal, and its oxide and hydroxide are almost wholly acidic ;
aluminium is a metal, and with its oxide and hydroxide basic proper-
ties predominate. Boron and its compounds really resemble carbon
and silicon and their compounds in all chemical properties except the
property of valence.
;
Occurrence. — Like silicon, boron is found in oxygen compounds,
namely, in boric acid (y.v.) and its salts. Of the latter, sodium tetra-
borate Na^B.^, or borax, came first from India under the name of tin-
cal. It constitutes a large deposit in Borax Lake in California. Cole-
manite, Ca^C^, 5H20, from California, and other complex borates,
furnish a large part of the commercial supply of compounds of boron.
Preparation. — When boric oxide is heated with powdered mag-
nesium (B203 + 3Mg — > 3MgO + 2B), amorphous boron can be sep-
arated with some difficulty from the borides of magnesium in the .
resulting mixture. When excess of powdered aluminium is used, hard
crystals of boron, containing aluminium, are found in the solidified
metal. They may be separated by interaction of the metal with dilute
hydrochloric acid.
Properties. — Amorphous boron is a black powder. It unites
with the same elements as does silicon (p. 519), but with somewhat
greater activity. Like carbon (p. 379), it is also oxidized by hot, con-
SILICON AND BORON 527
centrated sulphuric or nitric acid, the product being boric acid. The
crystalline variety is less rapidly attacked in each case. Both kinds
interact with fused potassium hydroxide, giving a borate :
2B + 6KOH -> 2K3BOS + 3H2.
Hydrides and Halides of Boron. — When magnesium boride
Mg3B2 is treated with hydrochloric acid, a gas containing much hydro-
gen, and possibly several hydrides of boron, is given off. By cooling
the mixture with liquid air, Ramsay obtained a white solid which,
when it resumed the gaseous condition, appeared to be B3H8.
By combined action of carbon and chlorine on boric oxide, using
the principle employed in preparing silicon tetr a chloride (p. 520), the
trichloride of boron may be made. It is likewise formed by direct
union of the free elements. It is also made easily by heating boric
acid and phosphorus pentachloride, the action being an example of the
behavior of the latter towards hydroxyl compounds (cf. p. 463) :
B(OH)3 + 3PC15 -* BC13 + 3POC18 + 3HC1.
The products are separated by fractional distillation. Boron trichloride
is a liquid which boils at 18°, fumes strongly in moist air, and is com-
pletely hydrolyzed by water.
Boron trifluoride BF8 is made by the interaction of calcium fluoride
and sulphuric acid with boron trioxide. The mode of preparation and
the properties of the substance recall silicon tetrafluoride (p. 521). It
interacts with water, like the latter, giving boric acid and hydrofluo-
boric acid :
4BF3 + 3H2O -* B(OH)3 + 3HBF4.
The boric acid, not being very soluble, is precipitated. Hydrofluoboric
acid is known only in solution, although many of its salts are stable.
Boric Acid. — Boric acid (boracic acid) is somewhat volatile with
steam (cf. p. 471), and is found in Tuscany in jets of water vapor
(soffioni) which issue from the ground. Water, retained in small
basins by brickwork, is placed over the openings, and from this water,
after evaporation by the help of the steam of the soffioni themselves,
boric acid is obtained in crystalline form. As boric acid is very
feeble, and withal little soluble, it may also be made by interaction of
sulphuric acid and concentrated borax solution :
Na2B407 + H,S04 + 5H2O <=» Na^SO, + 4H3B0
528 INORGANIC CHEMISTRY
Boric acid crystallizes from water in thin white plates, which are
soapy (like graphite and talc) to the touch. Its solubility in water is 4
parts in 100 at 19° and 34 in 100 at 100°. The solution scarcely affects
litmus. It confers a green tint on the Bunsen flame. The effect is
best seen by setting fire to the vapor of a boiling alcoholic solution of
the acid. This behavior is used as a test for the acid. At 100° the acid
slowly loses water, leaving metaboric acid HBO2, and at 140° tetra-
boric acid is formed : 4HB02 — H20 — » H2B407. Strong heating gives
the oxide B203. When dissolved in water, these dehydrated compounds
revert to boric acid. The solution of boric acid in water is used as an
antiseptic in medicine, and as a preservative for milk and other foods.
Borates. — Borates derived from orthoboric acid are practically
unknown. The most familiar salt is borax or sodium tetraborate.
The decahydrate Na2B4O7, 10H20, which crystallizes from water at 27°
in large, transparent prisms, and the pentahydrate which crystallizes
at 56°,* are both marketed. They are made by crystallization of native
borax. In Germany, borax is prepared from boracite, found at Stass-
furt, by decomposing a solution of the mineral with hydrochloric acid :
MgCl2, 2Mg3B8016 + 12HC1 + 18H20 -> 7MgCl2 + 16B(OH)8.
The boric acid is redissolved in boiling water, and sodium carbonate
is added :
4B(OH)8 + Na^COg-^Na^O, + 6H20 + C02.
In California it is made from colemanite by interaction with sodium
carbonate.
Since boric acid is a feeble acid, borax is extensively hydrolyzed by
water, and the solution has a marked alkaline reaction.
When heated with oxides of metals, sodium tetraborate behaves
like sodium metaphosphate (c/.p. 468), and is used in the form of beads
in analysis. If its formula be written 2]STaB02,B2O3, it will be seen
that a considerable excess of the acid anhydride is contained in it, and
that, therefore, a mixed metaborate may be formed by union with
some basic oxide. Thus, with a trace of cupric oxide, the bead is
tinged with green, from the presence of a compound like 2NaB02,
Cu(BO2)2. Cobalt compounds give a deep-blue color to the bead. For
the same reason, borax is used in hard soldering. The hard solder
* For explanation of the relation of temperature of crystallization to degree of
hydration, see under Manganous sulphate.
SILICON AND BORON 529
(brass) is placed, with a little borax, upon the joint between the objects
of copper or brass which are to be soldered. At the temperature pro-
duced by the blast-lamp the borax dissolves the superficial coating of
oxides, and the molten solder is able to " wet " the clean surfaces. A
substance used thus, to bring infusible bodies into a fusible form of
combination, is called a flux.
Boron Trioxide. — The oxide, as made by heating boric acid, is a
glassy white solid. It is obtained also by burning boron in oxygen.
Being almost perfectly involatile, it is able, when heated with salts, to
displace other acid anhydrides which can be vaporized :
K2S04 + BaO,<=> 2KBO2 + SO,J.
It has a slight tendency to act as a basic oxide. With fuming sul-
phuric acid it gives a boryl pyrosulphate BO,HS207 which is decom-
posed by water, and with phosphoric acid a phosphate BP04 which is
stable.
Nitride and Carbide. — One of the difficulties in making free
boron is due to its very great affinity for nitrogen, with which, when
heated in the air, it unites to form a nitride BK. This compound is
more easily made by heating borax with ammonium chloride :
Na2B407 + 4NH4C1 -* 4BN + 2NaCl + 7H20 + 2HC1.
The nitride is a white solid which is easily decomposed when heated
in a current of steam (cf. p. 417) :
BN + 3H20 -> B(OH)8 + NH8.
A carbide of boron B6C is made by heating the free substances in
the electric furnace. It is harder than carborundum, and stands
next to the diamond in respect to hardness.
•
Exercises. — 1. Why is the fact that carborundum is not affected
by water or acids worthy of mention?
2. Compare and contrast the elements carbon and silicon, and
their corresponding compounds.
3. Why are there no colloidal solutions of iron and zinc (p. 524) ?
4. What would be the interaction between aqueous solutions of an
ammonium salt and of sodium orthosilicate ?
CHAPTER XXXII
THE BASE-FORMING ELEMENTS
THERE are two ways in which the chemistry of a given set of ele-
ments may be described. We may take up the elements in succession,
and discuss under each its physical and chemical properties, and the
manufacture and behavior of a certain number of its compounds, such
as the oxide, hydroxide, nitrate, and sulphate. Or we may arrange our
major classification according to the properties and the forms of com-
bination, and detail the facts about the same set of elements under
each. Both methods have such advantages that neither can be sacri-
ficed entirely. We shall, therefore, adopt the former plan for our divis-
ion into chapters, following the usage already employed for the non-
metals. In the present chapter a preliminary view of the chemistry
of the metals will be given according to the second method.
Physical Properties of the Metals. — A knowledge of the phys-
ical properties of the metals is, to the chemist, of the greatest impor-
tance in connection with their manufacture and treatment. The
following brief statement in regard to some of these properties is illus-
trative rather than exhaustive. It should be noticed that the properties
of a metal vary according as the specimen has been prepared by roll-
ing, casting, or some other process. Numerical values, therefore, when
given, are only approximate.
Metals show what is commonly called a metallic luster, but, as a
rule, they do so only when in compact form. Magnesium and alumin-
ium exhibit it when powdered, but most of the metals when in this
condition are black. In compact masses the metals are usually silvery
white in color. Gold and copper, which are yellow and red respect-
ively, are the conspicuous exceptions.
The metals can all be obtained in crystallized form, when a fused
mass is allowed to cool slowly and the unsolidified portion is poured
off. In almost all cases the crystals belong to the regular system.
With the metals most nearly allied to the non-metals, however, they do
not. Thus, the crystals of antimony and bismuth belong to the hex-
agonal system, and those of tin to the square prismatic.
530
THE BASE-FORMING ELEMENTS
531
The metals vary in specific gravity from lithium, which is little
more than half as heavy as water (sp. gr. 0.59), to osmium, whose
specific gravity is 22.5. Those which have a specific gravity less than
5, namely, potassium, sodium, calcium, magnesium, aluminium, and
barium, are called the light metals, and the others the heavy metals.
Most metals are malleable, and can be beaten into thin sheets with-
out loss of continuity. Those which are allied to the non-metals, how-
ever, such as arsenic, antimony, and bismuth, are brittle, and can
be reduced to powder in a mortar. Zinc becomes malleable only when
heated to 150°. The order of the elements in respect to this property,
beginning with the most malleable, is : Au, Ag, Cu, Sn, Pt, Pb, Zn,
Fe, Ni.
The tenacity of the metals places them in an order different from
the above. It is measured by the number of kilograms which a piece
of the metal 1 sq. mm. in section can sustain without breaking. The
values are as follows : Fe 62, Cu 42, Pt 34, Ag 29, Au 27, Al 20, Zn 5,
Pb2.
The hardness is measured by the ease with which the material may
be disintegrated by a sharp, hard instrument. Potassium is as soft as
wax, while chromium is hard enough to cut glass.
The temperature at which the metal fuses has an important bearing
on its manufacture. Most of the following melting-points are only
approximate :
Mercury .
-40°
Zinc .
420°
Cast iron .
1150°
Potassium
62°
Antimony .
437°
Nickel. . .
1500°
Sodium .
96°
Aluminium .
700° (?)
Platinum .
1780°
Tin . . .
230°
Magnesium .
750°
Iron (pure) .
1800°
Bismuth .
264°
Silver
954°
Manganese
1900°
Cadmium
320°
Copper .
1057°
Chromium
2000°
Lead .
326°
Gold . . .
1075°
Iridium .
2200°
It will be seen that mercury is a liquid, that potassium and sodium
melt below the boiling-point of water, and that the metals down to the
foot of the second column can be melted easily with the Bun sen flame.
The metals osmium, molybdenum, uranium, tungsten, and vanadium*
have melting-points above the melting-point of platinum.
The methods of manufacture and the treatment of metals are much
influenced also by their volatility. The following are easily distilled :
Mercury, b.-p. 357°; potassium and sodium, b.-p. aboxit 700° ; cadmium,
b.-p. 770° ; zinc, b.-p. 950°. Even the most involatile metals can be
converted into vapor in the electric arc.
532 INORGANIC CHEMISTRY
In many cases molten metals dissolve in one another freely. The
mixtures are called alloys, and in some cases have the properties of
solid solutions. Sometimes, as in the case of lead and tin, mixtures can
be formed in all proportions. On the other hand, the solubility may be
limited, as in the case of zinc and lead, where only 1.6 parts of the
former dissolve in 100 parts of the latter. The colors of alloys are
not the average of those of the constituents. Thus, a mixture contain-
ing copper with 30 per cent of tin is perfectly white. A similar mix-
ture with 30 per cent of zinc is pale yellow. The nickel alloy used in
coining contains 75 per cent of copper and 25 per cent of nickel, yet it
shows none of the color of the former. Thirty per cent of gold may
be added to silver without conferring any yellow tint upon it.
Some of the properties of alloys are classifiable by the ordinary
laws of solution. Thus, a foreign body lowers the vapor tension of a
solvent (p. 161), and so the presence of a foreign metal diminishes the
ease with which particles can be torn from the surface by any means
whatever. That is to say, it increases the hardness. A foreign metal
also lowers the melting-point (p. 163). In many cases the metal be-
comes less active when alloyed. Thus, a mixture of gold and silver
containing twenty-five per cent of the latter does not interact visibly
with nitric acid. It is necessary to bring the amount of silver up to
75 per cent at least ("quartation") in order that the silver may be
freely attacked by the warm acid. The gold remains in any case un-
touched. The malleability and the conductivity for heat and elec-
tricity are diminished by addition of a foreign metal. Copper, whose
commercial applications depend largely on the first and third of these
properties, is much affected in respect to them by the presence of even
small traces of impurities.
Alloys in which mercury forms one of the components are known
as amalgams (Gk. /xaXay/wx, a soft mass), and are formed with especial
ease by the lighter metals. Of the common metals, iron is the least
miscible with mercury.
The good conductivity of metals for electricity distinguishes them
with some degree of sharpness from the non-metals. They show con-
siderable variation amongst themselves, silver conducting sixty times
as well as mercury. The conductivity increases as the temperature is
lowered, and this fact is taken advantage of in the measurement of the
temperature of liquefied gases. The platinum resistance thermometer
consists chiefly of a wire of platinum. The resistance of this metal
diminishes so rapidly, with decreasing temperature, that measurement
THE BASE-FOKMING ELEMENTS 533
of its resistance can be used for the accurate determination of the tem-
perature. In the following table the conductivities of the metals are
expressed in terms of the number of meters of wire 1 sq. mm. in sec-
tion which, at 15°, offer a resistance of one ohm :
Silver, cast .... 62.89 Nickel, cast . . . 7.59
Copper, commercial . 57.40 Iron, drawn . .7.55
Gold, cast . . . .46.30 Platinum . . 5.7-8.4
Aluminium, commercial 31.52 Steel 6.43
Zinc, rolled .... 16.95 Lead .... 4.56
Brass 14.17 Mercury .... 1.049
The resistance at 0° of a column of mercury 1 sq. mm. in section
and 1.063 meters long is called the ohm, and is employed in express-
ing the conductivities of solutions (p. 328). To compare the above fig-
ures with those given for solutions, however, it must be recalled that,
in the measurement of the conductivities of the latter, a column only 1
cm. in length and of 1 sq. cm. area was employed, so that the figures
representing the conductivities of solutions are on a scale approxi-
mately ten thousand times as great as those presented in the above
table. Thus, normal hydrochloric acid (p. 328) has a conductivity on
the above scale of 0.0301, or less than a thirtieth of that of mercury.
General Chemical Relations of the Metallic Elements.—
Since most of the compounds of the metals are ionogens, their solutions,
except when the metal is a part of a compound ion or of a complex ion
(see below), all contain the metal in the ionic state, and the resulting
substances, such as kalion and cuprion, have constant properties,
irrespective of the nature of the negative ion with which they may be
mixed. The properties of the ions, simple and compound, are much
used in making tests in analytical chemistry. On the other hand, the
chemical properties of the oxides and of the salts in the dry state are
of importance in connection with metallurgy.
There are six chemical properties which are more or less generally
characteristic of the metallic elements. The first two of them have
already been discussed somewhat fully.
1. The metals are able by themselves to form positive radicals of
salts, and, therefore, to exist alone as positive ions (pp. 337, 404).
2. The oxides and hydroxides of the metals are basic (pp. 119, 404).
3. The halogen compounds of the typical metals are little, if at all,
hydrolyzed by water (see next section).
534 INORGANIC CHEMISTRY
These three properties are characteristic of all metals, and form the
basis of the distinction between metals and non-metals. The remaining
three apply to many metals, but no one of them applies to all.
4. An oxide* or hydroxide which is basic may also be acidic, as, for
example, zinc hydroxide (p. 405). Even when this is not the case,
some other oxide of the metal may be acidic exclusively, as is manga-
nese heptoxide (p. 405). In consequence of either of these facts, a
metal may form part of the negative radical of a simple salt, and there-
fore be found in a negative ion, as, Zn02" or MnO/.
5. Some salts of certain metals combine with those of others to
give complex salts (p. 363 and next section below). Of this sort are
the complex cyanides, such as K.Ag(CN)2 and K4.Fe(CN)6. A metal
thus forms part of the negative radical of a salt of a complex acid,
and therefore is found in an anion like Ag(CN)2' or PtCl/'.
6. Some metals also form parts of complex cations which are con-
tained in solutions of molecular compounds (p. 443), chiefly those of
salts with ammonia. Thus, when AgCl, 3NH3 is dissolved in water, or
when ammonium hydroxide is added to a solution of a salt of silver, the
positive ion is found to be Ag(NH3)2* (see under Copper). (For a de-
tailed illustration of the application of these six criteria, see discussion
of the chemical relations in the nitrogen family, Chap, xli.)
Aside from these points, many features in the behavior of metals
and their compounds are summed up in the electromotive series (p. 362).
The reader should re-read all the parts referred to above before proceed-
ing farther. He should also reexamine the various kinds of chemical
changes discussed on p. 187 and particularly the varieties of ionic
chemical change on p. 360.
Hydrolysis of Halogen Compounds, Used to Distinguish
Metals from Non-Metals. — We have seen that the halogen com-
pounds of phosphorus (p. 463), of sulphur (p. 398), of silicon (p. 521),
and of other non-metals, are completely hydrolyzed by water, giving
the hydrogen halide, and an acid which contains the hydroxyl of the
water :
PCL, + 3HOH -t 3HC1 + P(OH)3.
Now, those elements whose halogen compounds are not hydrolyzed by
water, or, at all events, are only partly hydrolyzed, are the ones classed
as metals. Thus, sodium chloride is not decomposed appreciably by
THE BASE-FORMING ELEMENTS 535
water, and cupric chloride, like cuprie sulphate (p. 344), is but slightly
hydrolyzed, and its solution has a faint acid reaction :
CuCl2 + 2H20 <=; 2HC1 + Cu(OH)2.
In a few cases, as with the chlorides of antimony (q.v.~) and of bis-
muth (q.v.), a considerable proportion, but not all, of the halogen is re-
moved from each molecule :
SbCl3 + H20<=±2HC1 + SbOClJ.
The resulting mixture is strongly acid, and the product (antimony
oxychloride) is a definite compound, of the nature of a mixed salt
(p. 359), known as a basic salt. The difference is that, with the halides
of the metals, the action of water is notably reversible, while with halides
of the non-metals it is not so.
Hydrolysis of the halides of the metals is increased by rise in tem-
perature and by dilution of the solution (addition of more water), and
also gains headway when one of the products of hydrolysis is thrown
down as a precipitate. The last two influences are the ones which
normally permit a reversible action to approach completion (p. 259).
The halogen compounds are chosen as the basis of this criterion
because the halogen acids are active and would reverse the hydrolysis
completely, and leave no acid reaction, if the result depended upon
them alone. It is the lack of activity in the base, and the ten-
dency of its molecules to be formed from the metal ions of the salt
and the OH' of the water (p. 314), that determine the slight hydrolysis,
when it occurs. Thus, this criterion is simply another means of recog-
nizing whether or not the hydroxide of the element is a strong (much
ionized) base, and its application gives, therefore, the same result in
each case as does the employment of the second of the chemical charac-
teristics of metallic elements (see above).
Other, non-halide salts of the metals, even of the most active, may be
extensively hydrolyzed by water. Thus, sodium sulphide is decomposed
by it (p. 375) to the extent of one-half. But here the solution is
alkaline in reaction :
Na2S + H.,0 -»NaSH + NaOH,
owing to the small ionizatioii of the SH' ion, and the result is due to
the feebleness of the acid (H2S). Indeed, the great activity of the
base is demonstrated by the final reaction of the solution, and the
536 INORGANIC CHEMISTRY
metallic nature of sodium is therefore not impugned by the existence
of hydrolysis per se in such a salt as this, but is rather confirmed.
To sum up : An alkaline reaction shows that we have a solution
a salt of an active metal with a weak acid ; an acid reaction that
have a salt of an inactive metal, which may even verge on the non-
metallic, with an active acid. Salts of two active components give
neutral solutions. Thus, as a particular case, the halogen compounds
of the typical metals are not perceptibly hydrolyzed by water, and the
hydrolysis of the halide of a less pronounced metal takes place with
acid reaction, and is easily reversible by excess of either product. A
salt of an acid and base both of which are weak is hydrolyzed, often
completely. Aluminium carbonate and ammonium silicate are examples
of salts which, for this reason, are completely hydrolyzed. The result-
ing mixture may have an acid or a basic reaction, if the acid or the
base is sufficiently soluble and sufficiently active. Thus, ammonium
sulphide solution is alkaline.
Salts of Complex Acids. — These salts are of many classes, and
arise by direct union of two salts. Thus we have cyanides like
K.Ag(CN)2 (potassium argenticyanide), K.Cu(CN)2 (potassium cupro-
cyanide), and K4.Fe(CN)6 (potassium ferrocyanide) :
K.CN + AgCN | ->K.Ag(CN)2.
The complex sulphur compounds of arsenic, antimony, and tin (q.v.},
such as sodium sulphostannate Na2.SnS3, are made in the same way :
+ SnS <
Many double halides are of a like nature, as K2.PtClg (potassium chloro-
platinate), H2.PtCl6 (chloroplatinic acid), and Na.AuCl4 (sodium
chloraurate). In every case the metal of one of the original salts is
contained in the negative radical (the anion). Hydrofluosilicic acid
(p. 521) is a compound analogous to those of the last group, excepting
that SiF4, from which it is formed, is not a salt, and silicon is not a metal.
The characteristic of these compounds is that, when they are
ionized, the less positive metal is contained in a complex anion like
Ag(CN)/, SnS3", PtCl,."- In fact, they behave, in most respects, like
ordinary single salts. Thus, they undergo double decomposition in the
normal manner, the complex ion acting as a whole. For example, a
soluble ferrocyanide with a zinc salt gives zinc ferrocyanide :
K4.Fe(CN)6 + 2Zn.S04<z± Zn2Fe(CN)4 + 2K2.SO4.
THE BASE-FORMING ELEMENTS 537
This behavior distinguishes salts of complex acids from double salts
(p. 360). In typical cases the latter resemble the former, indeed,
only in their mode of preparation (by union of two simple salts), and
in solution are resolved once more into the component salts and their
separate ions. The distinction must be regarded, however, only as a
means of rough classification for practical purposes. In reality, all
complex ions give at least a trace of the simpler ions, and many are de-
composed to a noticeable extent. At the other extreme, the double salts
form complex ions in appreciable amounts in concentrated solutions.
In harmony with the above characteristic, salts of complex acids,
like potassium ferrocyanide, when treated with acids, like sulphuric
acid, undergo double decomposition, and give the free complex acid
(here ferrocyanic acid) :
K4.Fe(CN)6 + 2H2.S04<i±H4Fe(CN)6J + 2K2.S04.
But in many cases the acids are unstable, those of the cyanides, for
example, giving up hydrocyanic acid and those of the complex sulphur
compounds, hydrogen sulphide :
H2.SnS3 fc? H2S f + SnS, J.
The complex halogen acids, however, like H2.PtClg and H.AuCl4, are
stable, and are fairly active as acids.
The similarity of these compounds to oxygen acids, and their salts, may be seen
if we imagine them to be cases in which cyanogen, sulphur, chlorine, and other
radicals have taken the place of oxygen in the anion. Thus, [(CN)2]n, [C1J11, and
S11 are equivalent to Ou. Many of the corresponding oxygen compounds are
actually known, as sodium stannate Na2.Sn03 and sodium aurate Na.AuO2, corre-
sponding respectively to Na2.SnS3 and Na.AuCl4.
The behavior of complex ions is discussed further under Copper.
Classification of the Metallic Elements by their Chemical
Relations. — In treating of the metallic elements and their com-
pounds we shall use the groupings provided by the periodic system
(p. 411). A division into eleven sets, which are described briefly, and
in general terms below, will be sufficient for the purpose of this book :
1. Metals of the Alkalies. — Lithium, sodium, potassium, rubi-
dium, caesium, and the radical ammonium NH4. These metals are
univalent, and their oxides and hydroxides have strongly basic proper-
ties. Their salts with active acids are not hydrolyzed in solution.
2, Metals of the Alkaline Earths. — Calcium, strontium, barium,
538 INORGANIC CHEMISTRY
and, possibly, radium. These metals are bivalent in all their com-
pounds. Their oxides and hydroxides are strongly basic, but the
latter are not so soluble in water as are the hydroxides of the former
family. The salts with active acids are not hydrolyzed.
3. Copper, Silver, and Gold. — These metals occupy the right-
hand side of the second column of the table (p. 411). Their alliance
with the alkali metals, their neighbors, is rather remote. Each of
them, however, gives compounds in which it is univalent, although, in
their commoner compounds, copper and gold are bivalent and triva-
lent, respectively. The oxides and hydroxides of copper and gold have
rather weak basic properties ; those of silver are much more active.
4. Beryllium, Magnesium, Zinc, Cadmium, and Mercury. — These
metals, occupying the right-hand side of the third column, are biva-
lent, although mercury forms a series of compounds in which it is
univalent as well. The oxides and hydroxides are feebly basic, those
of magnesium being the most basic.
5. Aluminium and the other metals on both sides of the fourth
column. — The metals of these groups are trivalent, and the oxides and
hydroxides are feebly basic in character. The hydroxide of aluminium
has also a feebly acidic tendency.
6. Germanium, Tin, and Lead, and the Titanium Family. — In ac-
cordance with their position in the periodic table these metals are all
quadrivalent. At the same time they act also as bivalent elements,
the compounds of this class in the case of lead being the more familiar.
The oxides and hydroxides are feebly basic, and are able also to play
the role of acidic oxides towards strong bases.
7. Arsenic, Antimony, and Bismuth, and the metals of the Vanadium
Group. — These elements, like nitrogen and phosphorus, form two sets
of compounds in which .they are trivalent and quinquivalent, respec-
tively. The acidic tendency of the oxides and hydroxides, which in
the last column was noticeable, is here much more pronounced. Both
oxides of arsenic are almost wholly acidic in behavior, and the pent-
oxides of the other elements are likewise acid anhydrides exclusively.
The trioxides are basic in a feeble way, and their salts are much hy-
drolyzed by water.
8. Chromium, Molybdenum, Tungsten, and Uranium. — These ele-
ments, occupying the left-hand side of the seventh column, exhibit a
considerable variety of valence. The maximum, however, is six in each
case. The oxides of the form CrO and Cr203, in which the elements are
bivalent and trivalent, are base-forming. Those of the form Cr03, in
THE BASE-FORMING ELEMENTS 539
which the elements are sexivalent, resemble sulphur trioxide and are
acid anhydrides.
9. Manganese. — This element, the only one in the eighth column
which has not yet been treated, gives several series of compounds
in which its valence varies from 2 to 7. Compounds derived
from the basic oxide MnO are the salts in which manganese is
most distinctly a metallic element. The highest oxide, Mn207, is an
acid anhydride.
10. Iron, Nickel, and Cobalt. — These elements give oxides which
are feebly basic. Iron gives two extensive series of compounds in
which it is bivalent and trivalent, respectively. Those of the former set
resemble the bivalent salts of manganese and zinc. Those of the latter
resemble the salts of aluminium. Cobalt and nickel in most of their
compounds are bivalent elements, and the behavior recalls that of the
compounds of bivalent manganese and zinc.
11. Palladium and Platinum Families. — The metals of these fam-
ilies have little chemical activity, and their compounds are easily decom-
posed by heating. Along with gold, silver, and mercury, which have
similar characteristics, they are sometimes grouped together under the
name of the noble metals.
Occurrence of the Metals in Nature. — The minerals from which
metals are extracted are known as ores. They present a compara-
tively small number of different kinds of compounds. Most of the
metals are found in more than one of these forms, so that in the fol-
lowing statement the same metal frequently occurs more than once.
When the metal occurs free in nature it is said to be native. Thus
we have native gold, silver, metals of the platinum group, copper,
mercury, bismuth, antimony, and arsenic (cf. p. 362).
The metals whose oxides are important minerals are iron, man-
ganese, tin, zinc, copper, and aluminium. The metals are obtained
commercially from the oxides in each of these cases.
The metals whose sulphides are used as ores are nickel, cobalt,
antimony, lead, cadmium, zinc, and copper.
From* the carbonates we obtain* iron, lead, zinc, and copper. Sev-
eral other metals, such as manganese, magnesium, barium, strontium,
and calcium occur in larger or smaller quantities in the same form of
combination.
The metals which occur as sulphates are those whose sulphates are
not freely soluble, namely, lead, barium, strontium, and calcium.
540 INORGANIC CHEMISTRY
Compounds of metals with the halogens are not so numerous.
Silver chloride furnishes a limited amount of silver. Sodium and pot-
assium chlorides are found in the salt-beds, and cryolite 3NaF, A1F3 is
used in the manufacture of aluminium.
The natural silicates are very numerous, but are seldom used for
the preparation of the metals. Many of them are employed for other
commercial purposes, kaolin (p. 525) being a conspicuous example of
this class.
•
Methods of Extraction from tJie Ores. — The art of extracting
metals from their ores is called metallurgy. Where the metal is native,
the process is simple, since melting away from the matrix (p. 367) is all
that is required. Frequently a flux (p. 529) is added, which combines
with the matrix, giving a fusible slag. Since the slag is a melted salt,
usually a silicate, and does not mix at all with the molten metal, sepa-
ration of the products is easily effected. When the ore is a compound,
the metal has to be liberated by furnishing a material capable of com-
bining with the other constituent. The details of the process depend
on various circumstances. Thus the volatile metals, like zinc and
mercury, are driven off in the form of vapor, and secured by conden-
sation. The involatile metals, like copper and iron, run to the bottom of
the furnace and are tapped off.
Where the ore is an oxide it is usually reduced by heating with
carbon in some form. This holds for the oxides of iron and copper,
for example. Some oxides are not reducible by carbon in an ordinary
furnace. Such are the oxides of calcium, strontium, barium, magnesium,
aluminium, and the members of the chromium group. At the tempera-
ture of the electric furnace even these may be reduced, but the carbides
are formed under such circumstances, and the metals are more easily
obtained otherwise. Recently, heating the pulverized oxide with finely
powdered aluminium has come into use, particularly for operations on
a small scale. Iron oxide is easily reduced by this means, and even
the metals manganese and chromium may be liberated from their
oxides quite readily by this action. This procedure has received the
name aluminothermy (y.vS) on account of the great amounts of heat
liberated. In the laboratory the oxides of the less active metals are
frequently reduced in a stream of hydrogen (cf. p. 362).
When the ore is a carbonate, it is first heated strongly to drive out
the carbon dioxide (cf. p. 480) : FeC08 <=> FeO + C02 \ , and then the
oxide is treated according to one of the above methods. When the
THE BASE-FORMING ELEMENTS 541
ore is a sulphide, it has to be roasted (cf. p. 378) in order to remove the
sulphur, and the resulting oxide is then treated as described above.
Chlorides and fluorides of the metals can be decomposed by heat-
ing with metallic sodium (cf. p. 518). This method was formerly em-
ployed in the making of magnesium and aluminium.
The metals which are not readily secured in any of the above ways,
can be obtained easily by electrolysis of the fused chloride or of some
other simple compound. Aluminium is now manufactured entirely by
the electrolysis of a solution of aluminium oxide in molten cryolite.
Compounds of the Metals : Oxides and Hydroxides. — The
oxides may be made by direct burning of the metal, by heating the
nitrates (cf. p. 444), the carbonates (cf. p. 480), or the hydroxides :
Ca(OH)2 <i± CaO -f- H2O f. They are practically insoluble in water,
although those of the metals of the alkalies and of the metals of the alka-
line earths interact with water rapidly to give the hydroxides. They are
usually stable. Those of gold, platinum, mercury, and silver decompose
when heated, yet with increasing difficulty in this order. The metals,
like the non-metals, frequently give several different oxides. Those of
the univalent metals, having the form K20, if we leave cuprous oxide
and aurous oxide out of account, have the most strongly basic qualities.
Those of the bivalent metals of the form MgO, when this is the only
oxide which they furnish, are base-forming. Those of the trivalent
metals of the form A^Og are the least basic of the basic oxides. The
oxides of the forms Sn02, Sb20B, Cr08, and Mn207, in which the metals
have valences from 4 to 7, are mainly acid-forming oxides, although
the same elements usually have other lower oxides, which are basic.
The hydroxides are formed, in the cases of the metals of the alkalies
and alkaline earths by direct union of water with the oxides. They are
produced also by double decomposition when a soluble hydroxide acts
upon a salt (cf. p. 350). All hydroxides, except those of the alkali metals,
lose the elements of water when heated, and the oxide remains. In
some cases the loss takes place by stages, just as was the case with
orthophosphoric acid (p. 465). Thus lead hydroxide Pb(OH)2 (q.v.)
first gives the hydroxide Pb20(OH)2, then Pb302(OH)2, and then the
oxide PbO. With the exception of those of the metals of the alkalies
and alkaline earths, all the hydroxides are little soluble in water. The
hydroxides of mercury and silver, if they are formed at all, are evi-
dently unstable, for, when either material is dried, it is found to con-
tain nothing but the corresponding oxide.
542 INORGANIC CHEMISTRY
Compounds of the Metals : Salts. — It may be said, in general,
that each metal may form a salt by combination with each one of
the acid radicals. In the succeeding chapters we shall describe only
those salts which are manufactured commercially, or are of special in-
terest for some other reason. The various salts will be described under
each metal. Here, however, a few remarks may be made about the
characteristics of the more common groups of salts. The salts are
classified according to the acid radicals which they contain.
The chlorides may be made by the direct union of chlorine with
the metal (cf. p. 74), or by the combined action of carbon .and chlorine
upon the oxide (cf. p. 520). The latter method is used in making
chromium chloride. The general methods for making any salt (p. 186),
such as the interaction of a metal with an acid, or of the oxide, hydrox-
ide, or another salt with an acid, or the double decomposition of two
salts, may be used also for making chlorides. The chlorides are for
the most part soluble in water. Silver chloride, mercurous chloride,
and cuprous chloride are almost insoluble, however, and lead chloride
is not very soluble. Most of the chlorides of metals dissolve without
decomposition, but hydrolysis is conspicuous in the case of the chlorides
of the trivalent metals, such as aluminium chloride and ferric chlo-
ride (cf. p. 344). The chlorides of some of the bivalent metals are
hydrolyzed also, but, as a rule, only when they are heated with water.
This is the case with the chlorides of magnesium, calcium, and zinc.
Most of the chlorides are stable when heated, but those of the noble
metals, particularly gold and platinum, are decomposed, and chlorine
escapes. The chlorides are usually the most volatile of the salts of a
given metal, and so are preferred for the production of the spectrum
(q.v.) of the metal, and for fixing the atomic weight of the metal by
use of the vapor density. Some of the metals form two or more differ-
ent chlorides. For example, indium gives InCl, InCl2, and InClg.
The sulphides are formed by the direct union of the metal with sul-
phur, or by the action of hydrogen sulphide or of some soluble sulphide
upon a solution of a salt (cf. p. 375). In one or two cases they are made
by the reduction of the sulphate with carbon. The sulphides, except
those of the alkali metals, are but little soluble in water. The sul-
phides of aluminium and chromium are hydrolyzed completely by
water, giving the hydroxides, and those of the metals of the alkaline
earths are partially hydrolyzed (cf. p. 376).
Some of the metals, when they are in the molten form, simply dis-
solve carbon, and, when they are cooled once more, deposit it in the
THE BASE-FORMING ELEMENTS 543
form of graphite. This is true particularly of platinum and iron.
The carbides are usually formed in the electric furnace by interaction
of an oxide with carbon (cf. p. 478). Some of them are decomposed by
contact with water, after the manner of calcium carbide, giving a
hydroxide and a hydrocarbon. Of this class are lithium carbide Li2C2,
barium and strontium carbides BaC2 and SrC2, aluminium carbide
A14C3, manganese carbide MnC, and the carbides of potassium and
glucinum. Others, such as those of molybdenum CMo2 and chromium
Cr8C2, are not affected by water.
The nitrates may be made by any of the methods used for preparing
salts. They are all at least fairly soluble in water.
The sulphates are made by the methods used for making salts, and
in some cases by the oxidation of sulphides. They are all soluble
in water, with the exception of those of lead, barium, and strontium.
Calcium sulphate is meagerly soluble.
The carbonates are prepared by the methods used for making salts.
They are all insoluble in water, with the exception of those of sodium
and potassium. The hydroxides of aluminium and tin are so feebly
basic that these metals do not form stable carbonates (c/.pp. 119, 523).
The phosphates and silicates are prepared by the methods used in
making salts. The former are obtained also by special processes
already described (p. 467). With the exception of the salts of sodium
and potassium, all the salts of both these classes are insoluble.
Solubility of Bases and Salts. — The solubilities of a few salts
at various temperatures have already been given (p. 157). The fol-
lowing table includes a much larger number of substances (142), and
gives the solubility at 18°. Each square contains two numbers ex-
pressing the solubility of the compound whose cation stands at the
head of the column and whose an ion is indicated at the side. The
solubility is that of the hydrate stable at 18°, where such exists.
The upper number in each case gives the number of grams of the an-
hydrous salt held in solution by 100 c.c. of water. The lower number
shows the number of moles in 1 1. of the saturated solution, and indi-
cates therefore the concentration in terms of a molar solution as
unity — the molar solubility. In the cases of the less soluble com-
pounds the values are not exact, but they will serve to show roughly
the relative solubilities when several substances are compared. The
numbers for small solubilities have been abbreviated. Thus, 0.064 =
0.0000004.
544 INORGANIC CHEMISTRY
Solubility of Bases and Salts in "Water at 18°.
K
Na
Li
Ag
Tl
Ba
Sr
Ca
Mg
Zn
Pb
01
32.95
3.9
35.86
5.42
77.79
13.3
0.0316
0.0410
0.3
0.013
37.24
1.7
51.09
3.0
73.19
5.4
55.81
5.1
203.9
9.2
1.49
0.05
Br
65.86
4.6
88.76
6.9
168.7
12.6
0.041
0.066
0.04
0.0215
103.6
2.9
96.52
3.4
143.3
5.2
103.1
4.6
478.2
9.8
0.598
0.02
I
137.5
6.0
177.9
8.1
161.5
8.5
0.063T-
0.071
0.006
0.0317
201.4
3.8
169.2
3.9
200
4.8
148.2
4.1
419
6.9
0.08
0.0S2
F
92.56
12.4
4.44
1.06
0.27
0.11
195.4
13.5
72.05
3
0.16
0.0j92
0.012
0.001
00016
0.032
0.0076
0.0214
0.005
0.0,5
0.07
0.003
NO,
30.34
2.6
83.97
7.4
71.43
7.3
213.4
8.4
8.91
0.35
8.74
0.33
66.27
2.7
121.8
5.2
74.31
4.0
117.8
4.7
51.66
1.4
cio,
6.6
0.52
97.16
6.4
313.4
15.3
12.25
0.6
3.69
0.13
35.42
1.1
174.9
4.6
179.3
5.3
126.4
4.7
183.9
5.3
150.6
3.16
Br03
6.38
0.38
36.67
2.2
1525
8.20
0.59
0.025
0.30
0.009
0.8
0.02
30.0
0.9
85.17
2.3
42.86
1.5
58.43
1.8
1.3
0.03
IOS
7.62
0.35
8.33
0.4
80.43
3.84
0.004
0.0314
0.059
0.0216
0.05
0.001
0.25
0.0,57
0.25
0.007
6.87
0.26
0.83
0.02
0.002
0.043
OH
142.9
18
116.4
21.
12.04
5.0
0.01
0.001
40.04
1.76
3.7
0.22
0.77
0.063
0.17
0.02
0.001
0.032
0.0.,5
0.045
0.01
0.034
SO4
11.11
0.62
16.83
1.15
35.64
2.8
0.55
0.020
4.74
0.09
0.0,23
0.0410
0.011
0.036
0.20
0.015
35.43
2.8
53.12
3.1
0.0041
0.0313
CrO4
63.1
2.7
61.21
3.30
111.6
6.5
0.0025
0.0.,15
0.006
0.031
0.0,38
0.0415
0.12
0.006
0.4
0.03
73.0
4.3
0.042
0.085
C204
30.27
1.6
3.34
0.24
7.22
0.69
0.0035
0.032
1.48
0.030
0.0086
0.0338
0.0046
0.0326
0.0,56
0.0443
0.03
0.0027
0.036
0.044
0.0315
0.055
C03
108.0
5.9
19.39
1.8
1.3
0.17
0.003
0.031
4.95
0.10
0.0023
0.0311
0.0011
0.047
0.0013
0.0313
0.1
0.01
0.004?
0.033?
0.031
0.043
The following are the solubilities (number of grams in 100 c.c. water)
at 18° of two additional insoluble substances and of three acid salts.
Mercurous chloride, 0.032
(molar sol'ty, 0.041)
Mercuric iodide, 0.044
(molar sol'ty, 0.051)
Sodium bicarbonate 9.6
Potassium bicarbonate 26.1
Potassium bisulphate 50.0
THE BASE-FORMING ELEMENTS
545
It will be seen that some compounds, like zinc chloride and barium
iodide, are exceedingly soluble ; that others, like potassium chloride and
barium chlorate, are of medium solubility ; that still others, like calcium
hydroxide and calcium sulphate, are sparingly soluble ; and, finally, that
some, like calcium oxalate (CaC204) and barium chromate, are almost
insoluble. The reader should note the fact, however, that the differ-
ences in solubility even amongst the insoluble salts are as great as
amongst the soluble ones.
Hydrated Forms of Salts Commonly Used. — In the table
given below, the figures refer to the number of molecules of water in
the hydrates which are deposited by aqueous solutions of the salts in
the neighborhood of 18°. The letter h means that the compound is
stable when heated, the letter* a that it is not affected by the air, the
letter d that the salt is deliquescent, and the letter e that the hydrate
loses water spontaneously in an open vessel, i.e., is efflorescent.
Composition of Hydrates of Salts.
K
Na
Li
Ag
Ba
Sr
Ga
Mg
Zn
Cd
Cu
Pb
Cl .
Oh
Oh
Oh
Oh
2a
6e
6d
6d
l^d.
2}e
2d
OA
Br .
Oh
Oh
Oh
Oh
2a
6d
6d
Qd
2d
4e
4e
OA
I .
OA
Oh
Oh
Oh
2d
6d
Od
Sd
Od
Oa
OA
N03
Oa
Oa
Od
Oa
Oa
4e
4d
Qd
Qd
4d
6d
Oa
C103
Oa
Oa
Od
Oa
la
5d
2d
Qd
Qd
2d
6d
la
Br03
Oa
Oa
Od
Oa
la
la
la
Qe
6a
2a
6a
la
I0j
Oa
3e
Od
Oa
la
6e
6e
4a
2a
Oa
la
Oa
C2H302
Od
3d
2d
Oa
la
la
2d
4d
3a
3d
la
3a
S04.
Oh
lOe
Oh
Oa
Oh
OA
2a
le
le
2fa
5a
OA
CrO4
Oh
lOe
2a
Oa
Oh
OA
la
le
OA
CA
la
Oa
Oa
Oa
la
Oa
la
2a
2a
3a
la
Oa
C03
lid
lOe
Oa
Oe
OA
OA
Oa
3e
Oa
Oa
Oa
Isomorphism. — Substances which crystallize in one of the forms
belonging to the regular system (p. 138) must necessarily have identi-
cal crystalline shapes. Thus, crystallized specimens of sodium chlo-
ride and of lead sulphide (galena) in their natural shapes are cubical.
The forms found in other systems, however, are capable of assuming
an infinite diversity of shapes. The relative lengthening or shorten-
ing in one direction, shown by the square prismatic and hexag-
onal forms (p. 138), for example, makes it possible for each separate
substance to adopt proportions which are more or less different
546 INORGANIC CHEMISTRY
from those of every other substance. Each substance, not belonging
to the regular system, does, in fact, crystallize invariably in forms
based upon its own fundamental proportions, and differs therefore in
its angles from all other such substances in a way that is clearly
recognizable by refined measurement.
Now it is found that substances which are chemically somewhat
similar (see below) freqiiently crystallize in the same system and show
proportions which are almost, although not quite, identical. Further-
more, such substances, when the approach to identity in angles is not
accidental, can take part in the construction of one and the same
crystal. A crystal of one such substance placed in a solution of the
other will continue to grow, and in doing so will follow the pattern
already set, and simply increase in dimensions by accretion of the new
material. When a solution containing two such substances deposits
crystals, the structures are, not some of them of one material and some
of the other, but are all made up of both in a ratio determined by the
relative amounts of the substances in the solution.* Substances
related in these two ways, that is, baring, when separate, crystalline
forms which are closely alike and being capable of forming homoge-
neous crystals containing varying proportions of the two ingredients,
are called isomorphous substances (Gk. «ros, equal ; fj.op(f>^ form).
Thus, potassium permanganate KMn04 and potassium perchlorate
KC104 crystallize in the rhombic system, forming crystals with very
similar angles (Fig. 53, p. 139), and when a solution containing both is
allowed to evaporate there is formed but one set of crystals made up of
both substances. Similarly, potassium iodide and ammonium iodide
crystallize in cubes of the regular system, and, since all cubes are
alike, necessarily show absolutely identical angles. In addition to this,
however, they crystallize together from a solution containing both
salts. Other pairs from amongst substances belonging to the regular
system would not do this. The two salts are therefore isomorphous.
In the course of our study of the compounds of the metals we
shall have occasion to note many examples of isomorphism. Thus the
heptahydrates of the sulphates of many of the bivalent metals, such
as ZnSO4,7H20, FeS04,7H2O, MgS04,7IT2O, etc., belong to the rhombic
system, and form an isomorphous set of substances known as the
vitriols (q.v.).
* In general, two substances which are absolutely unrelated may be deposited
simultaneously from mixed aqueous solutions, but some of the crystals are pure
specimens of one substance, and the rest are pure specimens of the other.
THE BASE-FORMING ELEMENTS 547
The alums (q.v.) also constitute an important set, and crystallize,
separately and together, in the regular system (see, also, under Zinc
sulphate). Amongst minerals, lead sulphide (PbS) and silver sulphide
(Ag2S) form a common isomorphous pair, and nearly all natural speci-
mens of galena (<?.#.) contain at least a little silver sulphide.
The chemical significance of isomorphism was at first exaggerated.
Thus the elements magnesium and iron are not especially similar in
their chemical relations, excepting that both are bivalent ; yet they
form several pairs of compounds which, like the sulphates (above), are
isomorphous. Still, in practical chemical work a knowledge of the
relations of a substance in respect to isomorphism is indispensable.
It enables us to predict the probable impurities in a homogeneous-
looking material, for non-isomorphous substances would have given a
heterogeneous mixture with it. It assists us in separating and puri-
fying chemical substances, for non-isomorphous substances can be sep-
arated by recrystallization from water (p. 273), or by washing with
water or some other solvent (p. 276), at a temperature at which the
solubilities of the substances are different (see Potassium nitrate).
Isomorphous substances, however, can be separated only by conversion
into some other form of combination in which the property is lack-
ing. Thus, silver sulphide cannot be separated from lead sulphide by
Pattinsou's process (q.v.), and so the mixed metals, to the separation
of which the process is applicable, must first be secured by reduction.
Some chemists regard isomorphous mixtures as solid solutions.
Exercises. — 1. Compare the electrical conductivities of normal
sodium hydroxide and normal acetic acid with the conductivity of
copper. What length of copper wire will present the same resistance
as 1 cm. of each of these solutions when the cross-sections are alike ?
2. What do we mean by saying that an oxide is strongly or feebly
basic, or that it is acidic (p. 541) ?
3. What is meant by the same terms when applied to an hydroxide ?
4. Compare the molar solubilities at 18° (a) of the halides of silver
and (b) of the carbonates and (c) oxalates of the metals of the alkaline
earths, noting the relation between solubility and atomic weight.
5. What is the molar concentration of chloridion (cf. p. 149) in
saturated solutions of silver chloride and lead chloride at 18°, assum-
ing complete ionization in these very dilute solutions ?
6. How does the behavior of complex acids, like chloroplatinic
acid HaPtCl6, differ from that of acid salts ?
CHAPTER XXXIII
THE METALS OF THE ALKALIES: POTASSIUM AND
AMMONIUM
The Metals of the Alkalies. — The metals of this family form a
homogeneous group, and there is a very general similarity between the
properties of the corresponding compounds. Some of the physical
properties of the elements themselves can be presented best in tabular
form.
AT. WT.
SP. GK.
M.-P.
B.-P.
Lithium. . . "
7.0
0.59
186°
above red heat.
Sodium .
23.0
0.97
95.6°
742°
Potassium
39.1
0.87
62.5°
667°
Rubidium
85.5
1.52
38.5°
Caesium .
132.9
1.88
26.5°
270°
It will be seen that the specific gravities of the elements increase with
rising atomic weight, while the melting-points and boiling-points fall
(cf. p. 410). A table including all the physical properties, both of the
elements and their compounds, would show similar characteristics in a
general way, with here and there noticeable irregularities such as that
shown by the specific gravity of sodium.
The Chemical Relations of the Metallic Elements of the Alka-
lies. — The metals which are chemically most active are included in
this group, and the activity increases with rising atomic weight, caesium
being the most active positive element of all. A freshly cut surface of
any of these metals tarnishes by oxidation as soon as it is exposed to the
air. Indeed, there is scarcely time to see the metallic appearance in
the case of potassium and the metals following it. All of these metals
decompose water violently (cf. p. 97), liberating hydrogen. The
hydroxides which are formed by this action are exceedingly active
bases, that is to say, they give a relatively large concentration of
hydroxidion in solutions of a given molecular concentration (p. 349).
548
POTASSIUM AND AMMONIUM 549
Lithium hydroxide is the least active. In the dry form these hydrox-
ides are not decomposed by heating, while the hydroxides of all other
metals lose water more or less easily. All these metals seem to com-
bine with hydrogen, lithium giving the most stable compound. The
hydrides, however, unlike those of many of the non-metals, are not
ionogens, and consequently do not give acids when dissolved in water.
In all their compounds the metals of the alkalies are univalent.
The family may be subdivided into two minor groups. The com-
pounds of potassium, rubidium, and caesium resemble one another
closely, while those of sodium and lithium are sometimes largely diver-
gent in physical properties. Thus, the chlorides of the potassium set, not
only crystallize in cubes, but can form mixed crystals with one another
in all proportions. They are isomorphous (cf. p. 545). The same is
true of the bromides and of the iodides. Sodium chloride, although
crystallizing in cubes likewise, does not form mixed crystals with the
chlorides of the potassium set. In the case of lithium, the hydroxide
is not nearly so soluble as are the hydroxides of the other metals, and
the metal gives also an insoluble carbonate and phosphate, in which
respect it resembles magnesium and differs from all the other mem-
bers of the present group.
The compounds of ammonium will be discussed in connection with
those of potassium, to which they present the greatest resemblance.
The solubilities are often decisive factors in connection with the
preparation and use of salts. The reader will find most of these in the
table on p. 544, or the diagram on p. 157, and, as a rule, the values will
not be repeated in the descriptive paragraphs.
POTASSIUM.
Occurrence. — Silicates containing potassium, such as feldspar and
mica (p. 525), are constant constituents of volcanic rocks, and from the
weathering of these rocks, and of the detritus formed from them which
constitutes a large part of the soil, the potassium used by plants is
obtained. These minerals are not used commercially as sources of
potassium compounds. The salt deposits (see below) contain potassium
chloride, alone (sylvite) and in combination with other salts, and most
of the compounds of potassium are manufactured from this material.
Part of our potassium nitrate, however, is purified Bengal saltpeter
(p. 438). Potassium sulphate occurs also in the salt layers, and is used
directly as a fertilizer.
550 INORGANIC CHEMISTRY
Preparation. — Potassium was first made by Davy (1807) by
bringing the wires from a battery in contact with a piece of moist
potassium hydroxide. Globules of the metal appeared at the negative
wire. This process has just come into use commercially, molten potas-
sium chloride being the substance decomposed. The process which,
during the intervening years, furnished potassium, was the heating of
potassium carbonate with finely divided carbon in small retorts :
K2C03 + 2C -> 2K + SCO.
The vapor of the metal tends to combine with the carbon monoxide,
forming an explosive compound K6C606, and the yield is thus reduced.
Castner's improved process involves the heating of potassium hydroxide
with a spongy mass which is essentially a carbide of iron (CFe2). The
latter is made by heating together pitch and iron filings. !No carbon
monoxide is produced :
6KOH + 2C -> 2K2C03 + 3H2 + 2K.
Physical Properties. — Potassium is a silver-white metal which
melts at 62.5°. It boils at 667°, giving a greenish vapor. The metal
and its compounds confer a violet tint upon the Bunsen flame, and the
spectrum (q.v.) shows characteristic lines.
Chemical Properties. — The density of the vapor shows the
molecular weight of potassium to be about 40, so that the vapor is a mon-
atomic gas. The element unites violently with the halogens, sulphur,
and oxygen. In consequence of the latter fact it is usually kept under
petroleum, an oil which neither contains oxygen itself, nor dissolves a
sufficient amount of oxygen from the air to permit much oxidation of
the potassium to take place.
The Hydride. — When hydrogen is passed over potassium heated
to 360°, a hydride is formed. By washing the solid product with
liquefied, dry ammonia the excess of potassium is removed and white
crystals remain. These have the composition KH. On account of the
ease with which it decomposes, the substance behaves much like
potassium itself. When thrown into water, for example, it gives
potassium hydroxide, and the hydrogen is liberated.
Potassium Chloride. — Sea-water and the waters of salt lakes
contain a relatively small proportion of potassium compounds. During
POTASSIUM AND AMMONIUM 551
the evaporation of such waters, however, the potassium compounds
tend to accumulate in the mother-liquor while sodium chloride is being
deposited. Thus, when the Salt Lake in Utah shall have finally dried
up, the upper (last to be formed) part of the bed of salts which it will
leave behind will contain layers rich in compounds of potassium. This
condition is realized in geological deposits which have been formed in
the same way. Thus, at Stassfurt, near Magdeburg, there is a thickness
of more than a thousand meters of common salt, more or less mixed
with and intersected by layers of sedimentary deposits. Above this are
25 - 30 meters of salt layers in which the potassium salts are chiefly
found, while over all are several hundred meters of sandstone. For-
merly the upper layers were simply stripped off and rejected. Now,
however, the revenue obtained from the products of these layers is
many times greater than that coming from the rock salt below.
The chief sources of the potassium chloride found in the salt beds
are sylvite (KC1) and carnallite (KC1, MgCl2, 6H20). The latter is
heated with a small amount of water, or with a mother-liquor obtained
from a previous operation and containing sodium and magnesium
chlorides. The magnesium sulphate which it contains as an impurity
remains undissolved. From the clear liquid, when it cools, potassium
chloride is deposited first and then carnallite. The former is taken out
and purified, and the latter goes through the process again. This
potassium chloride is the source from which most of our potassium
hydroxide and potassium carbonate, as well as salts of minor commer-
cial importance, are made. It is a white substance crystallizing in
cubes, melting at about 750°, and slightly volatile at high temperatures.
Potassium Iodide. — When iodine is heated in a strong solution
of potassium hydroxide, the iodate and iodide are both formed (p. 277) :
6KOH + 3I2 -> 5KI + KIO3 + 3H20.
The dry residue from evaporation is heated with powdered carbon to
reduce the iodate, and all the iodide can then be purified by recrystal-
lization. Another method of preparation consists in rubbing together
iodine and iron filings under water. The soluble ferrous iodide (FeI2)
thus formed is then treated with additional iodine and gives a sub-
stance Fe3I8, intermediate in composition between ferrous and ferric
iodides. This is also soluble. When potassium carbonate is added to
the solution, a hydrated magnetic oxide of iron is precipitated, carbon
552 INORGANIC CHEMISTRY
dioxide escapes, and evaporation of the clear solution gives potassium
iodide :
Fe8I8 + 4K2C08 + 4H20 -> SKI + Fe8(OH)8 + 4COa.
The salt forms large, somewhat opaque cubes (m.-p. 623°). It is
used in medicine and for producing precipitates of silver iodide in
photography. In the laboratory it is used whenever an iodide is
required, for example, when experiments with iodidion are to be made
The Bromide and Fluorides. — Potassium bromide may be
made in either of the ways used for the iodide. It crystallizes in
cubes. It is used in medicine and for precipitating silver bromide in
making photographic plates (q.v.). In the laboratory it is always
employed when a bromide is needed as a source of bromidion.
The fluoride of potassium K2F2 may be obtained by treating the
carbonate or hydroxide with hydrofluoric acid. It is a deliquescent,
white salt. When treated with an equi-molecular quantity of hydro-
fluoric acid it forms potassium hydrogen fluoride KHF2, a white salt
which is also very soluble. This acid salt is used in the preparation
of pure hydrofluoric acid, since the latter is liberated from it as a vapor
at a high temperature.
Potassium chloride is the least soluble of the halides of potassium,
the bromide, iodide, and fluoride coming next in that order. The posi-
tion of the fluoride as the most soluble, when we should expect it to be
the least soluble (p. 244), shows that this compound is somewhat
exceptional. It is also slightly hydrolyzed by water, as if it were a
salt of a dibasic acid (cf. p. 344). These facts, together with the
existence of the acid fluoride, lead us to assign to it the formula K2F2.
Other acid fluorides of the formulae KH2F8 and KH8F4 have likewise
been made. Since potassium, hydrogen, and fluorine are always univa-
lent, and no ordinary valence is thus available for holding together
groupings more complex than KF and HF, we may regard all these
four fluorides of potassium as molecular compounds (p. 443).
Potassium Hydroxide* — This compound, known also as caustic
potash and sometimes as potassium hydrate (p. 120), was formerly
made entirely by boiling potassium carbonate with calcium hydroxide
suspended in water (milk of lime) :
K2C08 + | Ca(OH)2 <=> CaC08 J + 2KOH.
POTASSIUM AND AMMONIUM 553
The operation is conducted in iron vessels, because porcelain, being
composed of silicates, interacts with solutions of bases. The action is,
in theory, precisely similar to that of sulphuric acid upon barium diox-
ide (cf. p. 585). The potassium carbonate corresponds to the acid,
being completely dissolved from the beginning, and the calcium hydrox-
ide to the dioxide, since its relative insolubility enables the water to
take up fresh portions into solution only when the part dissolved has
already undergone chemical change. The calcium carbonate which is
precipitated is much less soluble than the hydroxide, and hence the
action goes forward. The action as a whole is reversible, for a
reason which will be explained later (see Ionic equilibrium in
Chap, xxxiv), and consequently such an amount of water is employed
that the solution at no time contains more than about ten per cent of
potassium hydroxide (sp. gr. 1.1). The conclusion of the action is
recognized when a clear sample of the liquid no longer effervesces
on addition of a dilute acid, and is therefore free from potassium
carbonate.
Recently much potassium hydroxide has been manufactured by
electrolytic processes. When a solution of potassium chloride is elec-
trolyzed, chlorine is liberated at the anode, and hydrogen and potassium
hydroxide at the cathode. The necessity of keeping those two sets of
products apart, since by their interaction potassium hypochlorite and
potassium chloride would be formed (cf. p. 266), has made the devising
of suitable apparatus extremely difficult. In one type of apparatus a
partition of asbestos cloth, especially prepared to resist the disinte-
grating effects of the alkali and the chlorine, divides the cell into two
parts. In some cases this is placed vertically, and in others horizon-
tally. In the latter case the anode is on the upper side of the partition,
in order that the chlorine as it is liberated may ascend to the surface
without stirring up the liquid or having occasion to pass near the par-
tition. In all cases the anode is made of graphite, since this substance
is less easily attacked by chlorine than is any other, and the cathode is
made of iron, a metal which best resists the action of alkalies. The
chlorine is used for making bleaching powder. Pure brine flows in con-
tinuously at one point, and a solution of the hydroxide containing much
undecomposed chloride flows out at another.
The same process is applied to sodium chloride, and, in some fac-
tories, the apparatus is used solely for making bleaching materials, and
most or all of the alkali is thrown away. This is on account of the
cheapness of the caustic soda made by non-electrolytic processes, and
554
INORGANIC CHEMISTRY
the expense involved in concentrating the rather dilute solution
obtained electrolytically.
The Castner-Kellner apparatus (Fig. 97) employs a different princi-
ple very ingeniously for the separation of the products. The two end
compartments are filled with brine and contain the graphite anodes.
The central compartment contains potassium hydroxide solution and
the iron cathode. The positive current enters by the anodes, and the
chlorine is therefore attracted to and liberated upon the graphite.
After rising through the liquid it is collected for the manufacture
of liquefied chlorine or of bleaching powder. The ions of potassium
or of sodium, as the case may be, are discharged upon a layer of mer-
FlG. 97.
cury which covers .the whole floor of the box, and the free metal
dissolves in the mercury, forming an amalgam (p. 532). The layer of
mercury extends beneath the partitions, and a slight rocking motion
given to the cell by the cam ((7) causes the amalgam to flow below
the partition into the central compartment. Here the sodium leaves
the mercury in the form of sodium ions and is attracted by the
cathode. Upon this, hydrogen from the water is discharged, and the
residual hydroxidion, together with the metallions, constitutes potas-
sium or sodium hydroxide. A slow influx of salt solution at one
point and overflow of the alkaline solution in the central cell at
another, is maintained. The overflowing liquid contains 20 per cent
of the alkali. Since in this form of the apparatus there is no unde-
composed chloride present in the part of the solution which contains
the hydroxide, simple evaporation to dryness furnishes the solid
alkali.
Potassium hydroxide is exceedingly soluble in water, and conse-
quently, instead of being crystallized from solution, the molten residue
POTASSIUM AND AMMONIUM 555
from evaporation is cast in sticks. When, for chemical purposes, the
hydroxide is required free from potassium carbonate and other impu-
rities, it is dissolved in alcohol, in which the other substances are not
soluble. Evaporation of this solution gives pure caustic potash. The
hydroxide, in consequence of the very low vapor tension of its solution
(cf. p. 162), is highly deliquescent. It also absorbs carbon dioxide from
the air, giving potassium carbonate. This salt is itself deliquescent,
and consequently a syrupy solution of the carbonate is the final result
of weathering. Solutions of the hydroxide have an exceedingly corro-
sive action upon the flesh, decomposing it into a slimy mass by hydro-
lyzing the albuminous and other substances. In solution, the base is
highly ionized, furnishing a high concentration of hydroxidion. Its
aqueous solution is therefore used with salts of other metals for precipi-
tation of less soluble bases. Commercially it is chiefly employed in the
making of soft soap.
The Oxides. — The simple oxide K20 may be made by heating
potassium nitrate or nitrite with potassium in a vessel from which air
is excluded : KN03 + 5K — » 3K20 + N. It interacts violently with
water, giving the hydroxide. When exposed to the air it unites spon-
taneously with oxygen, and K204 is formed.
When the metal burns in oxygen, K204, a yellow solid is the prod-
uct. This substance interacts violently with water, giving potassium
hydroxide, and the excess of oxygen is liberated. With perfectly dry
oxygen, potassium does not unite, even when it is heated strongly.
Potassium Chlorate* — The preparation of this salt (KC10g), by
interaction of potassium chloride with calcium chlorate, has already
been described (p. 273). It is also made by electrolysis of potassium
chloride solution, the potassium hydroxide and chlorine which are
liberated being precisely the materials required. All that is necessary
is to use a warm, concentrated solution and to provide for the mixing
of the materials generated at the electrodes. The salt crystallizes out
when the solution cools.
Potassium chlorate crystallizes in monoclinic plates. It melts at
about 359°, and at a temperature slightly above this the visible libera-
tion of oxygen begins. Since heat is given out by the decomposition,
the action may be almost explosive if large amounts of the material
are employed. This decomposition doubtless takes place at all tem-
peratures (p. 73), but below the melting-point the speed is so slight
55G INORGANIC CHEMISTRY
that the phenomenon is not perceived. On account of the ease with
which its oxygen is liberated, the salt is employed in making fire-
works and as a component, along with antimony trisulphide, of the
heads of Swedish matches. With acids it is used as an oxidizing agent
on account of the chloric acid which is set free (p. 272). It is also em-
ployed in medicine.
Potassium perchlorate KC104, formed by the heating of the chlorate
(p. 275), gives white crystals belonging to the rhombic system. Com-
pared with the chlorate, on account of the greater difficulty in liberat-
ing its oxygen by heat, it finds little practical application.
The Bromate and lodate. — These are the most familiar salts of
their respective acids. The mode of their preparation has already been
described (p. 277). Potassium iodate may be made also very conven-
iently by melting together potassium chlorate and potassium iodide at
a low temperature. The iodate is much less soluble than the chloride,
and the mixture may be separated by crystallization from water.
Potassium Nitrate. — The formation of this salt in nature and
its mode of extraction and purification have already been described
(p. 438). This source of supply proved insufficient, for the first time,
during the Crimean war (1852-55), and a method of manufacture from
Chili saltpeter (sodium nitrate), which is a much cheaper substance,
was introduced. Sodium nitrate and potassium chloride are heated
with very little water, and the sodium chloride produced by the action,
which is a reversible one, is by far the least soluble of the four salts. On
the other hand, at this temperature, the potassium nitrate is by far the
most soluble. Hence the hot liquid drained from the crystals contains
the required salt, and most of the sodium chloride is in the form of a
precipitate. If the solubility curve of potassium nitrate (p. 157) is ex-
amined, it will be seen that this salt is but slightly soluble in cold water,
and hence most of it is deposited when the solution cools. The crys-
tals are mixed with little sodium chloride, for, as the curve shows, com-
mon salt is little less soluble at 10° than it is at 100°.
Potassium nitrate gives long prisms belonging to the rhombic sys-
tem (Fig. 98). It melts at about 340°, and when more strongly heated
gives off oxygen, leaving potassium nitrite (p. 449). Although it does
not form a hydrate, the crystals inclose small portions of the mother-
liquor, and consequently contain both water and impurities. When
heated, the crystals fly to pieces explosively (decrepitate), on account of
POTASSIUM AND AMMONIUM
557
the vaporization of this water. All substances which form large crystals
and do not melt when warmed, behave in the same way and for the same
reason. In consequence of this, the purest salt is made by violent stir-
ring of the solution during the operation of crystallization, the result
being the formation of a crystal-meal.
Potassium nitrate is used chiefly in the manufacture of gunpowder,
whi^h contains 75 per cent of the highly purified salt. The other com-
ponents are 10 per cent of sulphur, 14 per cent of
charcoal, and about 1 per cent of water. The ingre-
dients are intimately mixed in the form of paste,
and the material when dry is broken up and sifted,
grains of different sizes being used for different pur-
poses. The chemical action which takes place when
gunpowder is fired in an open space probably results
chiefly in the formation of potassium sulphide, car-
bon dioxide, and nitrogen :
2KN03 + 30 + S -» K2S + 3C02 + Na.
The explosion occurring in firearms follows a much
more complex course, and half of the solid product
is said to be potassium carbonate. The pressure, at
the temperature of the explosion, if the gases could
be confined within the volume originally occupied by the gunpowder,
would reach about forty-four tons per square inch. In recent years
common gunpowder has been displaced largely by smokeless powder,
of which substances related to gun-cotton (p. 441) are the chief compo-
nents.
Paper saturated with potassium nitrate solution and dried, is known
as touch-paper. The salt interferes with the access of air to the cellu-
lose, and the oxygen for the combustion is obtained from the nitrate.
The product consequently does not blaze and cannot easily be blown
out.
Formerly nitric acid was all made from potassium nitrate, but now
the cheaper sodium nitrate is employed exclusively.
Potassiiim Carbonate. — This salt is manufactured from potas-
sium sulphate, from the Stassfurt deposits, by the Le Blanc soda
process (q.v.}. A certain amount is also obtained from the fatty
material, known as suint, which forms about 50 per cent of the weight
of sheep's wool, and is separated from the latter by washing. When
FIG. 98.
558 INORGANIC CHEMISTRY
this material, which contains the potassium salt of sudoric acid in
large proportions, is calcined, potassium carbonate remains, and is ex-
tracted from the ash with water. Some plants, like the sugar-beet,
take up exceptional quantities of potassium salts from the soil. The
molasses remaining from the crystallization of beet-sugar (p. 500)
is mixed with yeast and fermented. After the alcohol has been dis-
tilled off, the liquid, containing organic salts of potassium in solution,
is evaporated, and the residue is ignited. In some districts potassium
carbonate is still extracted from wood-ashes.
This salt is usually sold in the form of an anhydrous powder (m.-p.
over 1000°). When crystallized from water it gives a hydrate 2K2CO3,
3H20. It is extremely deliquescent. Its aqueous solution has a
marked alkaline reaction. The hydrolysis of the salt by the water is
exactly analogous to that of sodium sulphide (p. 375), although not so
extensive. The more elaborate scheme given in that connection may
be put in simpler form to show that the action consists essentially in
the formation of the ion HCO/, by union of the ion C03" with the
hydrion of the water. This takes place because the ionization of the
former ion is small enough to be commensurable with that of water it-
self : C03" + H* + OH' -> HCO/ -f- OH'. The commercial name of the
substance is pearl ash. It is used in making soft soap and hard glass.
It is also employed, by interaction with acids, in making salts of
potassium.
When a concentrated solution of the salt is electrolyzed in such a
way that the anode, towards which the KC03' ions travel, consists of
a thin platinum wire, the crowding together of the discharged material
results in the formation of the percarbonate (cf. p. 397) :
2KC08 + 20->K2C20..
The operation must be conducted between — 15° and 0°. When the
solution in the porous cell surrounding the anode is evaporated, the
product is obtained as an amorphous bluish-white powder. The sub-
stance liberates oxygen when heated, and in other respects behaves like
the persulphates. When it is treated with a dilute acid, a solution
containing hydrogen peroxide is formed. The compound is therefore
a mixed anhydride (p. 397) of hydrogen peroxide and potassium bi-
carbonate.
Potassium Cyanide. — Formerly this compound was made by
heating potassium carbonate with nitrogenous animal matter. So
POTASSIUM AND AMMONIUM 559
many other substances' were formed at the same time, however, that
the required product, which is very soluble, was difficult to isolate in
a state of purity. It is now made by heating together potassium ferro-
cyanide (q.v.} and potassium carbonate. The ferrocyanide acts as if it
were a mixture of potassium cyanide and ferrous cyanide : K4Fe(CN)6
— » 4KCN -+- Fe(CN")2. The latter, by interaction with the potassium
carbonate, would give potassium cyanide and ferrous carbonate, but this
in turn, is decomposed by heat into ferrous oxide, which is insoluble,
and carbon dioxide :
K4Fe(CN)6 + K2CO3 -> 6KCN + FeO + C02.
When the residue is extracted with water, only the potassium cyanide
dissolves, and it is easily crystallized in pure form from the solution.
Very interesting is the formation of potassium cyanide in the blast
furnace (q-v.). Carbon and nitrogen unite at a very high temperature
to form cyanogen (p. 260), and a sufficient amount of potassium is
found in the materials to complete the production of the salt.
Potassium cyanide crystallizes in cubes. It is extremely soluble in
water, and is therefore deliquescent. Its poisonous qualities are equal
to those of hydrocyanic acid. The acid is so feeble as to be liberated
even by the carbon dioxide of the air, and hence this salt always has a
distinct odor of hydrocyanic acid. Potassium cyanide has a great
tendency to form complex compounds with cyanides of other metals
(cf. p. 536). Complex compounds of this kind are used in the gal-
vanic deposition of silver and gold in commercial electroplating.
Large amounts of the cyanide are also used in extracting gold (q.v.} from
its ores, particularly in the Transvaal colony. The tendency to form
complex compounds is doubtless connected with the fact that the
cyanides are unsaturated compounds in which the carbon has two
free valences : K-N = C (p. 507).
Potassium cyanate KCNO is made by heating potassium cyanide
in the air, or, still better, with some easily decomposed oxide (p. 507).
It is a white, easily soluble, crystalline salt.
Potassium thiocyanate KCNS may be obtained by melting potas-
sium cyanide with sulphur (cf. p. 508). It is a white, deliquescent salt
which finds some applications in chemical analysis.
The Sulphate and Bisulphate. — The sulphate of potassium
is a constituent of several double salts found in the Stassfurt de-
posits. It is extracted from schoenite MgS04,K2S04, 6H20 and kainite
560 INORGANIC CHEMISTRY
MgS04, MgCl2, K2SO4, 6H20. The former is treated with potassium
chloride and comparatively little water, whereupon the relatively in-
soluble potassium sulphate crystallizes out, and the magnesium chlo-
ride remains in the mother-liquor. The crystals belong to the rhombic
system, contain no water of crystallization, and melt at 1066°. This salt
is employed in large quantities in making potassium carbonate by the Le
Blanc process and in preparing alum (q.v.). It is also much used as a
fertilizer. Since plants take up solutions through their cell walls, they
can absorb soluble compounds only. They are, therefore, dependent,
for the potassium compounds which they require, upon the weathering
out of soluble potassium compounds from the insoluble potassium
silicates contained in the soil. The weathering takes place too slowly
to furnish a sufficient supply for many crops, particularly that of the
sugar-beet. Hence potassium sulphate is mixed directly with the soil.
The mineral kainite itself is used for the same purpose.
Potassium hydrogen sulphate (bisulphate) KHS04 is made by
the action of sulphuric acid upon potassium sulphate : K,S04 +
H2S04— »2KHS04. It crystallizes from water, in which it is very
soluble, in tabular crystals. When heated to about 200° it melts, and
the elements of water are eliminated, the pyrosulphate remaining :
2KHS04 -4 H20 + K2S2O7. The latter, when still further heated,
yields sulphur trioxide and potassium sulphate. The bisulphate is
used in analysis for the purpose of decomposing oxides and silicates
and converting them into sulphates. The substance is more efficient
than sulphuric acid for this purpose, because the latter cannot be
heated above 330°, while the liberation of the active sulphur trioxide
from this salt takes place at a bright-red heat. The aqueous solu-
tion of the bisulphate is strongly acid on account of the considerable
ionization of the hydrosulphanion.
Sulphides of Potassium. — By the treatment of a solution of
potassium hydroxide with excess of hydrogen sulphide, a solution of
potassium hydrogen sulphide is obtained. Evaporation of the solution
gives a deliquescent solid hydrate 2KHS, H2O. When the solution,
before evaporation, is treated with an equivalent amount of potassium
hydroxide, and the water is driven off, the sulphide K2S remains
behind (cf. p. 375) :
KHS + KOH +± K2S + H2O.
With proper care, the very soluble hydrate K2S, 5H20 may be obtained.
Considerable amounts of sulphur can be dissolved in solutions of either
POTASSIUM AND AMMONIUM 561
of thsse sulphides. By evaporation of the resulting yellow liquids, vari-
ous polysulphides have been obtained. To some of these have been
ascribed the formulae K2S8, K4S7, K2S4, K4S9, and K2S5 (cf. p. 376).
Similar substances are produced, as a result of the liberation and
recombination of sulphur, when the solutions are exposed to the
oxidizing action of the air :
2KHS + 02 -> 2KOH + 28.
In most respects the corresponding compounds of potassium and
sodium are similar in their physical properties and chemical action.
Since, however, the latter are almost uniformly less expensive, they
find much wider application. In a few cases, however, the potassium
salt is more generally used. Thus, potassium chlorate and potassium
iodide are much less soluble than the corresponding sodium compounds,
and it is consequently possible in each of these two cases to separate
by crystallization, and to purify the potassium salt with greater ease.
Properties of Kalian: Analytical Reactions. — The positive
ionic material of the potassium • salts is a colorless substance. It
unites with all negative ions, and most of the resulting compounds are
fairly soluble. For its recognition we add solutions containing those
ions which give with it the least soluble salts. Thus, with chloroplat-
inic acid H2PtCl6 it gives a yellow precipitate of potassium chloro-
platinate K2PtCl6. Since nearly one part of this salt dissolves in 100
parts of water, the test is far from being a delicate one. The solubility
in alcohol is much smaller, and consequently the precipitate may fre-
quently be obtained from a dilute solution by adding more than an
equal volume of alcohol. Picric acid (p. 441) gives potassium picrate
KC6H2(IS[02)30, which is much less soluble in water (0.4 parts in 100
at 15°). Perchloric acid and hydrofluosilicic acid likewise give some-
what insoluble salts of potassium. Potassium hydrogen tartrate
KHC4H4O6 is precipitated by the addition of tartaric acid to a suffi-
ciently concentrated solution of a potassium salt. The normal tartrate
K2C4H406 is much more soluble. It may be obtained by treating the
precipitate with a solution of potassium carbonate or potassium hy-
droxide. Addition of an acid to this solution causes reprecipitatioii
of the bitartrate.
The Spectroscope. — A much more delicate test for the recognition
of a potassium compound consists in the examination by means of the
562 INORGANIC CHEMISTRY
spectroscope of the light given out by a Bunsen flame, in which a
little of the salt is held upon a platinum wire. When the amount of
potassium is considerable, and no other substance which would likewise
color the flame is present to mask the effect, the violet tint is recog-
nizable by the eye. In general, however, the light must be analyzed.
White light is composed of vibrations of every wave-length within a
certain range. If the light is made up of one or more wave-lengths
only, it appears to the eye to be colored. Now, when a narrow bundle
of rays of white light, coming through a slit, falls upon a three-sided
prism standing with its edges parallel to the slit, the rays of various
wave-length are retarded to different extents as they pass through the
glass, and in consequence are bent from their paths by varying amounts.
Fig. 99 shows a horizontal section through the slit ($) and prism, in
which the width of the slit and
of the beam of light are exagger-
Vioiet ated. The light emerging at the
Blue other side of the prism consists,
Yellow therefore, of a series of images
Bed of the slit arranged side by side.
The red light is least refracted,
and the red images of the slit,
therefore, are most nearly in the
same straight line with the ori-
ginal beam. The yellow, green,
blue, and violet images are displaced more and more from this direction,
and the resulting colored band is called a spectrum. The whole
series of images of the slit may be received upon a screen, or directly
upon an eye looking towards the prism. Now, when the light comes
from the vapor of potassium heated in a Bunsen flame, there are pro-
duced, not thousands of images of the slit, representing as many differ-
ent wave-lengths of light, but only two images, one red, and one deep
blue, corresponding to the two wave-lengths which are alone contained
in the original light. In a more powerful instrument other fainter
lines are seen also. Naturally the brightness of all these lines is
together equal to that of the original beam. No other substance gives
any of those particular lines, although many others give blue and red
light of somewhat different wave-lengths. Thus, strontium compounds
give a blue light along with several red tints, but when strontium and
potassium are used together, the lines are found not to be coincident.
In the case of strontium, all the lines lie nearer to the yellow than in
POTASSIUM AND AMMONIUM 563
that of potassium. Since the whole light of the compound is thus con-
centrated in one or two narrow strips easily visible against a dark
background, small amounts of the elements give effects which are
readily recognizable in the instrument. This remains true even when,
to the eye, the colors are completely obscured by the much more bril-
liant, yellow light which compounds of sodium produce. In the spec-
trum of sodium, this yellow light is all concentrated into two yellow
lines which lie very close together.
Helium gives many lines, but one orange line (D8), in particular, was
noted in the spectrum of the sun's photosphere many years before the
element was obtained from terrestrial sources by Ramsay. When the
spectra of helium and other gases are to be examined in the laboratory,
a little of the material is inclosed in a narrow, exhausted tube, through
which an electrical discharge can be passed between platinum wires.
Under this treatment helium shows its conspicuous orange line, and
hydrogen a red and two blue ones. In this apparatus compounds are
dissociated and give the spectra of their constituents. When a Bunsen
flame is used with the salts* of metals, however, the temperature is not
high enough to render visible the spectra of the non-metals contained
in them. Indeed, even of the metals themselves, only the members
of the alkali and alkaline-earth groups give distinct results.
RUBIDIUM AND CAESIUM.
Soon after the invention of the spectroscope by Bunseri and Kirch-
hoff, the instrument was applied to the examination of many sub-
stances. In 1860 Bunsen discovered several new lines in the spectrum
given by materials derived from the salts in Durkheim mineral water.
Two new elements of the alkali group were found to cause their
presence, and were named, from the colors of the lines which they gave,
rubidium (red) and caesium (blue). Both elements have since been
found in small quantities in various minerals. Rubidium is obtain-
able with relative ease from the mother-liquors of the Stassfurt works.
The metals may be obtained by heating their hydroxides with
magnesium powder. The salts of these two elements are, in crystal-
line form and solubility, very much like those of potassium. In some
cases the difference in solubility is sufficient to make separation pos-
sible. Thus, a mixture containing compounds of these two metals
and of potassium gives with chloroplatinic acid a yellow precipitate,
* The chlorides are preferred because of their volatility. The salts of the oxy-
gen acids are dissociated, and leave the highly involatile oxides (e.g. pp. 380, 444).
564 INORGANIC CHEMISTRY
consisting of the three insoluble chloroplatinates. The solubilities at
10°, however, are as follows : Potassium chloroplatinate 0.9, rubidium
chloroplatinate 0.15, caesium chloroplatinate 0.05. Hence, when the
mixed precipitates are carefully washed with small quantities of cold
water the potassium chloroplatinate can be almost entirely removed.
On similar principles the two other metals can be separated from one
another. The iodides of all three elements combine with iodine, giving
tri-iodides (cf. p. 235), of which the tri-iodide of caesium is the most
stable. Whether this is to be regarded as showing that the metals may
occasionally be trivalent, or whether the extra iodine must be held to
have entered into combination with the iodine of the compound, and not
with the metal, has not been determined. In the parallel case of
hydriodic acid, the union with extra iodine (p. 359) seems to show con-
clusively that iodidion can combine with iodine. While an inclination
to tri valence in one of the metals of the alkalies would furnish a very
acceptable link between the two sides of the first column in the peri-
odic table (p. 411), since gold is a trivalent element, the latter of the
two above assumptions is more probably the correct one.
AMMONIUM.
The compounds of ammonium claim a place with those of the
alkali metals because in aqueous solution they give the ion NH/, an
ion which in its behavior closely resembles kalion. Some of the
special properties peculiar to ammonium compounds have been dis-
cussed in detail already (pp. 420, 421).
Ammonium Chloride. — This salt, known commercially as
salammoniac, like all the other compounds of ammonium, is prepared
from the ammonia dissolved by the water used to wash illuminating-
gas (p. 418). It is purified by sublimation, and then forms a com-
pact fibrous mass. It crystallizes from solution in cubes or octahe-
drons, which are often arranged according to a feathery pattern.
When heated to 350° it volatilizes and is almost completely dissoci-
ated into ammonia and hydrogen chloride at this temperature (p. 421).
Ammonium bromide and ammonium iodide are white salts which
crystallize in cubes or octahedrons, and are isomorphous with the cor-
responding potassium salts. They are dissociated by heat, and, in
the case of the iodide, some of the hydrogen iodide is still further
decomposed, giving free iodine.
POTASSIUM AND AMMONIUM 565
Ammonium Hydroxide. — The nature and behavior of this sub-
stance have been fully discussed (p. 421). It may be remarked here
that its very small basic activity as compared with that of potassium
hydroxide is only in part due to the low degree of ionization of its
molecules. A normal solution of ammonia contains much free NH3,
besides the NH4OH produced by its union with water. Thus, some
part of the material — how much we have at present no means of de-
termining — is not actually in the form of a base and is not in directly
ionizable condition at all. There are indications that the amount of
uncombined ammonia may be considerable. Thus the organic deriva-
tive tetramethylammonium hydroxide N(CH3)4OH is a very active
base indeed, and one of the most conspicuous differences between it
and ammonium hydroxide is that it cannot decompose into water and
a non-ionizable substance. It is all available for ionization, while the
material in ammonia-water is not.
Ammonium Nitrate. — This is a white crystalline salt which
may be made by the interaction of ammonium hydroxide and nitric
acid. When heated gently it decomposes, giving nitrous oxide and
water (p. 450). It is used as an ingredient in fireworks and explo-
sives. It exists in no fewer than four solid physical states. The
melted salt solidifies at about 160°, giving crystals of the regular
system. When these are allowed to cool somewhat, and are held at
a temperature a little below 125.5°, they change gradually into a mass
of rhombohedral crystals, the specific gravity and all other physical
properties altering at the same time. This temperature is a transition
point like that at which monoclinic sulphur assumes the rhombic form
(p. 368). When these rhombohedral crystals, in turn, are held at a
temperature a little below 83° they change their form once more into
crystals which belong to the rhombic system and possess a third dis-
tinct set of physical properties. Finally, below 35° a fourth change,
into rhombic needles, takes place, and this condition of the substance
is the one familiar at ordinary temperatures. All these changes pro-
ceed in the reverse order when the temperature is elevated once more.
Ammonium Carbonate. — When ammonium hydroxide is treated
with excess of carbon dioxide the solution gives, on evaporation,
ammonium bicarbonate NH4HC03. This is a white crystalline salt
which is fairly stable at the ordinary temperature. It has, however,
a faint odor of ammonia, and its dissociation becomes very rapid when
566 INORGANIC CHEMISTRY
slight heat is applied. When a solution of this salt is treated with
ammonium hydroxide, the neutral carbonate is formed :
NH4HC03 + NH4OH <=> (NH4)2C03 + H20.
But this salt, when left in an open vessel, loses ammonia very rapidly,
and leaves the bicarbonate behind.
The substance commonly sold as ammonium carbonate is the so-
called sesquicarbonate, and is made by sublimation from a mixture of
ammonium chloride or ammonium sulphate and chalk or powdered
limestone. It is a mixture, in approximately equi-molar proportions,
of ammonium bicarbonate and ammonium carbamate. The latter is a
substance related to urea, and formed when ammonia and carbon
dioxide gases are mixed :
Ammonium cyanate is interesting on account of its rapid trans-
formation, when warmed, into urea (p. 488). Ammonium thiocyanate
is a white salt which finds some application in analysis.
Ammonium Sulphate. — This is a white salt, crystallizing in
rhombic prisms, which is used chiefly as a fertilizer. By electroylsis
of a concentrated solution of the bisulphate, ammonium, persulphate,
which is less soluble, is formed and crystallizes out (cf. p. 397).
Sulphides of Ammonium. — When gaseous hydrogen sulphide
and ammonia are mixed in equi-molar proportions and compressed or
strongly cooled, ammonium hydrogen sulphide NH4HS is formed as
a crystalline deposit on the vessel. In an open vessel, at the ordinary
temperature, this solid dissociates slowly into its constituents. The
sulphide, (NH4)2S, can be produced under similar conditions by using
twice as much ammonia. But it is much less stable and gives up half
its ammonia, producing the acid sulphide very quickly. Solutions of
these sulphides, made by passing hydrogen sulphide gas into ammo-
nium hydroxide, are much used in analysis. The sulphide is almost
completely hydrolyzed by water into the acid sulphide and ammonium
hydroxide, its behavior being like that of sodium sulphide (p. 375) :
- H2S <=> (NH4)2S t=f 2NH4' + S" )
'
POTASSIUM AND AMMONIUM 567
It is used for the precipitation of sulphides, such as zinc sulphide,
which are insoluble in water. Although the S" ions are not numerous
at any moment, disturbances of the equilibrium by their removal, when
they pass into combination, causes displacements which result in the
generation of a continuous supply. The liquid smells strongly of
ammonia and hydrogen sulphide on account of the dissociation of the
parent molecules. Because of this dissociation the >salt is preferred to
potassium or sodium sulphide in analysis. The excess of the reagent
can be driven out by simply boiling the mixture for a few minutes, all
of the above equilibria being reversed. Another application in analy-
sis depends on the tendency of this salt to unite with certain insoluble
sulphides, particularly those of tin, arsenic, and antimony (q.v.^), giving
soluble complex salts.
The solution dissolves free sulphur, giving yellow polysulphides
similar to those of potassium (p. 561). The same yellow substances
are also obtained by gradual oxidation of ammonium sulphide when the
solution of this salt is allowed to stand in a bottle from which the air is
imperfectly excluded.
Microcosmic Salt. — This salt would be named, systematically,
the tetrahydrate of secondary sodium-ammonium orthophosphate
(NaNH4HP04, 4H20). When ammonium chloride and ordinary sodium
phosphate are mixed in strong solution the hydrate crystallizes out.
The substance is used in bead tests (cf. pp. 467, 468).
Ammonium Amalgam. — As we have seen (p. 550), a potassium
ion may be discharged and the element which it contains secured in the
free condition. When a salt of ammonium is decomposed by electro-
lysis, however, the NH4 ion upon its discharge gives ammonia and hy-
drogen, and no substance NH4 is obtained. If, however, a pool of
mercury is used as the negative electrode, the NH4 forms an amalgam
with it, and there seems to be no doubt that this substance is actually
present in solution in the mercury. While the amalgam is being formed
it swells up and gives off the decomposition products above mentioned,
so that the existence of the substance is only temporary. The same
material may be obtained by putting sodium amalgam into a strong
solution of a salt of ammonium. The action is a displacement of
one ion by another (p. 361) :
Ka (diss'd in mercury) -+- NH4* —> NH4 (diss'd in mercury)-)- Na*.
568 INORGANIC CHEMISTRY
This behavior is interesting since it is in harmony with the idea that
ammonium, if it could be isolated, would have the properties of a metal.
Substances, other than metals, are not miscible with mercury.
Ammonion : Analytical Reactions. — Ionic ammonium is a color-
less substance. It unites with negative ions, giving salts, which, in the
majority of cases, are soluble. Ammonium chloroplatinate, and to a
less extent ammonium hydrogen tartrate, are insoluble compounds, and
their precipitation is used as a test. The surest means of recognizing
ammonium compounds, however, consists in adding a soluble base to
the substance (cf. p. 421). The ammonium hydroxide, which is thus
formed, gives off ammonia, and the latter may be detected by its odor.
The quantity of the ammonium salt present may be determined by dis-
tilling the mixture and catching the distillate in a measured volume of
normal hydrochloric acid. Determination of the amount of the acid
remaining unneutralized, by titration with a standard alkali solu-
tion, then gives, by difference, the quantity of ammonium hydroxide.
Exercises. — 1. What kind of metals will, in general, interact with
solutions of bases (cf. p. 553) ?
2. Why should a mixture of potassium chlorate and antimony tri-
sulphide be explosive ?
3. How does the direct vision spectroscope differ from the arrange-
ment here described (cf. any work on physics) ?
4. Why is not ammonium carbamate (p. 566) formed by the neutral-
ization method?
5. How should you set about making : a borate of potassium,
potassium pyrophosphate, ammonium nitrite ?
6. Why is the cleaning of platinum wires, as usually effected by
holding them in the Bunsen flame, assisted by periodical dipping
into hydrochloric acid (p. 542) ?
CHAPTER XXXIV
SODIUM AND LITHIUM. IONIC EQUILIBRIUM CONSIDERED
QUANTITATIVELY
SODIUM chloride forms more than two-thirds of the solid matter
dissolved in sea- water, and the great salt deposits are largely composed
of it. Sea-plants contain sodium salts of organic acids, just as land-
plants contain potassium salts. Chili saltpeter, cryolite, and albite (a
soda feldspar) are important minerals.
Preparation. — Sodium was first made by Davy (1807) by electro-
lysis of moist sodium hydroxide. It is manufactured by Castner's
process, which is used also for potassium (p. 550), and by the electro-
lysis of fused sodium hydroxide by a method likewise invented by
Castner. In the latter case the negative electrode projects through the
bottom of the iron vessel containing the fused hydroxide. This electrode
is surrounded by a wire-gauze partition, which is surmounted by a bell-
shaped vessel of iron. The positive electrode is an iron cylinder sur-
rounding the gauze. The sodium and hydrogen liberated at the cathode,
being lighter than the fused mass, ascend into the iron vessel, under
the edge of which the hydrogen escapes. Oxygen is set free at the
anode.
Properties. — Sodium is a soft, shining metal, melting at 95.6°
and boiling at 742°. The vapor is a monatomic gas. The metal is
soluble in liquefied ammonia, giving a blue solution. The amalgam with
mercury, when it contains more than a small amount of sodium, is
solid, and probably contains one or more compounds of the two elements.
This amalgam is often used instead of the metal sodium, since the
dilution or combination with mercury makes the interactions of the
metal more easily controllable. Sodium is used in the manufacture of
many complex carbon compounds which are employed as drugs and dyes.
Sodium Hydride. — - When hydrogen is led over sodium at 340°
in such a way that the upper part of the tube is cooler, a matted mass
of fine white crystals of the hydride is deposited on the cool part of
6GU
570 INORGANIC CHEMISTRY
the tube. The temperature must not rise beyond 430°, since the com-
pound dissociates rapidly at this temperature. The properties of the
substance are similar to those of potassium hydride (p. 550).
Sodium Chloride. — Common salt is obtained from the salt de-
posits of Stassfurt, Reichenhall (near Salzburg), in Cheshire, at Syra-
cuse and Warsaw in New York, at Salina in Kansas, in Utah, Cali-
fornia, and many other districts. Natural brines are obtained from
wells in various parts of the world. Since the salt can seldom be
used directly, on account of impurities which it contains, it is purified
by re crystallization from water. Natural brines, which are sometimes
dilute, are often concentrated by dripping over extensive ricks com-
posed of twigs. When the resulting brine is allowed to evaporate
slowly by the help of the sun's heat, large crystals, sold as " solar
salt," are obtained. By the use of artificial heat and stirring, smaller
crystals of greater purity can be secured. Salt intended for table
use must be freed from the traces of magnesium chloride (g.v.') present
in the original brine or deposit, for this impurity causes it to absorb
moisture more vigorously from the air. The purest salt for chemical
purposes is precipitated from a saturated solution of salt by leading
into it hydrogen chloride gas. Explanation of this effect will be
given presently (see p. 584).
Common salt crystallizes in cubes, the faces of which are usually
hollow. The crystals decrepitate (p. 556) when heated, and melt at
about 820°. Common salt is the source of all sodium compounds, with
the exception of the nitrate. From it come also most of the chlorine
and hydrogen chloride used in commerce.
The Hydroxide and Oxides. — Sodium hydroxide is prepared
both by the action of slaked lime upon sodium carbonate and by the
electrolysis of a solution of sodium chloride, precisely as is potassium
hydroxide (p. 554). An interesting method used at Niagara Falls,
the Acker process, employs molten sodium chloride as the electrolyte
and molten lead as the cathode. The sodium forms an alloy with the
lead, which continually circulates along the trough-shaped cell and
back beneath a partition. On issuing at one end it comes in contact
with a jet of steam, and this removes the sodium from the lead, forming
sodium hydroxide. The latter flows in the molten condition from the
surface of the lead, and the metal returns to the trough.
Sodium hydroxide is a highly deliquescent substance, which, when
SODIUM 571
exposed to the air, first liquefies and then becomes solid on account of
the formation of sodium carbonate. Its general chemical properties
are identical with those of potassium hydroxide. It is used in the
manufacture of soap, in the preparation of paper pulp, and in many
other chemical industries.
Sodium peroxide Na202 is made by passing slices of sodium, rest-
ing upon trays of aluminium, through a tubular vessel. In this it comes
in contact with a current of air which has been freed from carbon
dioxide and is maintained at a temperature between 300° and 400°.
This oxide when thrown into water decomposes in part, in consequence
of the heat developed, giving sodium hydroxide and oxygen. With
careful cooling, however, much of it can be dissolved. By interaction
with acids it yields hydrogen peroxide. Sodium peroxide is now used
commercially for oxidizing and bleaching. The ordinary sodium oxide
is made in the same way as is potassium oxide (p. 555).
The Nitrate and Nitrite* — The occurrence and purification of
sodium nitrate have already been described (p. 438). Its crys-
tals are of rhombohedral form (Fig. 9, p. 14). This salt is one
of the best of fertilizers, since it furnishes to plants the nitrogen
which they require in a very easily absorbed form. It is used also in
the manufacture of potassium nitrate, of nitric acid, and of sodium
nitrite.
Sodium nitrite is formed by heating sodium nitrate with metallic
lead and recrystallizing the product (p. 449). Although very soluble
it is less so than potassium nitrite, and is therefore more easily pre-
pared in pure condition. It is used as a source of nitrous acid by
manufacturers of organic dyes.
Manufacture of Sodium Carbonate. — Natural sodium carbon-
ate is found in Egypt and in other parts of the world. At Owen's
Lake, California, it is secured by solar evaporation of the water.
The sesquicarbonate Na2C08,ISraHC08,2H20, being the least soluble
of the carbonates of sodium, is the one deposited. Locally, small quan-
tities of sodium carbonate are still made by the burning of sea-weed.
Up to the close of the eighteenth century this was the only source of
the compound, and the product from Spain, known commercially as
barilla, was ten times as expensive as the carbonate now is. Hence
glass and soap were proportionately dearer than at present.
In 1791 the French Academy offered a prize for the discovery of
572
INORGANIC CHEMISTRY
an inexpensive method for the preparation of sodium carbonate from
common salt, and Le Blanc proposed the process which bears his name
and is still in use. During the Eevolution his factory was destroyed,
his patents were declared to be public property, and the inventor died
by suicide. The chief stages of Le Blanc's process involve three chemi-
cal actions. In the first place, sodium chloride is treated with an equi-
valent amount of sulphuric acid in a large cast-iron or earthenware
FIG. 100.
pan. The bisulphate thus produced (cf. p. 178), together with the
unchanged sodium chloride, is raked out on to the hearth of a rever-
beratory * furnace (Fig. 100) and heated more strongly, while being
continually worked by means of rakes, until the action is completed :
NaCl + NaHS04 +± Ka2S04 + HCl.f . ;
The product of this treatment is called salt-cake. The hydrogen
chloride, which is liberated in both stages, passes through towers
containing running water in which it is absorbed. The second and
third actions which follow are conducted in one operation. They con-
sist in the reduction of the sodium sulphate by means of powdered
coal and the interaction of the resulting sulphide of sodium with
chalk or powdered limestone :
NaS
+ 2C -
CaC0
Na2S + 2C0
+ Na2C08
2,
CaS.
In the less modern factories the salt-cake, limestone, and coal are
stirred upon the hearth of a reverberatory furnace and worked by hand.
* So called because the heated gases from the fire are deflected by the roof and
play upon the materials spread on the bed of the furnace.
SODIUM 573
The material is finally collected into balls, and the end of the action is
recognized by the fact that bubbles of carbon monoxide begin to force
their way to the surface and cause little jets of blue flame. The gas
is produced by the action of the coal upon the calcium carbonate,
excess of both of these substances being present :
CaC08 + G -> CaO + 2CO.
The production of this gas gives a porous texture to the material,
which facilitates the solution of the sodium carbonate in the final
stage. The porous product is called black-ash. In modern factories
hand labor is saved by giving the black-ash furnace the form of a
rotating cylinder, in which projections from the walls assist in bring-
ing about complete mixing of the materials during the action.
The black-ash varies very much in composition. It commonly
contains 45 per cent of sodium carbonate, 30 per cent of calcium
sulphide, 10 per cent of calcium oxide, and a number of other products
and impurities. The coal used in the operation is selected so as to be
as free as possible from combined nitrogen, the presence of which
leads to the formation of cyanides.
Calcium sulphide is not very soluble in water, and is but slowly
hydroly/ed by it (p. 376), especially when calcium hydroxide is present.
The sodium carbonate is therefore extracted from the black-ash by a
systematic treatment of the ash with water. The ash is placed in a
series of vessels at different levels, and a stream of water flows from
one vessel to another, until, when it issues from the last, it is com-
pletely saturated with sodium carbonate. A temperature of 30° to 40°,
at which the solubility of sodium carbonate is at a maximum, is
employed. When the material in the first of the vessels has been
exhausted, the water is allowed to enter the second vessel directly, and a
vessel containing fresh black-ash is added at the lower end of the series.
In this way the most nearly exhausted ash comes in contact with pure
water, .which is in the best position to dissolve the remaining sodium
carbonate rapidly, while the fresh black-ash encounters a solution
already almost at the point of saturation.
The saturated solution is evaporated in shallow pans placed in the
flues of the furnaces, and the monohydrate Na2C03, H20, which crystal-
lizes from the hot liquid, is raked out and dried by heat, leaving cal-
cined soda. When this material is recrystallized from water and is
allowed to deposit itself from the solution at the ordinary temperature,
the decahydrate Na.,CO8,l()H20, soda crystals or washing soda, appears.
574 INORGANIC CHEMISTRY
The solid residue from the extraction of the black-ash is known as
tank waste, and contains 35-55 per cent of calcium sulphide. This
material contains the sulphur of the original sulphuric acid, and its
treatment involves a problem of some difficulty. If it is dumped near
the factory the sulphur is lost, and by slow weathering yellow solutions
containing poly sulphides flow from the decomposing heap into the
streams, and offensive odors of hydrogen sulphide fill the air. The
most effective process for the recovery of the sulphur and consequent
abatement of this nuisance is that of Chance. The product is arranged
in a series of cylinders through which is passed carbon dioxide from
a kiln of special form. The hydrogen sulphide liberated in the first
cylinder forms the acid sulphide with the material contained in the
second :
2CaS + C02 + H2O -> Ca(SH)2 + CaC03.
The further action of the carbon dioxide on this product gives, finally,
a mixture of gases containing a larger proportion of hydrogen sul-
phide. By burning this mixture with a limited supply of air, the
sulphur is then secured in free condition :
2H2S + 02 -> 2H20 + 2S.
About 70,000 tons of sulphur are thus recovered annually.
The Solvay, or ammonia-soda process, invented in 1860, is a seri-
ous rival of the Le Blanc process. It differs from the latter by
involving almost nothing but ionic actions. A solution of salt con-
taining ammonia and warmed to 40° fills a tower divided by a number
of perforated partitions. Carbon dioxide, which is forced in below,
makes its way up through the liquid. The ammonium bicarbonate
formed by its action undergoes double decomposition with the salt, and
sodium bicarbonate which is precipitated settles upon the partitions :
NaCl+NH4HC03 <± NaHC03 [ + NH4C1, or HC03' + Na' <=> NaHC08|.
The solid sodium bicarbonate, after being freed from the liquid, is
heated strongly and leaves behind sodium carbonate :
2NaHC03 -> Na,C03 + H20 f + CO2 f.
The carbon dioxide which is liberated passes through the operation
once more. The mother-liquor from the sodium bicarbonate contains
ammonium chloride. This is decomposed by heating with quicklime,
and the ammonia which is thus obtained is available for the treatment
of another batch. The supply of carbon dioxide is generated in
SODIUM 575
lime-kilns of special form. The lime produced in these kilns serves
for the liberation of the ammonia.
This process is cheaper than that of Le Blanc, and furnishes a
much purer product. The latter process continues to be used, how-
ever, because of the hydrochloric acid which is produced at the same
time. This finds remunerative application in the liberation of its
chlorine for the manufacture of bleaching powder.
A possible rival of these two processes threatens to arise in the
treatment of electrolytic sodium hydroxide with carbon dioxide gas.
Properties of Sodium Carbonate* — The common form of
sodium carbonate consists of large monoclinic crystals of the decahy-
drate. This substance has a fairly high aqueous tension, and loses
nine of the ten molecules of water which it contains when it is exposed
in an open vessel. When warmed it melts at 35.2°, giving a solution
of sodium carbonate in water. The residue from evaporation, above
35.2°, is the monohydrate. A.t higher temperatures, or with low at-
mospheric aqueous tension (p. 121), this in turn can be completely dehy-
drated (see Chap. xxxv). The decahydrate increases in solubility up
to 35.2°, when it ceases to exist. Just above this temperature the mono-
hydrate is the only substance which is stable. Its solubility is the
same as that of the decahydrate at 35.2°, and diminishes as the tem-
perature is raised. The relations are of the same nature as in the case
of sodium sulphate (p. 158). In aqueous solution, sodium carbonate is
hydrolyzed, and shows a marked alkaline reaction. The compound is
used in large amounts for the manufacture of glass and soap, and
is applied in innumerable ways in the scientific industries for pur-
poses akin to cleansing.
Nearly all the familiar compounds of sodium are formed in the
course of one or other of the processes by which sodium carbonate is
manufactured, or are made by the treatment of sodium carbonate or
sodium hydroxide with acids.
Sodium Bicarbonate. — This salt can be prepared in a state of
purity by passing carbon dioxide over the decahydrate of sodium
carbonate :
CO?=>2NaHC0 + 9H0.
The hydrate is spread upon a grating, through which the water generated
by the action drips away. This action is reversible, and sodium bicar-
576 INORGANIC CHEMISTRY
bonate shows, even in the cold, an appreciable tension of carbon diox-
ide. Even a solution of this salt gives off carbon dioxide, when boiled.
An aqueous solution of pure sodium bicarbonate is neutral to phenol-
phthalein on account of the small degree of ionization of the ion HCO/.
The salt is used in the manufacture of baking powder and in medicine.
Sodium Sulphate. — Anhydrous sodium sulphate (thenardite)
crystallizes in the rhombic system, and is found in the salt layers.
The same salt is contained in mineral waters, such as those of Frie-
drichshall and Karlsbad. It is formed in connection with the manu-
facture of nitric acid from sodium nitrate, and as an intermediate
product in the making of sodium carbonate. Some of it is also pre-
pared at Stassfurt by dissolving kieserite (MgS04,H20) in water
along with sodium chloride. When the solution is cooled to 0°, the
decahydrate of sodium sulphate crystallizes out, and magnesium chlo-
ride remains in solution.
The decahydrate of sodium sulphate, Glauber's salt, forms large
monoclinic crystals which give up ten molecules of water when kept in
an open vessel. When heated the crystals melt at 32.4°, and are re-
solved into the sulphate and water. The relations of the hydrate and
anhydrous substance in respect to solubility have been fully discussed
already (p. 158). When the decahydrate is mixed with concentrated
hydrochloric acid, it is decomposed, and a part of the sulphate is con-
verted into sodium chloride, the second action being a reversible one.
This is one of those actions which proceed spontaneously, and there-
fore involve a diminution in the store of available energy in the
system, although, so far as heat is concerned, a marked absorption of
this form of energy takes place (cf. p. 27). The combination is used,
in fact, as a freezing mixture.
Sodium Thiofsulphate. — This salt is made by boiling a solution of
sodium sulphite with sulphur. It is also obtained by boiling sulphur
with caustic soda, and crystallizes from the mixed solution :
48 + 6NaOH -* Na^SA + 2^8 + 3H20.
It gives large, transparent monoclinic crystals of a pentahydrate.
When heated it first loses the water of hydration, and then decomposes,
giving sodium sulphate, which is the most stable oxygen-sulphur com-
pound of sodium, and sodium pentasulphide :
j- Na2S8.
LITHIUM 577
From the latter, four unit- weights of sulphur can be driven by stronger
heating. Sodium thiosulphate is used for fixing negatives in photog-
raphy (q.v.), and by bleachers as antichlor (p. 396).
Phosphates of Sodium. — Common sodium phosphate is a do-
decahydrate of the secondary orthophosphate, Na2HP04,12H20. It is
made by neutralization of phosphoric acid with sodium carbonate, and
crystallizes from the solution in large, transparent monoclinic prisms.
Its properties have already been discussed (pp. 466, 467).
Sodium metaphosphate NaP08 is used for bead reactions (cf. p. 468).
Sodium Tetraborate. — This salt combines with ten molecules of
water, forming large, transparent prisms, Na2B4O7, 10H20. When
heated, it loses, water, and leaves the easily fusible anhydrous salt in
glassy form. Its sources have already been discussed under boric acid
(p. 528). It is used as an ingredient in glazes for porcelain, in
soldering, and for bead reactions (cf. p. 528).
Sodium Silicate. — A salt essentially of the composition of the
metasilicate NagSiOg (cf. p. 522) is used for fire-proofing wood and other
materials. Sand which is moistened with it and pressed in molds
forms, after baking, a serviceable artificial stone. Since silicic acid is
a feeble acid, this salt is much hydrolyzed, and gives a strongly alka-
line solution (p. 523).
Properties of Natrion : Analytical Reactions. — Natrion is
a colorless ionic material which unites with all negative ions. Practi-
cally all the salts so formed are soluble in water. The only ones which
can be precipitated are sodium fluosilicate, made by the addition of
hydrofluosilicic acid to a strong solution of a sodium salt, and sodium
hydrogen pyroaiitimoniate Na2H2Sb207, made by similar addition of the
corresponding potassium salt. For the recognition of natrion the
solution is evaporated, and the residue examined with a spectroscope.
If the yellow light persists longer than could be accounted for by the
ordinary deposit of dust on the wire, a sodium compound is present in
the material.
LITHIUM.
The compounds of lithium are made from amblygonite, a mixed
phosphate and fluoride of aluminium and lithium. It occurs in lepido-
lite (a lithia mica) and in other rare minerals. Traces of compounds
578 INORGANIC CHEMISTRY
of the element are found widely diffused in the soil, and are taken up
by plants, particularly tobacco and beets, in the ashes of which the
element may be detected spectroscopically.
The metal is liberated by electrolysis of the fused chloride, the
manipulation being facilitated by the addition of some potassium chlo-
ride to lower the melting-point of the lithium salt. The melting-point
and boiling-point of the free element are higher than those of any
other alkali-metal, and the specific gravity (0.59) is lower than that
of any other metal whatever. Lithium not only floats upon water,
but also in the petroleum in which it is preserved.
The metal behaves towards water and oxygen like sodium (p. 548).
It unites directly and vigorously with hydrogen (LiH), nitrogen
(Li8N), and oxygen (LijO), forming stable compounds. The chloride
crystallizes in octahedra (p. 549). The relative insolubility (p. 544)
of the hydroxide (LiOH), the carbonate (LijCOg), and the phosphate
(Li8P04,2H20), is in sharp contrast to the easy solubility of the cor-
responding compounds of the other alkali-metals, and links lithium
with magnesium. The compounds of lithium give a bright-red color
to the Bunsen flame, and a bright-red and a somewhat less bright orange
line are seen in the spectrum. The carbonate is used in medicine.
IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY.
The Simple Case. — In view of the predominance of ionic actions
in the chemistry of the metals, and of the determinative effect of ionic
equilibria on many actions, it is essential that we should be prepared
in future for a more exact study of these phenomena than we have
hitherto attempted. The whole basis for this exact study has already
been supplied, and only more specific application of the principles is
demanded.
In the first place, the principles themselves must be recalled.
When acetic acid, for example, is dissolved in water, it is ionized thus :
Ct C2 63
The amount of molecular acetic acid dissociated per second in a given
amount of the solution is proportional to the concentration of the mole-
cules ((7j), while the amount of the two ionic materials, hydrion and
acetanion, uniting to form molecules of acetic acid depends on the
frequency of the encounters of the two kinds of ions and is propor-
IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 579
tional to the ionic concentrations (p. 297). The unit of concentration
(p. 250) is 1 mole per liter, or, in the present case, 60 g. of the acid,
1 g. of hydrion, and 59 g. of acetanion respectively, per liter, for these
numbers represent the weight of one mole of each component. Accord-
ing to the law of concentration (p. 252) :
<±*^ = K, (1)
and the numerical value of this fraction, or of K, remains unchanged
whatever the total concentration of the solution may be. If the solution
is diluted, for example, (72 and Cs diminish relatively less quickly than
Cl in order that the value of the whole expression may remain the
same. This is accomplished by ionization of a part of the material
whose concentration is Cl and its transference to the ionic forms whose
concentrations are C2 and C8, respectively (p. 297).
A numerical example will show that this law of concentration ex-
presses the facts with considerable exactness. The data in regard to
acetic acid are as follows (p. 328) :
Acetic Acid.
I.
Molar Con-
centra-
tion.
II.
Proportion
of Whole Mate-
rial Ionized.
III.
Molar Concentrations
of H' (Cz)
and of 02H,02' (C3).
(I X II.)
IV.
Molar Concentra-
tion of
H<JfiW
Uni-molar
1.0
0.004
0.004
1.0 -0.004
Deci-molar .
0.1
0.013
0.0013
0.1 - 0.0013
Centi-rnolar .
0.01
0.0407
0.000407
0.01 -0.000407
Now C 2 = Cs, since the ions are produced in equal numbers. Also,
for our purpose, the numbers to be subtracted in column iv are rela-
tively so small that the values 1, 0.1, and 0.01 may be taken to repre-
sent Cl without appreciable error. Hence, substituting the data in
equation (1) above, we have :
(.004)2
= .0,160.
(.0013)2
= .04169.
(.000407)2
= .04165.
In other words, although the last solution is a hundred times more
dilute than the first, and the degree of ionization has increased ten
times, the whole expression remains close to the value 0.04165 and is
essentially constant.
580 INORGANIC CHEMISTRY
When conductivity data, like the above, are applied in the same
way to the cases of more highly ionized substances, the values of K
are less nearly constant. It is supposed that with this class of sub-
stances the measurements of degrees of ionization are less accurate,
although the cause of the discrepancy has not been fully determined.
However, in the very general applications of the data, which are all
that we shall be required to make, the conclusions will not be affected
by this fact.
The expression for ionic equilibrium in terms of concentration is
often written in the following fashion :
[H']x[C,H.O/]_
[HC2H302]
Here [H*] represents the molar concentration of hydrion ((72),
[C2H302'] that of acetanion ( (73), and [HC2H302] that of molecular
acetic acid (C^). This form is more convenient for many reasons, and
will be employed frequently.
Excess of One Ion. — When, through the presence of two sub-
stances with a common ion, (72 is not equal to CS) the above law (form-
ula 1) still holds. For example, if with one mole of sodium acetate and
one mole of acetic acid a single liter of solution is prepared, the solu-
tion will be uiii-molar in respect to the acid and to the salt as well.
Thus, all the acetanion will be available for uniting both with the hydrion
and the natrion. Hence (73, in the expression for the acetic acid equi-
librium, will be abnormally large, and the ionization of the acid will be
repressed. In uni-molar sodium acetate, 0.53 of the salt is ionized
(p. 331), and, initially in the mixture of acid and salt, the concentration
of acetanion will be .53 + .004 = .534, or nearly 134 times larger than
in the acid alone.' Hence, in order that the product C2 X C8 may re-
cover, as it must, a value much nearer to the old one, (7, must be
diminished to something' like T^¥ of its former magnitude. That is,
C2 will become equal to about 0.00003, the rest of the hydrion uniting
with a corresponding amount of the acetanion to form molecular acetic
acid. The effect of adding this amount of sodium acetate therefore is
to reduce the concentration of the hydrion below the amount which
can be detected by use of an indicator like methyl-orange. There is
no less acetic acid present than before, but the number of acid ions is
very much smaller.
This action is of course reciprocal, and the ionization of the sodium
IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 581
acetate will be reduced also. But the acetauion furnished by the
acetic acid is relatively so small in amount (.004 against .53) that the
effect it produces on the ionization of the salt is imperceptible.
It will be noted that the acetanion and hydrion disappear in equi-
valent quantities, for they unite. There is, however, so much of the
former that its loss goes unremarked, while there is so little of the
latter that almost none of it remains. When substances of more nearly
equal degrees of ionization are used, both effects are equally incon-
spicuous. Thus, sodium chloride and hydrogen chloride in uni-molar
solutions yield approximately equal concentrations of chloridion
(.784 and .676). Hence, if One mole of sodium chloride were to be
dissolved in the portion of water already containing one mole of hydro-
gen chloride, the concentration of the chloridion, at a very rough
estimate, would be nearly doubled. If this doubling of the concentra-
tion of chloridion almost halved that of the hydrion (.784), in order
that the expression [Cl'] x [H'] -i- [HC1] might remain constant, the
concentration of the hydrion would still be about .400 and therefore
100 times as great as in molar acetic acid. It is thus altogether im-
possible to reduce the concentration of the hydrion given by an active
acid like hydrochloric acid below the limit at which indicators are
affected, for there is no way of introducing the enormous concentra-
tion of the other ion which the theory demands.
With more crude means of observation than indicators afford,
effects like this may sometimes be rendered visible. This was the case
with cupric bromide solution, to which potassium bromide was added
(p. 335). The blue of the cuprion disappeared from view, while much
cuprion was still present, because the brown color of the molecular
cupric bromide covered it up completely.
. Special Case of Saturated Solutions. — The commonest as
well as the most interesting application of the conceptions developed
above is met with in connection with saturated solutions, especially
those of relatively insoluble substances.
The situation in a system consisting of the saturated solution and
excess of the solute has been discussed already (p. 158). In the case
of potassium chlorate, for example, we have the following scheme of
equilibria :
KC103 (solid) <± KC103 (diss'd) <=> K' + C108'.
Solution of the solid is promoted by the solution pressure of the mole-
cules, while it is opposed by the osmotic pressure of the dissolved
582
INORGANIC CHEMISTRY
substance, and the solution is saturated when these tendencies produce
equal effects (p. 152). Now it must be noted that the tendency
directly opposed to the solution pressure is the partial osmotic press-
ure of the dissolved molecules alone. The chief contents of the solu-
tion, the molecules and two kinds of ions of the salt, and any foreign
material that may be present, are like a mixture of gases, and the
principle of partial pressure (p. 88) is to be applied. The ions and
the foreign material do not deposit themselves upon the solid, and take,
therefore, no part directly in the equilibrium which controls solubility.
In respect to this the ions are themselves foreign substances. Hence
the conclusion may be stated that in solutions saturated at a given
temperature by a given form (p. 160) of the solute, the concentration of
the dissolved molecules considered by themselves will be constant
•whatever other substances may be present.
The total solubility of a substance, as we have used the term
hitherto, is made up of a molecular and an ionic part. The latter, as
we shall presently see, is not constant when
a foreign substance containing a common ion
is already in the liquid. Since the treatment
of the subject requires us now to distinguish
between the two portions of the solute, a dia-
gram (Fig. 101) will assist in emphasizing
the distinction. The material at the bottom
is the salt. The molecules and ions are to
be thought of as being mixed and as being
present in numbers represented by the fac-
tors n and m. Since no foreign body is pres-
ent, the two ions in this case are equal in
number.
When we now apply these ideas to the
mathematical expression of the relation :
r1 v f1
\Jn A \J* -f-r
mK'+wCKV
u
n KC1O3
n
FIG. 101.
we perceive that, in a saturated solution, Cv the concentration of the
molecules is constant. Transposing, we have C2 X Cs = KCr Hence
the relation leads to the important conclusion that in a saturated
solution the product of the molar concentrations of the ions is constant.
The practical bearing of this inference appears at once when we
suppose an excess of one of the ions to be added, and remember that
jn chemical experiments mixtures of substances, including those with
IONIC EQUILIBRIUM, CONSIDERED* QUANTITATIVELY 583
common ions, are encountered much oftener than are pure solutions.
If the original substance is not very soluble, and the values of <72 and
Cs are therefore small, one may easily increase the value of C8 a
hundredfold by adding a material which is sufficiently soluble and
sufficiently ionized. To preserve the value of the product C2 X Ct,
the value of C2 will then have to be diminished at once to one-
hundredth of its former value. This can occur only by union of
the ionic material it represents with an equivalent amount of that
for which Cs stands. The molecular material so produced will thus
tend at first to swell the value of Cr But the value of Cl cannot
be increased, for the solution is already saturated with molecules, so
that the new supply of molecules, or others in equal numbers, will be
precipitated. Hence the ionic part of the dissolved substance may be
diminished, and the total solubility of the substance is seen to be in
reality a variable quantity. We have actually precipitated a part of
the dissolved material without introducing any substance, which, in
the ordinary sense, can interact with it.
The product of the concentrations of the ions C2 x C3 is called the
ion-product constant, or solubility product, because these two values
jointly determine the magnitude of the solubility of the substance.
The solubility of the molecules is irreducible, but the ionic part of
the dissolved material may become vanishingly small if the value
of C2 or (78 is very minute. The ionic part of any particular substance
is made up of the smaller of the two concentrations C*2 and Ca, plus
an equivalent amount, and no more, of the other. The rest of the
other is part of the solubility of some other component.
An Illustration : Potassium Chlorate. — A numerical ex-
ample will show how this operates. The molar solubility of potassium
chlorate at 18° is .52 (p. 544). At this concentration about .70 of the
salt is ionized. The concentration of each ion is therefore .70 x .52
= .36. That of the molecules is .52 - .36 = .16. The solubility may
therefore conceivably be reduced almost to .16 by admixture of a sub-
stance with a common ion, but cannot be carried below this value.
The ionic product Cz x €„ or [K*] x [C108'], is .36 x .36 = .13.
The molar solubility of sodium chlorate (p. 544), which is very
soluble, is 6.4. When, to the saturated solution of the potassium salt,
an equal volume of that of the sodium salt is added, some potassium
chlorate (the less soluble substance) is at once thrown down :
584 INORGANIC CHEMISTRY
Each solution is diluted by the volume of the other. The concen-
trations after mixing are therefore : KC103, .26 ; NaC103, 3.2. The
degree of ionization of the potassium chlorate is diminished somewhat
by the addition of the more concentrated sodium chlorate solution, that
of the latter somewhat increased by the mixing with the more dilute
potassium chlorate solution.* Leaving out of consideration the minor
change in the degrees of ionization, the concentration of the K" ions
would be approximately .70 x .26 = .18 ; that of the part of the C1O3'
ions coming from the potassium chlorate would be the same ; that of
the molecules .26 — .18 = .08. Of the sodium chlorate about .27 is
ionized, so that the concentration of each of its ions is roughly .27 x 3.2
= .86. Thus, the total concentration of chloranion is momentarily
.18 + .86 = 1.04. The product [K>] x [C103'], .18 x 1.04, is larger
than .13, and so each of the factors will be reduced by an equal amount
until the product equals .13. Leaving once more the changes in the
degrees of ionization out of account, for our rough calculation, the
equation (.18 — x) (1.04 — x) = .13 gives x = .05. Subtracting this
quantity, we have the concentrations .13 x .99 = .13 (nearly). That
is to say, .05 moles per liter will be precipitated from double the
original volume of the potassium chlorate solution, or .1 moles of the
original .52 moles in the solution. Thus, roughly, the final disposition
of the material originally present in one liter is :
Precipitated 10 moles.
Remaining dissolved, molecular .08 x 2 = .16 moles.
Remaining dissolved, ionic . .13 x 2 = .26 moles.
.52 moles.
Other Illustrations. — The precipitation of sodium chloride from
a saturated solution, by the introduction of gaseous hydrogen chloride
(p. 570), is to be explained in the same manner. A rough calculation,
similar to that made above, would show, however, that the fraction of
the whole material precipitated is smaller in the case of salt than in the
case of potassium chlorate. This is because salt is much more soluble.
With a substance even less soluble than potassium chlorate the effects
are, conversely, relatively much greater.
* The degree of ionization of a salt (e.g. KC1O3) in mixtures of salts which all
ionize with equal readiness, is approximately that which it would be in a pure solu-
tion of the salt (KC1O3) of a molar concentration equal to the total molar concen-
tration of all the salts present. For example, in the above case, for potassium
chlorate it would be that of a .26 + 3.2, or 3.46 molar solution, and for sodium
chlorate it would be the same.
IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 585
The evolution of a steady stream of hydrogen chloride is often
accomplished by allowing concentrated sulphuric acid to flow into
saturated hydrochloric acid (p. 182). The effect is due in part to repres-
sion of the ionization of the hydrogen chloride and elimination of mole-
cules of the gas from the water which is already saturated with
molecules of the same kind. The " salting out " of soap (p. 505) is
another phenomenon similar to those described above (see p. 721).
Direction of a, Chemical Change Determined by Ionic Con-
centrations. Rule for Precipitation or Solution of Insoluble
Substances. — In discussing the manufacture of potassium hydroxide
(p. 553), it was mentioned that the action of milk of lime on potassium
carbonate is reversed when the concentration of the KOH becomes too
great :
Ca(OH)2(solid) ^±Ca(OH)2(diss'd) «=*20H' + Ca" 1 <=> CaC03(diss'd) «=»
K2C03 «=> 2K- + C03" / CaC08(solid).
The whole theory of the most common influences of ionic concen-
tration may be illustrated by this one example.
We will imagine this action to be carried out at 18°, instead of
100°, since exact data in regard to the situation at the latter tempera-
ture are lacking. The molar solubility of calcium hydroxide at 18° is
.02, and about .88 of the dissolved molecules are ionized.
Thus the molar concentration of the Ca** ions is .88 of .02 or .0176.
That of the OH' ions, since each molecule gives two of them, will be
2 x .0176, or .0352. The solubility product is [Ca"] x [OH']2, in
which, when the substance is present by itself, the values are .0176 x
(.0352)2.
Calcium carbonate is precipitated by the action and the liquid is
therefore saturated with this salt. Its molecular solubility is .00013
(p. 540). In so dilute a solution we may regard the ionization as being
complete. Hence the molar concentrations of Ca" and of CO/' are each
.00013. The solubility product of calcium carbonate is, therefore,
(.00013)2. Now the hydroxide at first furnishes Ca" continuously with
the concentration .0176, and the potassium carbonate, which is all dis-
solved from the beginning, gives CO3" in relatively high concentration.
Thus both the ions of calcium carbonate are present in great excess, and
this compound is precipitated rapidly. Leaving other matters out of
consideration, we should expect the precipitation to go on until one
or other of the ingredients was almost all used up and the product
586 INORGANIC CHEMISTRY
[Ca-] x [CO,"] had fallen to (.00013)2. This presents the theory of
precipitation, in its relation to ionic concentration.
Further reflection shows, however, that, in this particular action,
the hydroxidion continually accumulates in the solution. It begins at
a concentration of .0352 and grows until it equals that of the kalion
(whose concentration is constant throughout and need not be consid-
ered). The final result is a solution of potassium hydroxide. Let us
consider the effect of this upon the dissolving of the calcium
hydroxide :
Ca(OH)2 (solid) <=» Ca(OH)2 (diss'd) <± Ca" + 20H'.
Cl C2 Ca
When the potassium hydroxide has become di-molar, corresponding to
the formation of an 11.2 per cent solution of caustic potash, .67 of its
molecules are ionized, and therefore the concentration of hydroxidion is
2 x .67 or 1.34. In other words, -one of the factors in the solubility
product of calcium hydroxide has become very large. As we have
seen, that product is normally constituted of the factors .0176 x
(.0352)2 and equals .04218. Hence when the hydroxidion reaches .0352
+ 1.34 or 1.3752, both factors must be reduced by an equal amount in
order that the product may retain the value .04218. Thus (.0176 — x)
(1.3752 - 2x)* = .0000218. Solving this cubic equation by trial, we
find x = .0175. This gives the value of the concentration of the cal-
cion, .0176 — x = .0001. The concentration of Ca" is thus less than
.00013, that given by calcium carbonate, and, if the potassium carbonate
has by this time been exhausted, the concentration of Ca" and CO3" will
be smaller than those forming its solubility product, and this salt can
no longer be precipitated. In fact, the precipitation may cease before
so much as 11.2 per cent of potassium hydroxide has been formed.
It thus appears that the action goes forward because calcium car-
bonate is less soluble than calcium hydroxide, but, in presence of a
sufficient concentration of potassium hydroxide, this relation will be
reversed, calcium carbonate may dissolve and calcium hydroxide be
precipitated. When the direction of an action is thus determined by a
difference in solubility amongst slightly soluble substances, ionic con-
centrations may exercise a decisive influence on the result.
This illustrates the rule which describes the conditions under which
solution or precipitation proceeds : If the product of the molar concen-
trations of any pair of ions (one an anion, and one a cation) in a solution
is less than the solubility product for the saturated solution of the
IONIC EQUILIBRIUM, CONSIDERED QUANTITATIVELY 587
substance formed by union of those ions, the substance, if present in
excess, will dissolve. In the case just described this condition is ful-
filled at first for Ca(OH)2 and afterwards for CaC03. Conversely, if
the product of the concentrations of any pair of ions in a solution is
greater than the solubility product for the saturated solution of the sub-
stance, the latter will be precipitated. In the present instance, first
CaC03 and then Ca(OH)2 were the substances in this situation. This
rule defines the conditions for all precipitations of ionogens in chemis-
try, and is therefore continually in application.
Exercises. — 1. The vapor density of sodium peroxide has not
been determined. Why is the formula Na^ nevertheless assigned to
it ? What is its graphic formula ?
2. Construct a scheme of equilibria (p. 371) showing the hydroly-
sis of calcium sulphide. Why does the presence of calcium hydroxide
diminish the tendency to hydrolysis ?
3. In terms of the principle stated on p. 587 explain the precipita-
tion of insoluble hydroxides by the addition of a soluble hydroxide to
a salt. What will be the effect of using an excess either of the salt or
of the soluble hydroxide ?
4. Calculate (p. 583) the effect of adding 2 volumes of a saturated
(3.9 molar) solution of potassium chloride to a saturated (.52 molar)
solution of potassium chlorate. Assume that 1.95 molar KC1 is
.71 ionized.
5. What will be the effect of adding a concentrated solution of
sodium sulphate to a saturated (p. 581) solution of silver sulphate?
CHAPTER XXXV
THE METALS OF THE ALKALINE EARTHS
The Chemical Relations of the Elements. — The metals of this
group, calcium (Ca, at. wt. 40.1), strontium (Sr, at. wt. 87.6), and ba-
rium (JBa, at. wt. 137.4), constitute a typical chemical family both in the
qualitative resemblance to one another of the elements and correspond-
ing compounds and in the quantitative variation in the properties with
increasing atomic weight. The metals themselves displace hydrogen
vigorously from cold water, giving hydroxides. The solutions of these
hydroxides, although dilute, on account of a rather small solubility,
are strongly alkaline in reaction. The high degree of ionization of the
hydroxides recalls the hydroxides of the metals of the alkalies, and
their relative insolubility the hydroxides of the " earths " (q.v.~).
In all their compounds, calcium, strontium, and barium are bivalent.
The hydroxides are formed by union of the oxides with water, and, ex-
cept in the case of barium hydroxide, are easily decomposed again by
heating. The carbonates, when heated, yield the oxide of the metal
and carbon dioxide, barium carbonate being the most difficult to decom-
pose. The nitrates are broken up by heating and yield the oxide of the
metal, nitrogen tetroxide, and oxygen. In these and other respects
the compounds of the metals of the alkaline, earths resemble those of
the heavy metals and differ from those of the metals of the alkalies.
Barium approaches the latter most nearly.
The table on p. 544 shows that the chlorides and nitrates of calcium,
strontium, and barium are all soluble in water, the solubility dimin-
ishing in the order given. The sulphates and hydroxides cover a wide
range from slight solubility to extreme insolubility. Of the sulphates,
2100, 110, and 2.3 parts, respectively, dissolve in one million parts of
water. In the case of the hydroxides the order of magnitude is reversed,
and the corresponding numbers are 200, 630, and 2200. The carbon-
ates are almost as insoluble as is barium sulphate. The new element,
radium (Ra, at. wt. 225), appears to belong to this family (see under
Uranium).
688
THE METALS OF THE ALKALINE EAilTHS 589
CALCIUM.
Occurrence. — The fluoride, and the various forms of the carbon-
ate, sulphate, and phosphate which are found in nature, are described
below. As silicate, calcium occurs, along with other metals, in many
minerals and rocks. It is found also in plants, and its compounds
are important constituents of the bones and shells of animals.
The Metal. — Although the alkali metals can be liberated by heating
the carbonates with carbon, the metals of the present family are not
obtainable by this means. This may be due, in part, to imperfect con-
tact between the materials in consequence of the infusibility of the ox-
ides. Calcium is most easily made by electrolysis of the molten
chloride. A hollow cylinder made of blocks of carbon bolted together
and open above, forms the anode. A rod of copper hanging so that its
end dips into the melt forms the cathode. The melting of the anhy-
drous calcium chloride with which the cylinder is filled is started by
means of a thin rod of carbon laid across from the anode to the cath-
ode. When the heat generated by the passage of the current through
this highly resisting medium has melted a sufficient amount of the salt,
the rod is removed, and the resistance of the fused material suffices to
maintain the temperature. The calcium rises round the cathode and
collects on the surface of the bath. By slowly elevating the copper
cathode, the calcium, which adheres to it, may be drawn out of the
fused mass in the form of a gradually lengthening, irregular rod. The
rod of calcium is kept constantly in contact with the metal which
accumulates on the surface, and thus forms one of the electrodes.
Calcium is a silver-white, crystalline metal (m.-p. 760°, sp. gr. 1.85)
which is a little harder than lead, and can be cut, drawn, and rolled.
It interacts violently with water. When dry and cold it is inactive,
but when heated it unites vigorously with hydrogen, oxygen, the halo-
gens, and nitrogen. It burns in the air, giving a mixture of the oxide
and nitride. The presence of the latter may be shown by the libera-
tion of ammonia when water is brought in contact with the residue :
Ca3N2 + 6H2O -+ 3Ca(OH)2 + 2NH3.
Calcium Hydride. — The hydride CaH2, when formed by direct
union of the constituents, is a white crystalline body. It interacts
vigorously with water, liberating hydrogen.
590 INORGANIC CHEMISTRY
Calcium Chloride. — Calcion is present in small amount in
sea-water, and hence compounds containing calcium chloride, such as
tachydrite CaCl2,MgCl2,12H20, are found in salt deposits. The salt,
for which there is no extensive commercial application, is formed in
large quantities as a by-product in several industrial operations. Thus,
it arises in the liberation of ammonia from ammonium chloride by the
action of lime, in the manufacture of potassium chlorate (p. 273), and
in the Solvay soda process (p. 574). By evaporation of any solution the
hexahydrate of the salt, CaCl2,6H20, is obtained in large, deliquescent,
six-sided prisms. On account of the great concentration of a satu-
rated solution of this compound, the solid and solution do not reach a
condition of equilibrium with ice (cf. p. 164) until the temperature has
fallen to —48°. The freezing mixture must be made with the hydrate,
and not with the anhydrous salt, as the latter gives out much heat in
dissolving. The former, on the other hand, absorbs heat in liquefying,
as all solids do.
There are several other hydrates of calcium chloride containing
less than 6H20, and those containing less water have lower aqueous
tensions (cf. p. 122) than those containing more. By elevating the tem-
perature, however, it is easy to raise the aqueous tension even of the
monohydrate until it exceeds the partial pressure of water vapor in
ordinary moist air, and so to drive out the water. To perform this
rapidly, a temperature of over 200° is required. The dehydrated
calcium chloride forms a porous mass which is used in chemical labo-
ratories for drying gases and liquids. Usually the dehydration is left
incomplete, as, at the temperature required to complete it rapidly, some
interaction with the water occurs (CaCl2 -+- H20 — » CaO + 2HC1), and
a little free alkali is present in the product. When calcium chloride
is used as a drying agent, it is naturally able to reduce the partial
pressure of the water vapor only to the value of the aqueous tension of
the hydrate which is present, and no further. Even at low tempera-
tures the aqueous tensions of hydrates are always perceptible (cf.
p. 121). Concentrated sulphuric acid is a more thorough drying agent
than calcium chloride, and phosphorus pentoxide, whose hydrated form
(metaphosphoric acid) has no observable aqueous tension, is better still.
Calcium chloride forms molecular compounds, not only with water,
but also with ammonia (CaCl2, 8!STH3) and with alcohol. For drying
these substances, therefore, quicklime is employed. Hydrogen sulphide
interacts with the salt, giving hydrogen chloride, which renders the
gas impure. This gas is therefore dried with phosphorus pentoxide.
THE METALS OF THE ALKALINE EARTHS 591
Calcium Fluoride. — This compound occurs in nature as fluorite
or fluor-spar CaF2. It crystallizes in cubes, is insoluble in water, and
when pure is colorless. Natural specimens often possess a green tint
or show a violet fluorescence. It is formed as a precipitate when a
soluble fluoride is added to a solution of a salt of calcium.
Fluorite is used in the etching of glass, as the source of the hydrogen
fluoride (p. 241). It is easily fusible, as its name indicates (I^sit.ftuere,
to flow), and is employed in metallurgical operations, for the purpose
of lowering the melting-point (cf. p. 163) of the slag (g.v.), and so facili-
tating the separation of the latter from the metal.
Calcium Carbonate. — This compound is found very plentifully in
nature. Limestone is a compact, indistinctly crystalline variety, while
marble is a distinctly crystalline form. Chalk* is a deposit consist-
ing of the calcareous parts of minute organisms ; and egg-shells, oyster-
shells, coral, and pearls are other varieties of organic origin. f A
laminated kind of limestone found at Solnhofen is used for lithographic
work. Calcite and Iceland spar (Ger. spalten, to split) are pure crys-
tallized calcium carbonate. The former occurs in flat rhombohedrons,
or in pointed, six-sided crystals (Fig. 52, p. 138) known as scalenohe-
drons ("dog-tooth" spar) belonging to the same system. All the
crystals split with ease parallel to three planes of cleavage, giving rhom-
bohedrons of the shape shown in Fig. 9 (p. 14), but this nearly cubical
form is itself seldom found in nature. An entirely different crystal-
lized variety is known as aragonite. This belongs to the rhombic
system, although complex crystals (" twins ") of hexagonal outline con-
stitute the most familiar specimens. Aragonite, when heated strongly,
resolves itself into a mass of minute crystals of calcite, and the latter
is the more stable form of the substance. When calcium carbonate is
produced by precipitation it is at first amorphous but slowly becomes
crystalline. In cold liquids the resulting crystals are calcite ; but in
warm solutions the less stable form, that of aragonite, is first assumed.
When heated, calcium carbonate dissociates, giving carbon dioxide
and quicklime :
CaC08 <=> CaO + C02.
At ordinary temperatures the decomposition is imperceptible. On the
contrary, atmospheric carbon dioxide, in spite of its very low partial
* Blackboard " crayon " is usually made of gypsum and not of chalk.
t The hard coverings of Crustacea and insects are not made of this substance,
but of an organic material called chitin.
592 INORGANIC CHEMISTRY
pressure, combines with quicklime, giving " air-slaked " lime. As
the temperature rises, however, the tension of carbon dioxide coming
from the carbonate increases, and has a fixed value for each temperature.
If it is continually allowed to escape, so that the maximum pressure is
not reached, the whole of the salt eventually decomposes. If, on the
other hand, the gas is confined, the system reaches a condition of equi-
librium. Attempts to increase the pressure on the gas beyond the
dissociation pressure proper to the existing temperature, result in
recombination. The phenomenon is precisely similar to the dissocia-
tion of barium dioxide (p. 257) and to the evaporation of a liquid
(p. 117). The following data show the actual tensions at various
temperatures :
Temperature 547° 625° 745° 810° 812° 865°
Tension in mm. 27 56 289 678 753 1333
Thus at 812° the pressure almost reaches one atmosphere, and at 865°
it approaches two atmospheres.
Limestone is used in the manufacture of quicklime (q.v.) and of
glass. It is employed largely as a flux in metallurgy, when minerals
rich in silica are brought into fusible form by the production of calcium
silicate (CaSi08). Large amounts also find application as building-stone.
The Phase Rule, a Method of Classifying all Systems in
Equilibrium. — The formal resemblance that we have just shown to
exist between the modes of behavior of a system composed of water
and water-vapor in physical equilibrium, on the one hand, and of a
system made up of calcium oxide, carbon dioxide, and undecornposed
calcium carbonate in chemical equilibrium, on the other, is not a co-
incidence. • A study of all kinds of systems in equilibrium shows that
their different modes of behavior are limited in variety and can be
classified in a very simple way.
The categories used for classification are : (1) the independent
components in the system, and their number; (2) the distinct,
physically separable parts or phases (p. 156) of the system, and their
number; and (3) the conditions — temperature, pressure, and concen-
tration (or volume) — and the degree of variability in the conditions
which is possible without the occurrence of a change in the number of
the phases.
•The mode of employment of these three categories may be illus-
trated in the order of their mention :
THE METALS OF THE ALKALINE EARTHS 593
1. In the water and water-vapor system, water is the only com-
ponent. In the calcium carbonate system, the independent components
are two in number, calcium oxide and carbon dioxide.
2. In the water and water-vapor system there are two phases, the
liquid phase and the vapor phase. In the calcium carbonate system
there are three phases — two solid phases, the carbonate and oxide, and
one gaseous phase, the carbon dioxide. .
3. In the water and water-vapor system either the temperature or
the pressure may be altered, within certain limits, at will. But,
whichever one of these two conditions it be that is thus changed, the
preservation of the two phases will at once require a simultaneous
modification in the other condition, of such a nature as will suit the
new value of the first. Thus, if the pressure upon the vapor is raised,
the vapor phase will be destroyed (p. 117) unless the temperature is
simultaneously elevated to a certain definite point (p. 116). Sim-
ilarly, if the temperature is raised, the liquid phase will all pass into
vapor unless a sufficient increase in the pressure is simultaneously
effected. There is therefore one, and only one degree of variability in
the conditions — the system is univariant. By a study of the cal-
cium carbonate system, as described above, it will be seen that it also
is a univariant system.
A partial generalization of these results leads to the conclusion
that when the number of the phases exceeds the number of the compo-
nents by one, the system is univariant. Additional illustrations are
now required for reaching a still more general statement.
If ice be added to the water and water- y^por system, and the system
be allowed to reach equilibrium with all three phases present, we find
on analyzing as before : one component (water), three phases (solid,
liquid, and gaseous), and no variability in the system. Neither tem-
perature nor pressure may be altered without ensuing disappearance
of one or other of the three phases. This system is therefore
invariant.
It thus appears that with an equal number of components, the
more phases we have, the more restricted are the possibilities of
change in the conditions. When the number of phases equals the
number of components, the system is bi variant ; when the number of
phases exceeds the number of components by one, the system is uni-
variant ; when the number of phases exceeds the number of compo-
nents by two, the system is invariant, and, in general,
Phases + Variable conditions = Components + 2.
594 INORGANIC CHEMISTRY
The law, of which this equation is the most compact expression, is
knawn as the phase rule, and was first formulated by Willard Gibbs, of
Yale University. It applies to physical and chemical equilibria
without distinction, and involves no consideration of molecular or other
hypotheses.
Thus, the formal resemblance between the dissociation phenomena
exhibited by calcium carbonate and other compounds, on the one hand,
and the behavior of a liquid in contact with its vapor on the other, is
due simply to the fact that in each case the number of phases exceeds
the number of components by one. This will be found to hold in all
cases where there is at each temperature a constant dissociation press-
ure. A decomposing hydrate, for example, furnishes such a case.
The system is made up of one gaseous phase (water-vapor) and two
solid phases (the hydrate, and the anhydrous substance or a lower
hydrate). It has three phases and two components (water and the
anhydrous substance), and is, therefore, uni variant.
Again, when we have a sharp transition-point at a fixed temper-
ature, that is, a unique temperature at which alone several different
states of aggregation of a substance can co-exist (p. 115), the system
is always invariant. Thus, ice and water (and vapor) co-exist at the
melting-point of ice : three phases and one component. Again, two
liquid forms of sulphur (p. 369) co-exist with sulphur vapor at 160°.
which is the transition-point : three phases and one component as before.
Still again, at 96° two solid forms of sulphur (p. 368) co-exist with
sulphur vapor. In these cases the change which takes place at the
transition point is purely physical. Analogous cases in which the
change is a chemical one are equally familiar. The decahydrate of
sodium carbonate decomposes (p. 575) above 35.2°. At this temper-
ature the decahydrate and the monohydrate co-exist with the saturated
solution and water-vapor : four phases and two components. The sys-
tem is, therefore, invariant. The cases of gypsum (see p. 603) and
sodium sulphate (p. 576) are similar.
Hard Water. — Calcium carbonate is almost insoluble in water.
In the cold the solubility is a little over 1 part in 100,000 ; in hot
water the solubility is even smaller. Water containing carbonic acid
dissolves it more freely, on account of the formation of the more
soluble calcium bicarbonate (cf. p. 482) :
CaCO8 + H2C03 <=± Ca(HC08)2.
THE METALS OF THE ALKALINE EARTHS 595
A considerable excess of carbonic acid is required, as the action is
markedly reversible. At 15 D a liter of water saturated with carbon
dioxide (at 760 mm. pressure) dissolves 0.385 g. of the carbonate,
whereas a liter of pure water dissolves but .013 g. With a higher
pressure of carbon dioxide still larger amounts may be dissolved. Con-
versely, when the carbon dioxide is driven out by boiling, the carbon-
ate is reprecipitated.
Water containing salts of lime or magnesia in solution is known as
hard water. The carbonate, which may be thrown down by boiling,
gives " temporary hardness," while the sulphate of calcium, since it is
soluble per se and is not affected by boiling, gives "permanent hard-
ness." Soap will not produce a lather with hard 'water until the
calcion has all been precipitated in the form of the calcium salts of
stearic, palmitic, and oleic acids (p. 506). Commercially, the hardness
is therefore measured by the quantity of a standard soap solution
which is just sufficient to produce a persistent lather in a given amount
of the water.
The temporary hardness may be removed by adding to the water a
quantity of slaked lime sufficient to convert the excess of carbonic acid
into calcium carbonate. The permanent hardness may be destroyed by
the addition of sodium carbonate. In both cases the precipitated car-
bonate is allowed to settle or is separated by nitration.
When evaporated in steam-boilers hard water leaves behind a heavy
deposit of boiler-crust containing all the salts formerly in solution.
Similar water, when it dries slowly on the roofs of caverns, gives rise to
stalactites. The drippings form stalagmites on the floors.
Calcium Oxide. — Pure oxide of calcium may be made by ignition
of pure marble or calcite. For commercial purposes limestone is con-
verted into quicklime in kilns. In the United States the " long-flame "
process, in which the kiln is first charged with limestone and a fire is
then kindled in a cavity left at the bottom, is the one most commonly
used. Elsewhere, the limestone and coal are thrown, in alternate
layers, into the kiln, and the products are withdrawn at the bottom. The
latter, the " short-flame " method, demands less fuel, since the operation
is continuous, and the structure is never allowed to cool, but the quick-
lime is mixed with the ash of the coal. In both cases a free passage of
excess of air through the mass is desirable in order that, by reduction
of the partial pressure of the carbon dioxide, the reverse action (p. 257)
may be minimized. The better the ventilation, the lower the tempera-
596 INORGANIC CHEMISTRY
ture at which rapid decomposition can be effected (cf. p. 592). The
lowest temperature that will suffice is employed, as strong heating need-
lessly favors the interaction of the calcium oxide with the clay which
most limestone contains. The product of the interaction with clay,
calcium silicate, is fusible, and by closing many of the pores of the
quicklime renders the subsequent slaking imperfect.
Pure calcium oxide is a white, porous solid. It is infusible even in
the oxyhydrogen flame, but may be melted and boiled in the electric
arc. It is not reducible by sodium, or by carbon excepting at the
temperature of the electric furnace.
Calcium Hydroxide. — When water is poured upon quicklime
it is first absorbed into the pores mechanically. The chemical union
by which the hydroxide is formed :
CaO + H20 <± Ca(OH)2,
proceeds slowly. When it is complete the product is a fine powder
occupying a much larger volume than the original materials. The
action is accompanied by the development of much heat, and during
its progress a part of the water is driven off as steam. The action is
reversible, and at a high temperature the hydroxide can be dehy-
drated. Quicklime from pure limestone slakes easily, and is known as
" fat " lime. That made from material containing clay or magnesium
carbonate is " poor " lime. The latter slakes slowly and often incom-
pletely, and, when used for mortar, does not harden so satisfactorily.
Calcium hydroxide is slightly soluble in water, 600 parts of water
dissolving 1 part of the hydroxide at 18°, and about twice as much
water being required at 100°. The solution, relatively to its concen-
tration, is strongly alkaline. On account of its cheapness, this sub-
stance is used by manufacturers in almost all operations requiring a
base, and it thus occupies the same position amongst bases that sul-
phuric acid does amongst acids. When the presence of much water is
unnecessary or undesirable, a suspension of the solid hydroxide in the
saturated solution ("milk of lime"), or even a paste, is employed.
In such cases, as in the manufacture of caustic alkalies (p. 585), the
action takes place with the part which is at the moment in solution,
and proceeds through the continual readjustment of a complex set of
equilibria. Some of the industries in which caustic lime plays a part
are : the manufacture of alkalies, bleaching powder, and mortar, the
removal of the hair from hides in preparation for tanning, and the
purification of illuminating-gas (p. 513).
THE METALS OF THE ALKALINE EARTHS 597
Mortar and Cement. — Mortar is made by mixing water with
slaked lime and a large proportion of sand. The " hardening " process
consists in an interaction of the carbon dioxide of the air with the
calcium hydroxide :
C02 + Ca(.OH)2 -> CaC03 + H20.
After the superficial parts have been changed, the process goes on
very slowly, and many years are required before the deeper layers have
been transformed. The minute crystals of calcium carbonate which
are formed are interlaced with the sand particles, and a rigid mass is
finally produced. The sand is useful in two ways. In the first place,
it makes the whole material more porous, and so facilitates the
diffusion of the gas into the interior. In the second place, since the
sand is not itself altered, its presence prevents the formation of large
cracks which would otherwise arise from the shrinkage that accom-
panies the formation of the carbonate. The " hardening " does not
begin until the excess of water used in making the mortar has evapo-
rated, and hence ordinary mortar is unsuitable for use in damp places
such as cellars.
Cement is made by strongly heating a mixture of limestone, clay,
and sand, and pulverizing the product. Some natural limestones, con-
taining over 20 per cent of clay, give cements without the addition of
other ingredients. When the cement is mixed with water, it gradually
sets to a solid mass which appears to consist of a mixture of silicates
of calcium and aluminium. The change proceeds throughout the
whole material simultaneously, since it is not dependent on access of
any gas, and not, as in the case of mortar, from the surface inwards.
For this reason the hardening of cement occurs just as well under
water as in any other locality.
Calcium Oxalate. — This salt may be observed under the micro-
scope in the cells of many plants. It appears in the form of needle-
shaped or of granular crystals. Since it is the least soluble salt of cal-
cium, its formation is used as a test for calcium ions. Calcium is es-
timated quantitatively by adding ammonium oxalate to the neutral or
slightly alkaline solution of the calcium salt. The precipitate is sep-
arated by filtration, washed with water, and then heated strongly (ignited)
in a crucible. The product weighed is calcium oxide, CaC,O4— >CaO
+ C02 + CO. More often, perhaps, the oxalate is ignited with sul-
phuric acid, and the calcium weighed as sulphate.
598 INORGANIC CHEMISTRY
Interaction of Insoluble Salts with Acids, Resulting in So-
lution of the Salt. — It was noted above that for the complete precip-
itation of calcium oxalate a neutral solution must be used. The salt
passes into solution when in contact with acids, especially active acids.
Thus, with hydrochloric acid, it gives calcium chloride and oxalic acid,
both of which are soluble :
CaC204 1 + 2HC1 <=» CaCL, + H2C204. ( 1 )
The action of acids upon insoluble salts is so frequently mentioned in
chemistry and is so important a factor in analytical operations that it
demands separate discussion. The present case is a typical one and
may be used as an illustration.
According to the rules already explained (p. 587), calcium oxalate
(or any other salt) is precipitated when the product of the concentra-
tions of the two requisite ions [Ca**] x [C204"] exceeds the solu-
bility product for a saturated solution of calcium oxalate in pure water.
If we assume complete ionization, this value in the present instance
(p. 544) will be [0.0443] x [0.0443]. When, on the contrary, the
product of the concentrations of the two ions falls below the limit-
ing value, a condition which may arise from the removal in some way
either of the Ca" or of the C204" ions, the undissociated molecules
will ionize, and the solid will dissolve to replace them until the ionic
concentrations necessary for equilibrium with the molecules have been
restored or until the whole of the solid present is consumed. Now an
ion may be removed in two ways. It may combine with another ionic
material to form either (1) a still more insoluble substance, or (2) a
substance which although soluble is less highly ionized than the original
salt. The former of these causes determined the solution of calcium hy-
droxide and precipitation of calcium carbonate in the manufacture of po-
tassium hydroxide (p. 586). Here it is the other cause which comes
into operation. The oxalanion from the calcium oxalate combines with
the hydrion of the acid (usually an active one) which has been added,
and forms molecular oxalic acid :
C204" + 2H'^H2C204. (2)
Hence dissociation of the dissolved molecules of calcium oxalate pro-
ceeds, being no longer balanced by encounters and unions of the now
depleted ions, and this dissociation in turn leads to solution of other
molecules from the precipitate.
It will be seen that this second plan for the removal of ions will be
THE METALS OF THE ALKALINE EARTHS 599
successful only when the new acid is a feebly ionized one. Here, to be
specific, the concentration of the C204" in the equilibrium (2) above
must be less than that of the same ion in a saturated calcium oxalate
solution. Now oxalic acid does not belong to the least active class
of acids. Its ionization is chiefly into H" and HC204', as is usual
with dibasic acids, but the latter ion is broken up to an appreciable
extent into H* and C204". There is, however, a decisive factor in the
situation which we have not yet taken into account. The hydrochloric
acid which we used for dissolving the precipitate is a very highly
ionized acid and gives an enormously greater concentration of hydrion
than does oxalic acid. Hence the hjrdrion is in excess in equation (2),
. , [H']2 x [C204"]
and the condition or equilibrium tor oxalic acid, - — ^ L * — = K
[H2O2U4J
in this particular case, will be satisfied by a correspondingly small
concentration of C204". The whole change, therefore, depends for its
accomplishment not only on the mere presence of hydrion, but on the
repression of the ionization of the oxalic acid by the great excess of
hydrion furnished by the active acid that has been used. As a matter
of fact, we find that a weak acid like acetic acid has scarcely any effect
upon a precipitate of calcium oxalate. An acid stronger than oxalic acid
must be employed. The whole scheme of the equilibria is as follows :
2HC1
,.,,.
When excess of an acid sufficiently active to furnish a large concen-
tration of hydrion is employed, the last equilibrium is then driven
forward and the others follow. With addition of a weak acid, only
a slight displacement occurs, and the system comes to rest again when
the molecular oxalic acid has reached a sufficient concentration.
A generalization may be based on these considerations : an insoluble
salt of a given acid will in general interact and dissolve when treated
with a solution containing another acid, provided that the latter acid
is a more highly ionized (more active) one than the former (see below).
This principle furnishes one of the means of measuring the relative
activity of different acids, for the more active the acids the larger will
be the amount of the insoluble salt (say calcium oxalate) which equal
quantities of equivalent solutions will decompose.
Even active acids frequently fail to bring salts of weak acids into
solution, especially when the weak acid is itself present also. Here
the cause lies in the fact that such salts are less soluble than those of
600 INORGANIC CHEMISTRY
the calcium oxalate type, and give so low a concentration of the nega-
tive ion that the utmost repression of the ionization of the correspond-
ing acid does not give a lower value for the concentration of this ion
than does the salt itself. Thus we have seen (p. 375) that even hydro-
chloric acid (dilute) will not dissolve a number of sulphides. For ex-
ample, in the case of cupric sulphide in presence of excess of hydrogen
sulphide, the S" factor in the solubility product [Cu**] x [S"] remains
smaller than that in the scheme defining the hydrogen sulphide equi-
rH"]2 X rS"l
librium = — ^ L — * even when the S" factor in the latter is dimin-
lH2b]
ished in consequence of great addition of hydrion. In this case the
first link in the chain of equilibria :
CuS (solid) <=t CuS (diss'd) <=> Cu" + S
2HC1 ^2Cl'+2
tends so decidedly backward that only the use of concentrated acid
will increase the concentration of the H* to an extent sufficient to
secure any marked advance of the whole action. We must add, there-
fore, to the above rule : provided also that the salt is not one of
extreme insolubility. This point will be illustrated more fully in
connection with the description of individual sulphides (see under
Cadmium).
Illustrations of the application of these generalizations are count-
less. Carbonic acid is made from marble (p. 480), hydrogen sulphide
from ferrous sulphide (p. 371), hydrogen peroxide from sodium perox-
ide (p. 303), and phosphoric acid from calcium phosphate (p. 456). In
each case the acid employed to decompose the salt is more active than
the acid to be liberated. On the other hand, calcium phosphate (except
when freshly precipitated) and calcium oxalate are insoluble in acetic
acid because this acid is weaker than are phosphoric and oxalic acids.
We have thus only to examine the list of acids showing their degrees
of ionization (p. 330) in order to be able to tell which salts, if insoluble
in water, will be dissolved by acids, and, in general, what acids will be
sufficiently active in each case for the purpose. In chemical analysis
we discriminate between salts soluble in water, those soluble in
acetic acid (the insoluble carbonates and some sulphides, for exam-
ple), those requiring active mineral acids for their solution, and those
insoluble in all acids.
The influence of solubility is shown not only by the sulphides, but
also, for example, by the sulphates. Thus, barium sulphate is not ap-
THE METALS OF THE ALKALINE EARTHS 601
preciably dissolved even by the most active acids (p. 390), being a salt
of a rather highly ionized acid, and being itself very insoluble. Yet
calcium sulphate, being much less insoluble (p. 544), is dissolved to a
noticeable extent by the same acids.
Phosphates, oxalates, and other insoluble salts of weak acids, when they have
been dissolved by active acids, are reprecipitated on addition of a basic solution.
The hydroxyl combines with the hydrion of the acid, and so removes the agent by
whose influence solution was effected. Hence, in analysis, the phosphates of ba-
rium, strontium, calcium, and magnesium are dissolved by acids and are repre-
cipitated by ammonium sulphide (an alkaline substance), along with the iron
group.
Precipitation of Insoluble Salts in Presence of Acids. — The
converse of solution, namely, precipitation, depends upon the same
conditions : an insoluble salt which is dissolved by a given acid cannot
be formed by precipitation in th,e presence of this acid. Thus, calcium
oxalate can be precipitated in presence of acetic acid but not in pres-
ence of active mineral acids in ordinary concentrations. Cupric sul-
phide or barium sulphate can be precipitated in presence of any acid,
but ferrous sulphide and calcium carbonate in the absence of acids only.
Dissolving of Insoluble Salts by Solutions of Salts. — The effect
of salt solutions on insoluble salts is seldom very marked. The dissolv-
ing of salts by acids depends, as we have seen, ultimately upon the wide
differences, in degree of ionization which characterize acids. Now, salts
are nearly all highly ionized substances, and therefore the same drastic
effects in the way of solution of insoluble substances are not to be ex-
pected. In making this statement, we are purposely leaving the forma-
tion of complex ions out of consideration since it will be taken up later
(see Copper). So far as the simple crosswise exchange of ions is con-
cerned, and of this nature have been all the cases so far considered, the
effect of adding a solution of a salt to a precipitate is simply to pro-
duce a very limited amount of dissolved un-ionized molecules of two
new salts. This change does not usually consume much of the insol-
uble material. Thus, calcium carbonate with sodium chloride solution
will give a little of the molecular forms of calcium chloride and sodium
carbonate :
Ca" + 2Cl'«=jCaCls and 2Na' + CO8" <=s Na^CO,,
and a very little of the insoluble carbonate will dissolve. It is only
when the insoluble salt is of the less insoluble class that the effects are
602 INORGANIC CHEMISTRY
otherwise than negligible. Thus, calcium sulphate is somewhat more
soluble in many solutions of salts than it is in water.
Calcium Carbide. — The manufacture of this compound has been
described (p. 478), and the formation of acetylene by its interaction
with water has already been discussed (p. 496). The substance was
discovered by Wohler in 1862, was first prepared by the use of elec-
trical heating by Borchers in 1891, and was made on a large scale in
1892 by Wilson, a Canadian engineer.
Bleaching Powder. — This substance (of. p. 266) is manufactured
by conducting chlorine into a box-like structure containing slaked lime
spread upon perforated shelves. When the transformation is complete,
the supply of the gas is shut off, and some lime-dust is blown into the
chamber to absorb the remainder of the free chlorine. The action is
represented by the equation already given, or by the following :
2Ca(OH)2 -f 2CL, -» CaCL, -f Ca(OCl)2 + 2H2O.
While pure lime should thus yield a product containing 49 per cent of
chlorine, in practice the proportion is always less. Good bleaching
powder should contain 36-37 per cent of chlorine.
That bleaching powder is a mixed salt CaCl(ClO) rather than an
equi-molar mixture of calcium chloride and calcium hypochlorite, which
would have the same composition, is rendered probable by the facts
that the material is not deliquescent as is calcium chloride, and that
calcium chloride cannot be dissolved out of it by alcohol.
Bleaching powder is somewhat soluble in water, and in solution the
ions Ca**, Cl', and CIO' are all present. Addition of acids causes the
formation of hydrochloric and hypochlorous acids. The oxidizing
and, incidentally, the bleaching properties (p. 269) of the latter are
characteristic of the acidified liquid. Weak acids like carbonic acid
displace the hypochlorous acid only (cf. p. 267), and hence the dry
powder, when exposed to the air, has the odor of the latter substance
rather than that of chlorine.
The substance is largely used by bleachers (cf. p. 270), and as a
disinfectant to destroy germs of putrefaction and disease.
Calcium Nitrate. — This salt is found in the soil (p. 438), and
may best be prepared in pure form by treating marble with nitric acid
and allowing the product to crystallize from the solution. Calcium
THE METALS OF THE ALKALINE EARTHS 603
nitrate forms several hydrates. The tetrahydrate Ca(N03)2,4H205
which forms transparent monoclinic prisms, is the one deposited at
ordinary temperatures. The anhydrous salt is easily soluble in alco-
hol. It is used in the laboratory for drying nitrogen peroxide. When
heated it decomposes (cf. p. 444), giving nitrogen peroxide, oxygen,
and quicklime.
Calcium Sulphate. — This salt is found in large quantities in
nature. The mineral anhydrite CaSO^ occurs in the salt layers (see
under Manganous sulphate). It contains no water of crystallization,
and its crystals belong to the rhombic system. The dihydrate,
CaS04, 2H20, is more plentiful. In granular masses it constitutes ala-
baster. When perfectly crystallized it is named gypsum or seleuite.
The same hydrate is formed by precipitation from solutions. Its
solubility is about 1 in 500 at 18°. Its solubility varies in an unusual
manner with temperature, increasing slowly to 38° and then falling
off.
When its temperature is raised, the dihydrate quickly shows an
appreciable aqueous tension. After three-fourths of the water has
escaped, a definite hydrate (CaS04)2, H2O remains. This hemihydrate
shows a much smaller tension of water vapor (cf. p. 122).
The transition temperature at which the dihydrate passes sharply into the hemi-
hydrate is 107°. It corresponds to the temperature of 35.2° at which the decahy-
drate of sodium carbonate turns into the monohydrate and water. At 107° both
of the hydrates are in equilibrium with water (p. 594). Naturally this state of
affairs can be realized only in a tube sealed up to prevent the escape of the water.
Plaster of paris is manufactured by heating gypsum until nearly
all the water of hydration has been driven out. When it is mixed
with water, the dihydrate is quickly re-formed and a rigid mass is pro-
duced. If, in course of manufacture, the water is all removed, or the
temperature is allowed to rise much above the most favorable one
(about 125°), the product when mixed with water does not set quickly
and is said to be " dead-burnt." In explanation of this it should be
noted that natural anhydrite combines very slowly with water.
Apparently good plaster of paris must contain some unchanged parti-
cles of the dihydrate which may act as nuclei. They fulfil the same
role as the crystal which is added to a supersaturated solution (p. 159),
without which crystallization may be long delayed or may even fail to
take place. Probably with moderate heating the product is a mixture
of the dihydrate and the hemihydrate with anhydrous salt, while the
604 INORGANIC CHEMISTRY
more rapid decomposition at higher temperatures destroys all of the
first. The former mixture must be an unstable system, and the dihy-
drate loses water to the anhydrous salt. At ordinary temperatures,
however, this transference must be very slow, and hence the property
of setting is not lost by prolonged storage.
That the plaster sets, instead of forming a loose mass of dihydrate,
is due to the fact that the anhydrous salt is more soluble than the
dihydrate, and so a constant solution of the one and deposition of the
other goes on until the hydration is complete :
CaSO, (solid) ^ CaSO (diss'd)
This process results in the formation of an interlaced and coherent
mass of minute crystals.
Plaster of paris is used for making casts and in surgery. The
setting of the material is accompanied by a slight increase in volume,
and hence a very sharp reproduction of all the details in the structure
of the mold is obtained. An " ivory " surface, which makes washing
practicable, is conferred by painting the cast with a solution of paraffin
or stearine in petroleum ether. The waxy material, left by evaporation
of the volatile hydrocarbons, fills the pores and prevents solution and
disintegration of the substance by water. Stucco is made with plaster
of paris and rubble, and is mixed with a solution of size or glue in-
stead of water.
Calcium Sulphide. — This compound is most easily made by
strongly heating pulverized calcium sulphate and charcoal. The
sulphate is reduced :
4C + CaS04 -> CaS + 4CO.
Calcium sulphide is meagerly soluble in water, but is nevertheless
slowly dissolved in consequence of its decomposition by hydrolysis
into calcium hydroxide and calcium hydrosulphide (cf. p. 375). It is
probable that the action would be less nearly complete than it is if
the reverse action were not weakened by the precipitation of the cal-
cixim hydroxide :
2CaS + 2H20 <=> Ca(OH)2 J+ Ca(SH)2.
Since calcium sulphide is thus decomposed by water it cannot be
precipitated from aqueous solution by adding a soluble sulphide, such
THE METALS OF THE ALKALINE EARTHS 605
as ammonium sulphide, to a solution of a salt of calcium. Only the
soluble hydrosulphide can be formed.
Ordinary calcium sulphide, after it has been exposed to sunlight,
usually shines in the dark. Barium sulphide behaves in the same way.
On this account these substances are used in making luminous paint.
They apparently owe this behavior to the presence of traces of com-
pounds of vanadium and bismuth, for the purified substances are not
affected in the same fashion.
Phosphates of Calcium. — The tertiary orthophosphate of cal-
cium Ca^PO^, known as phosphorite, is found in many localities. It
is probably derived from the remains of animals. Guano contains
some of the same substance, along with compounds of nitrogen.
Apatite, 3Ca3(P04)2,CaF9, a double salt with calcium fluoride, is a
common mineral and frequent component of rocks. The orthophos-
phate forms about 83 per cent of bone-ash, and is contained also in
the ashes of plants. It may be precipitated by adding a soluble
phosphate to a solution of a salt of calcium.
Since it is a salt of a weak acid, and belongs to the less insoluble
class of such salts, calcium phosphate is dissolved by dilute mineral
acids (cf. p. 598), the ions HP04" and H2PO/ being formed. When
a base, such as ammonium hydroxide, is added to the solution, the
calcium phosphate is reprecipitated (cf. p. 601).
Calcium phosphate is chiefly used in the manufacture of phos-
phorus and phosphoric acid (p. 456), and as a fertilizer. The supply
of calcium phosphate in the soil arises from the decomposition of rocks
containing phosphates, and is gradually exhausted by the removal
of crops. Bone-ash is sometimes used to make up the deficiency.
It is almost insoluble in water, however, and, although somewhat
less insoluble in natural water containing salts like sodium chloride
(cf. p. 601), is brought into a condition for absorption by the plants
too slowly to be of much service. In consequence of this the " super-
phosphate" (see below) is preferred.
Primary calcium orthophosphate is manufactured in large quan-
tities from phosphorite by the action of sulphuric acid. The uncon-
centrated " chamber acid " is used for this purpose, as water is re-
quired in the resulting action. The amounts of material employed
correspond to the equation :
^ + 2H2S04 + 6HaO -> Ca(H2P04)2,2H20 + 2CaSO4,2H20.
606 INORGANIC CHEMISTRY
As soon as mixture has been effected, the action proceeds with evolu-
tion of heat, and a large cake of the two hydrated salts remains. This
mixture, after being broken up, dried, and packed in bags, is sold as
" superphosphate of lime." The primary phosphate which it contains
is readily soluble in water, and is therefore of great value as a fer-
tilizer.
Calcium Silicate. — Calcium metasilicate CaSi03 forms the min-
eral wollastonite, which is rather scarce, and enters into the composi-
tion of many complex minerals, such as garnet, mica, and the zeolites.
It may be made by precipitation with a solution of sodium metasilicate
(p. 577), or by fusing together powdered quartz and calcium carbonate
or quicklime :
Si02 + CaCO3 -» CaSiO. + CO,.
'2 | Vv't-WVyV-'g 7- \_^ *M K_/ A V^- g ~^ V^V-'n,
Glass. — Common glass is a complex silicate of sodium and cal-
cium, or a homogeneous mixture of the silicates of these metals with
silica. It has a composition represented approximately by the formula
Na20,CaO, 6SiO2, and is made by melting together sodium carbonate,
limestone, and pure sand :
Na2C03 + CaCO3 + 6Si02 -* Na^O, CaO, 6Si02 + 2COa.
For the most fusible glass, a smaller proportion of quartz is employed.
This variety is known as soda-glass, or, from its easy fusibility, as
soft glass. First, the materials are heated to a temperature high
enough to produce chemical action without bringing about complete
melting. This permits the ready escape of the gases. Then the tem-
perature is raised to about 1200° until fusion is complete and all the
bubbles have escaped. Finally, the crucible and its contents are
allowed to cool to 700-800° to allow the latter to acquire the viscosity
required for working.
Plate-glass is made by casting the material in large sheets and
polishing the surfaces until they are plane. Window-glass is prepared
by blowing bulbs of long cylindrical shape, and ripping them down one
side with the help of a diamond. The resulting curved sheets are then
placed on a flat surface in a furnace and are there allowed to open out.
Beads are made, chiefly in Venice, by cutting narrow tubes into very short
sections and rounding the sharp edges by fire. Ordinary apparatus is
made of soft soda-glass, and hence when heated strongly it tends to
soften and also to confer a strong yellow tint (<*/. p. 563) on the flame.
THE METALS OF THE ALKALINE EARTHS 607
In all cases the articles are annealed by being passed slowly through
a special furnace in which their temperature is lowered very grad-
ually. Glass which has been suddenly chilled is in a state of tension
and breaks easily when handled.
Bottles are made with impure materials, and owe their color chiefly
to the silicate of iron which they contain. In making cheap glass,
sodium sulphate is often substituted for the much more expensive
carbonate. In this case powdered charcoal or coal is added to reduce
the sulphate :
C + 2Si0 -> 2^81 0 + C0 + 2S0
2 - 3 2 2.
Soft glass is partially dissolved by water. When powdered glass
is shaken with water the solution soon dissolves enough sodium silicate
to give the alkaline reaction (pink color) with phenolphthalein (cf.
p. 355).
Bohemian, or hard glass, is much more difficult to fuse than soda-
glass, and is also much less soluble in water. It is manufactured by
substituting potassium carbonate for the sodium carbonate, and is used
for making apparatus for special purposes where inf usibility or insolu-
bility are desirable. When lead oxide is employed instead of lime-
stone, a soda-lead glass known as flint glass is produced. This has a
high specific gravity, and a great refracting power for light, and is
employed for making glass ornaments. By the use of grinding
machinery, cut glass is made from it. It is easily fusible, and dissolves
in water like soda-glass.
Colored glass is prepared by adding small amounts of various
oxides to the usual materials. The oxides combine with the silica, and
produce strongly colored silicates. Thus, cobalt oxide gives a blue,
chromium oxide or cupric oxide a green, and uranium a yellow glass.
Cuprous oxide, with a reducing agent, and compounds of gold, give the
free metals, suspended in colloidal solution in the glass, and confer a
deep-red color upon it. Milk-glass contains finely powdered calcium
phosphate in suspension, and white enamels are made by adding
stannic oxide.
Glass is a typical amorphous substance (cf. p. 139). From the facts
that it has no crystalline structure, and that it softens gradually when
warmed, instead of showing a definite melting-point, it is regarded as
a supercooled liquid of extreme viscosity. Most single silicates crys-
tallize easily, and have definite freezing- (and melting-) points. Glass
may be regarded as a solution of several silicates. When kept for a
608 INORGANIC CHEMISTRY
considerable length of time at a temperature insufficient to render it
perfectly fluid, some of its components crystallize out, the glass becomes
opaque, and " devitrification " is said to have occurred. The absence of
such changes in cold glass may be attributed to that general hampering
of all molecular movements and interactions which is characteristic of
low temperatures. The word " crystal " popularly applied to glass is
thus definitely misleading.
Calcion : Analytical Reactions. — Ionic calcium is colorless. It
is bivalent, and combines with negative ions. Many of the resulting
salts are more or less insoluble in water. Upon the insolubility of
the carbonate, phosphate, and oxalate are based tests for calcion in
qualitative analysis (see below). The presence of the element is most
easily recognized by the brick-red color its compounds confer on the
Bunsen flame, and by two bands — a red and a green one — which are
shown by the spectroscope (p. 561).
STRONTIUM.
The compounds of strontium resemble closely those of calcium,
both in physical properties and in chemical behavior.
Occurrence' — The carbonate of strontium SrC08 is found as
strontianite (Strontian, a village in Argyleshire), and is isoniorphous
with aragonite. The sulphate, celestite SrS04, is more plentiful. It
shows rhombic crystals, which are isoniorphous with those of anhydrite,
often have a blue color, and are commonly associated with native sul-
phur in specimens from Sicily. The metal may be isolated by elec-
trolysis of the molten chloride.
Compounds of Strontium. — The compounds are all made from
the natural carbonate or sulphate. The former may be dissolved
directly in acids, and the latter is first reduced by means of carbon to
the sulphide, and then treated with acids.
Strontium chloride, made in one of the above ways, is deposited
from solution as the hexahydrate. The nitrate comes out of hot solu-
tions in octahedrons which are anhydrous. From cold water the tet-
rahydrate is obtained (see under Manganous sulphate). The anhy-
drous salt is. mixed with sulphur, charcoal, and potassium chlorate to
make " red fire." The oxide SrO may be secured by igniting the car-
THE METALS OF THE ALKALINE EARTHS 609
bonate, but on account of the low dissociation tension of the compound
it is obtained with greater difficulty than is calcium oxide from cal-
cium carbonate. It is generally made by heating the nitrate or
hydroxide.
Strontium hydroxide is made by heating the carbonate in a
current of superheated steam:
SrC08 + H20 -> Sr(OH)2 + C02.
This action takes place more easily than does the mere dissociation of
the carbonate, because the formation of the hydroxide liberates energy,
and this partially compensates for the energy which has to be pro-
vided to decompose the carbonate (cf. p. 78). The lowering of the
partial pressure of the carbon dioxide by the steam also contributes to
the result (cf, p. 595).
The hydrate crystallizes from water as Sr(OH)2,8H20, and is
employed in separating crystallizable sugar from molasses. By evapo-
ration of the extract from the sugar-cane or beet-root, as much of the
sugar as possible is first secured by crystallization. Then the molasses
which remains is mixed with a saturated solution of strontium hydrox-
ide. The resulting precipitate of sucrate of strontium, C12H22O11,
SrO, or C^H^On^SrO, is separated by a filter-press, made into a paste
with water, and treated with carbon dioxide. A second filtration parts
the insoluble carbonate of strontium from the solution of sugar, and the
latter is evaporated and allowed to crystallize. Calcium hydroxide,
which gives a tricalcium sucrate, is often employed in the same way.
Strontion is a bivalent ion, and gives insoluble compounds with
carbonanion, sulphanion, and oxalanion. The presence of stronti am is
recognized by the carmine-red color which its compounds give to the
Bunsen flame (see also below). Its spectrum shows several red bands
and a very characteristic blue line.
BARIUM.
The physical and chemical properties of the compounds of barium
recall those of strontium and calcium. All the compounds of barium
which are soluble in water, or can be brought into solution by the weak
acids of the digestive fluids, are poisonous.
Occurrence. — Like strontium, barium is found in the form of
the carbonate, witherite BaC08, which is rhombic and isomorphous with
610 INORGANIC CHEMISTRY
aragonite, and the sulphate BaS04, heavy-spar or barite (Gk. /3apvs,
heavy), which is rhombic and isomorphous with anhydrite. The
specific gravity of the sulphate is 4.5, while the specific gravity of other
compounds of the light metals does not generally exceed 2.5. ' The free
metal, which is silver-white, may be obtained by electrolysis of the
molten chloride.
The compounds are made by treating the natural carbonate with
acids directly, or by first reducing the sulphate with carbon to sulphide,
or converting the carbonate into oxide, and then treating the products
with acids.
Barium Carbonate. — This carbonate demands so high a tempera-
ture (about 1500°) for the attainment of a sufficient dissociation ten-
sion, and is so apt then to be partially protected from decomposition
by the melting of the oxide, that special means is employed for its
decomposition. It is heated with powdered charcoal (cf. p. 485) :
BaC03 + C -> BaO + 2CO.
The precipitated form of the carbonate is made by adding sodium car-
bonate to the aqueous extract from crude barium sulphide (q.v.). The
compound is also obtained by fusing pulverized barite with excess of
sodium carbonate, and dissolving the sodium salts out of the residue.
The Sulphate and Sulphide. — The natural sulphate is the
source of many of the compounds of barium. The precipitated sul-
phate, made by adding sulphuric acid to the aqueous extract from •
barium sulphide, is used in making white paint, in filling paper for
glazed cards, and sometimes as an adulterant of white lead. The salt is
highly insoluble in water and is hardly at all affected by aqueous solu-
tions of chemical agents. It is somewhat soluble in hot, concentrated
sulphuric acid, and the solution yields crystals of a compound
BaS04, H2S04, or Ba(HS04)2. Calcium and strontium sulphates behave
in the same way. All three compounds are decomposed by water, and
give the insoluble sulphates. ,
Barium sulphide, like the sulphides of calcium and strontium
(p. 604), is slightly soluble in water, but slowly passes into solution
owing to hydrolysjis and formation of the hydroxide and hydrosulphide.
It is made by heating the pulverized sulphate with charcoal :
BaSO4 + 40 -* BaS + 4CO.
THE METALS OF THE ALKALINE EARTHS 611
The Chloride and Chlorate. — Barium chloride is generally
manufactured by heating the sulphide with calcium chloride. The
whole treatment of the heavy-spar is carried out in one operation :
BaS04 + 4C + CaCl2 -> 4CO + BaCLj + CaS.
By rapid treatment with water, the chloride can be separated from the
calcium sulphide before much decomposition of the latter (cf. p. 604)
has taken place. Barium chloride crystallizes in rhombic tables as a
dihydrate BaCl2, 2H20. Aside from the difference in composition, this
compound differs from the ordinary hydrated chlorides of calcium and
strontium in being non-hygroscopic and in being capable of dehydration
by heat without the formation of any hydrogen chloride (cf. p. 535).
Barium chlorate is made by treating the precipitated barium car-
bonate with a solution of chloric acid. It is deposited in beautiful
monoclinic crystals, and is used with sulphur and charcoal in the prepa-
ration of " green fire."
The Oxides and Hydroxide. — The oxide of barium is manufac-
tured from the carbonate or sulphide. In the latter case, moist carbon
dioxide is passed over the sulphide, and the resulting carbonate is then
treated with steam. The compound may be obtained in pure form by
heating the nitrate. The oxide unites vigorously with water to form
the hydroxide. When heated in a stream of air or oxygen it gives the
dioxide BaO,. This change and its reversal constitute the basis of
Brin's process for obtaining oxygen from the air (p. 63). To protect
the oxide from conversion into the carbonate and hydrate, which are
not decomposable at the temperature employed, the air must be care-
fully purified from carbon dioxide and moisture.
Barium dioxide, when made by union of oxygen with the monoxide,
is a compact gray mass. A hydrated form is thrown down as a crystal-
line precipitate when hydrogen peroxide solution is added to a solution
of barium hydroxide :
Ba(OH)2 + H20, *=? BaO, J + 2H20.
The crystals have the formula BaO2,8H20. Similar hydrates of the
dioxides of strontium and calcium may be made in the same way.
In all three cases the pure dioxides are obtained as white powders by
removal of the water of hydration by very gentle heating in vamo
(cf. p. 276). The dioxides of strontium and calcium are not formed by
direct union of oxygen with the oxides. Barium dioxide is used in.
the manufacture of hydrogen peroxide (p. 303).
612 INORGANIC CHEMISTRY
Barium hydroxide is the most soluble of the hydroxides of this
group, and gives, therefore, the highest concentration of hydroxidion.
The solution is known as " baryta-water." It is also the most stable of
the three hydroxides, and may be melted without decomposition. It
crystallizes when a warm, saturated solution cools in the form
Ba(OH)2, 8H20. It is much used in quantitative analysis for making
standard alkali-solutions. Solutions of sodium or potassium hydrox-
ide may acquire varying proportions of carbonate by the action of
carbon dioxide from the air, and their action on indicators loses there-
by in sharpness. With barium hydroxide this is impossible, for the
carbonate is insoluble, and is precipitated from the solution.
Barium Nitrate. — This salt is made by the action of nitric acid
on the sulphide, oxide, hydroxide, or carbonate of barium. It crys-
tallizes from aqueous solution without water of hydration.
Analytical Reactions of the Calcium Family. — Barion is a
colorless, bivalent ion. Many of its compounds are insoluble in water,
and the sulphate is insoluble in acids also. The spectrum given by
the salts contains a number of green and orange lines.
In solutions of salts of calcium, strontium, and barium, the ions cal-
cion, strontion, and barion may be distinguished by the fact that
calcium sulphate solution will precipitate the strontium and barium
as sulphates, but will leave salts of calcium unaffected. Similarly,
strontium sulphate solution precipitates barium sulphate, and does
not give any result with salts of the two first. The oxalate of calcium
is precipitated in presence of acetic acid, while the oxalates of
strontium and barium are not (cf. p. 599), and there are other differ-
ences of a like nature in the solubilities of the salts.
Exercises. — 1. Arrange the chromates of the metals of this
family in the order of solubility (p. 644). Compare the solubilities
with those of the carbonates, oxalates, and sulphates of the metals of
the same family.
2. What must be the approximate total molar concentration of the
solution of calcium chloride freezing at — 48° (p. 291) ?
3. What is meant by fluorescence (cf. any book on physics) ?
4. What will be the ratio by volume, at 150°, of the nitrogen per-
oxide and oxygen given off by the decomposition of calcium nitrate ?
THE METALS OF THE ALKALINE EARTHS 613
What would be the nature of the difference between the ratio at 150°
and that at room temperature ?
5. What fact about the heat of solution of gypsum is indicated by
its change of solubility with temperature (p. 260) ?
6. What is the significance of the fact that hydrated barium
chloride gives no hydrogen chloride when heated ?
7. What are the advantages of removing water of hydration in
vacuo (p. 611) ?
8. Explain in terms of the ionic hypothesis the precipitation of
the sulphate of strontium by calcium sulphate solution, and the ab-
sence of precipitation when the latter is added to the solution of a
soluble salt of calcium.
9. Construct a table for the purpose of comparing the properties of
the free elements of this family and also the properties of their corre-
sponding compounds.
10. Are the elements of this family typical metals? If not, in
what respects do they fall short (p. 533) ?
11. What inference do you draw from the fact that the oxalates
of barium and strontium are not precipitated in presence of acetic
acid, while the oxalate of calcium is so precipitated ? Is the infer-
ence confirmed by reference to the data ?
CHAPTER XXXVI
COPPER, SILVER, GOLD
THE three metals of this family, being found free in nature, are
amongst those which were known in early times. They are the metals
universally used for coinage and for ornamental purposes. They are
the three best conductors of electricity (p. 533), and each represents
the maximum of conductivity in the periodic series to which it belongs.
In malleability and ductility silver is intermediate between gold and
copper (p. 531), but in electrical conductivity it exceeds both.
The Chemical Relations of the Copper Family. — Copper (Cu,
at. wt. 63.6), silver (Ag, at. wt. 107.93), and gold (Au, at. wt. 197.2),
occupy the right side in the second column of the table of the periodic
system (p. 411), and the chemical relations (p. 226) of these elements
are in many ways in sharp contrast to those of the alkali metals, their
neighbors, on the left side :
ALKALI METALS. COPPER, SILVER, GOLD.
All univalent and give but one series Cu1 and Cuir : two series. Ag1 : one
of compounds. Halides all soluble series. Au1 and Au111 : two sei'ies.
in water. Halides of univalent series insoluble.
Very active ; rapidly oxidized by air ; Amongst least active metals ; only
displace all other metals from com- copper is oxidized by air ; displaced
bination (E. M. series, p. 362). by most other metals. Hence,
found free in nature (p. 362).
Oxides and hydroxides strongly basic, Oxides and hydroxides feebly basic
and halides not hydrolyzed (p. 534). (except Ag2O) ; halides hydrolyzed
(except Ag-halides). Hence, basic
salts are numerous.
Never found in auion. Give no com- Frequently in anion, e.g., K.Cu(CN)2,
plex cations. K.Ag(CN)2, K.AuO2, K.Au(CN)2,
and in complex cation, e.y.,
Ag(NH3)2.OH and Cu(NH3)4.(OH)2.
On account of their inactivity towards oxygen, and their easy
recovery from combination by means of heat, silver and gold,
together with the platinum family, are known as the " noble metals."
Univalent copper and gold resemble in some ways Hg1 and Tl1,
614
COPPER, SILVEK, GOLD 615
while bivalent copper resembles Zn11, Mn", Fe", and Ni", and trivalent
gold resembles Al111 and Fem. This family is, in fact, not homoge-
neous, and the close relation which, amongst metals, subsists between
valence and chemical properties makes comparisons with elements of
entirely different families often the most suggestive.
COPPER.
Chemical Relations of the Element. — Copper is the first metal-
lic element showing two valences which we have encountered. In such
cases two more or less complete, independent series of salts are known.
These are here distinguished as cuprous (univalent) and cupric (biva-
lent) salts. The methods by which a compoiind of one series may be
converted into the corresponding compound of the other series should
be noted.
The chief cuprous compounds are Cu2O, CuCl, CuBr, Cul, CuCN,
Cu2S. There are no cuprous salts of oxygen acids. The cuprous com-
pound is in each case more stable (p. 119) than the corresponding cupric
compound, and is formed from it either by spontaneous decomposition, as
in the cases of the iodide and cyanide (2CuI2 — » 2CuI + I2), or on heat-
ing. The cuprous halides and cyanide are colorless and insoluble in
water ; but the chloride, being easily oxidized by air to the cupric con-
dition, quickly turns green. The ion Cu* (monocuprion) seems to be
colorless.
The cupric compounds are more numerous, as they include also
salts of oxygen acids, like CuS04, Cu(NO3)2, etc. CuI2 and Cu(dST)2
cannot be obtained in pure form, as they decompose, giving the cuprous
salts. The anhydrous salts are usually colorless or yellow, but the
ion Cu" (dicuprion) is blue, and so, therefore, are the aqueous solu-
tions of the salts. The cupric are more familiar than the cuprous
compounds, since cupric oxide, sulphate, and acetate are the compounds
of copper which most frequently find employment in chemistry and
in the arts. All the soluble salts of copper are poisonous.
In electrolyzing salts of copper, a given amount of electricity will
deposit twice as much copper from a cuprous salt as from a cupric salt
(p. 316), since monocuprion carries only half as great a charge, weight
for weight, as dicupriou.
Writers on chemistry frequently double (Cu2Cl2, etc.) the formulae
of cuprous salts. The molecular weights in organic solvents (cf. p. 292),
however, in many cases accord with the simple formulae. Some higher
616 INORGANIC CHEMISTRY
molecular weights observed in solution and the vapor density of cuprous
chloride (6.6, corresponding nearly to Cu2Cl2) might be regarded as
being due to association (imperfect dissociation, cf, p. 205). The for-
mation of numerous double or complex compounds like HCl,CuCl
(= HCuCl2), which may be regarded as acid salts, however, lends sup-
port to the view that the formulae should be doubled. Inasmuch as the
behavior of the salts is sufficiently well represented by the simple for-
mulae, these are here used throughout (see under Silver, p. 626).
Occurrence. — Copper is found free in the Lake Superior region,
in China, and in Japan. The sulphides, copper pyrites CuFeS2 and
chalcocite Cu2S, are worked in Montana, in southwest England, in
Spain, and in Germany. Malachite, Cu2(OH)2C03 (= Cu(OH)2,CuC03),
and azurite, Cu8(OH)2(COs)2(= Cu(OH)2,2CuC03), both basic carbon-
ates, are mined in Siberia and elsewhere. Cuprite or ruby copper
CujO is also an important ore. The name of the element comes from
the fact that in ancient times copper mines, long since worked out, ex-
isted in Cyprus. The element is found in the feathers of some birds,
in the haemocyanin of the blood of the cuttle-fish, which is blue when
arterial and colorless when venous, and elsewhere in living organisms.
Extraction from Ores. — For isolating native copper it is only
necessary to separate the metal, by grinding and washing, from the
rock through which it ramifies, and to melt the almost pure powder of
copper with a flux (p. 540). The carbonate and oxide ores require
coal, in addition, for the removal of the oxygen.
The liberation of copper from the sulphide ores is difficult, and often
involves very elaborate schemes of treatment. This arises from two
causes. Other metals, such as iron and zinc, unite with oxygen more
readily than with sulphur, and hence the sulphides are easily converted
into oxides by roasting with free access of air ; with copper it is just the
reverse. Thus, even great excess of air and repeated roasting produce
only a gradual diminution in the amount of sulphur in the mass. Then,
many copper ores contain a large amount of the sulphides of iron, and
these have to be removed by conversion into oxide (by roasting) and
then into silicate (with sand). The silicate forms a flux, and separates
itself from the molten mixture of copper and copper sulphide
("matte"). In Montana it is found possible to abbreviate the treat-
ment. The ore is first roasted until partially oxidized. It is then
melted in a cupola or a reverberating furnace, and placed in large iron
COPPER, SILVER, GOLD 617
vessels like Bessemer converters (q.v.~) provided with a lining rich in
silica. A blast of air mixed with sand is now blown through the
mass. The iron is completely oxidized to FeO and made into silicate,
the sulphur escapes as sulphur dioxide, and arsenic and lead are like-
wise removed by this treatment. The silicate of iron floats as a slag
upon the copper when the contents of the converter are poured out.
The resulting copper is pure enough to be cast in large plates and
purified by electrolysis (see below).
Wet processes are used for poor ores. In one of these the ore is
roasted with salt. The copper is thus converted into cupric chloride,
and can be dissolved away from the oxide of iron and other materials
by means of water. Any traces of silver which may have been present
pass also into solution (as AgCl), and are precipitated by addition of
potassium iodide. The copper is then displaced by means of scrap
iron, and forms a brown sludge (Cu" + Fef — > Fe" + Cu J).
The properties of copper are seriously affected by small amounts of
impurities, such as cuprous oxide or sulphide, which are soluble in the
molten metal. Then, too, the difficulties in the way of its preparation
are aggravated by the high standard of purity demanded in order that it
may have the maximum tenacity, ductility, and conductivity which its
various applications require. Hence a large proportion of the copper
on the market is purified by electrolysis, a method which meets the
case by giving a copper in which foreign materials can be shown to be
present only by the most refined tests. In this method of refining the
metal, thin sheets of copper, coated with graphite to permit easy re-
moval of the deposit, form the cathodes, and thick plates of copper the
anodes. These are suspended alternately and close together in large
troughs filled with cupric sulphate solution. The cathodes are all
connected with the negative wire of the dynamo, and the anodes with
the positive one. The Cu" is attracted to the cathodes and is deposited
upon them. The SO/' migrates towards the anodes, where copper
from the thick plate becomes ionized in equivalent amount. The stock
of cupric sulphate thus remains the same, and the practical effect of
the electrolysis is to carry copper across from one plate to the other
(cf. p. 325). The cathodes are removed from time to time, and the
deposit of copper is stripped from their surface. Fresh anodes are
substituted when the old ones are eaten away. Since there is no polar-
ization, a current of less than 0.5 volt suffices. The copper is deposited
in pure form, although an impure anode is employed, because metals
like gold, silver, and antimony, which succeed copper in the electro-
618 INORGANIC CHEMISTRY
motive series (p. 362), and compounds like cuprous sulphide, do not
become ionized. They fall to the bottom of the tank as a fine powder.
Similarly, metals like zinc, which displace copper, although they are
dissolved, are not again deposited. This will be understood when it is
considered that, if they were to be deposited, they would displace
copper from the solution and dissolve. Since the copper, as deposited
on the cathodes, is crystalline and porous, it is afterwards melted and
cast into blocks or bars.
In 1900 the United States produced 60 per cent of the world's
copper, Great Britain 16 per cent, and Germany 6 per cent. The pro-
portions of the whole consumed in each of these countries were 33, 22,
and 22 per cent, respectively.
Physical Properties. — Copper is red by reflected and greenish by
transmitted light. Native copper shows crystals of the regular system
(p. 530). It melts at 1057°, and therefore much more easily than pure
iron (1800°). When steel draw-plates are used it can be drawn into
wire with a diameter of only 2 mm., and by means of plates provided
with perforated diamonds the diameter of the wire can be reduced to
.03 mm. (1 kilometer weighs only 7 g.). The metal after drawing is
more tenacious but conducts electricity less well.
Chemical Properties. — In dry oxygen, copper does not rust.
In moist oxygen a thin film of cuprous oxide is formed, and in ordinary
air a green basic carbonate (not verdigris, q.v.). It does not decom-
pose water at any temperature or displace hydrogen from dilute acids
(p. 362). On the other hand, hydrogen, absorbed in platinum or even
in charcoal, liberates copper (Cu " + H2 — » Cu + 2H') when immersed
in solutions of copper salts. The metal attacks oxygen acids (pp.
379, 446), however. Again, acids like hydrochloric acid, in conjunc-
tion with oxygen from the air, do act slowly upon copper (2Cu + 4HC1
+ 02 — » 2CuCl, -f- 2H20). This sort of simultaneous action of two
agents is frequently used, as in making silicon tetrachloride (p. 520).
In a similar way sea-water and air slowly corrode the copper sheath-
ings of ships, giving a basic chloride, Cu4(OH)6Cl2, H,0(= 3Cu(OH)2,
CuCl2,H2O), which is found in nature as atakamite.
On account of its resistance to the action of acids, copper is used
for many kinds of vessels, for covering roofs and ships' bottoms, and
for coins. It furnishes also electrotype reproductions of medals, of
engraved plates, of type, etc. For this purpose a cast of the object
COPPER, SILVER, GOLD 619
is first made in gutta percha, plaster of paris, or wax. This is then
coated with graphite to give it a conducting surface, and receives an
electrolytic deposit of copper. Great quantities of the metal are used
iu electrical plants and appliances.
Alloys. — The qualities of copper are modified for special purposes
by alloying it with other metals. Brass contains 18-40 per cent of
zinc, and melts at a lower temperature (p. 532) than does copper. A
variety with little zinc is beaten into thin sheets, giving Dutch-metal
(" gold-leaf "). Bronze contains 3-8 per cent of tin, 11 or more per
cent of zinc, and some lead. It was used for making weapons and
tools before means of hardening iron were known, and later, on account
of its fusibility, continued to be employed for castings until displaced
largely by cast-iron (discovered in the eighteenth century). For works
of art it is preferred to copper because of its fusibility, its color, and
its more rapid acquirement of a much prized " patina " due to surface
corrosion. Artificial "bronzing" of brass is effected by applying a
solution of arsenious oxide in hydrochloric acid (AsCl3). The zinc
displaces some arsenic, which combines with the copper. Brass instru-
ments are " bronzed " by means of a dilute solution of chloroplatinic
•acid (q.v.), from which the zinc displaces platinum. Gun-metal con-
tains 10 per cent, and bell-metal 25 per cent of tin. German silver
contains 19-44 per cent of zinc and 6-22 per cent of nickel, and shows
none of the color of copper. Aluminium-bronze contains 5-10 per cent
of aluminium, and resembles gold in color. When it contains some
iron it can be worked at a red heat, but not welded. Silicon-bronze
contains not more than 5 per cent of silicon, and is made by adding
silicide of copper (made in the electric furnace, p. 457) to copper. It
has usually only 60 per cent of the conductivity of pure copper, but is
nearly twice as tenacious, and is used for telephone and over-head
electric wires. Phosphor-bronze contains copper and tin (100 : 9) with
£-1 part of phosphorus, and is employed for certain parts of machines.
Ships' propellers are made of manganese-bronze (30 per cent manganese).
Cupric Chloride. — This compound is made by union of copper
and chlorine, by treating the hydrate or carbonate with hydrochloric
acid, or by heating copper with hydrochloric acid and some nitric acid
the latter being used simply as an oxidizing agent (Cu + 2HC1 + 0 —>
CuCL, + H20). The .blue crystals of a hydrate, CuCl^H/), are de-
posited by the solution. The anhydrous salt is yellow. Dilute solu-
620 INORGANIC CHEMISTRY
tions are blue, the color of dicuprion, but concentrated solutions are
green on account of the presence of the yellow molecules (p. 335).
The aqueous solution is acid in reaction (p. 344). When excess of
ammonium hydroxide is added to the solution, the basic chloride, cupric
oxychloride Cu4(OH)6CLj (p. 618), which is at first precipitated, redis-
solves, and a deep-blue solution is obtained. This on evaporation yields
deep-blue crystals of hydrated ammonio-cupric chloride Cu(NHs)4.
CLj, 1^0. The deep-blue color of the solution, which is given by all
cupric salts, is that of the Cu(NH3)4" ion. The dry salt also absorbs
ammonia, giving CuCLj, 6NH3. This and the preceding compound have
an appreciable tension of ammonia, and when warmed leave CuC^, 2NH8.
Eeduction of pressure or rise of temperature results in the final loss
of all the ammonia (cf. p. 122).
Cuprous Chloride. — This salt is formed when cupric chloride is
heated (2CuCLj «z± 2CuCl + CLj). It may be made by adding hydro-
chloric acid to- cupric chloride solution, and boiling the mixture with
copper turnings :
CuCL, + Cu -> 2CuCl or Cu" + Cu _» 2Cu*.
The solution contains compounds of cuprous chloride with hydrogen
chloride HCl,CuCl or HCuCLj and H2CuClg, which are decomposed when
water is added. The cuprous chloride is insoluble in water, and forms
a white crystalline precipitate.
Cuprous chloride is hydrolyzed quickly by hot water, giving,
finally, red, hydrated cuprous oxide, Cu20. When dry it is not
affected by light, but in the moist state becomes violet and, finally,
nearly black. The action is said to be 2CuCl — > CuCl2 + Cu. In
moist air it turns green, and is oxidized to cupric oxychloride (p. 618).
It is dissolved by hydrochloric acid, giving the colorless complex acids
HCuCl2 and JELjCuCls. The solution is oxidized by the air, turning
first brown and then green, and finally depositing the cupric oxychlo-
ride. Cuprous chloride also dissolves in ammonium hydroxide, giving
Cu(NH3)2.Cl, the ion Cu(NH3)2* being colorless. The solution is
quickly oxidized by the air, turns deep-blue, and then contains
Cu(NH3)4". The solution of cuprous chloride in hydrochloric acid is
used for absorbing carbon monoxide from gaseous mixtures. A crys-
talline compound (CuCl, CO, 2H20 ?) has been isolated from the solu-
tion.
COPPER, SILVER, GOLD 621
The Bromides and Iodides of Copper. — By treatment of cu-
pric oxide with hydrobromic acid and slow evaporation of the solution,
jet-black crystals of anhydrous cupric bromide (CuBr2) are obtained.
A concentrated aqueous solution is deep-brown in color, and the grad-
ual ionization of the molecules as the solution is diluted is well shown by
this salt (p. 335). The ionization is here accompanied by evolution of
heat (p. 330), as it is also in the cases of cupric chloride and cupric sul-
phate, and in the ionized condition the substances contain less available
energy than in the molecular. In these cases, therefore, when the tem-
perature is raised the ionization diminishes (p. 260). It will be remem-
bered that ordinary thermal dissociation invariably increases with rise
in temperature, and is always accompanied by absorption of heat.
Ionic dissociation, however, may be either endothermal or exothermal.
When cupric bromide is heated, bromine is given off, and cuprous
bromide CuBr remains.
Cupric iodide appears to be unstable at ordinary temperatures.
When a soluble iodide is added to a cupric salt, a white precipitate of
cuprous iodide and free iodine are obtained :
The iodine may be dissolved in excess of the soluble iodide (p. 235), or
reduced to hydrogen iodide with sulphurous acid (p. 394).
The Solution of Insoluble Salts when Complex Ions are
Formed. — The solution of an insoluble salt like cuprous chloride by
hydrochloric acid or ammonium hydroxide is typical of a great variety
of actions of which we here meet one of the first examples (cf. p. 363).
Since a salt is normally less soluble in an acid having the same
anion (p. 581), the dissolving of cuprous chloride in hydrochloric acid
requires a special explanation, namely, the fact that here two complex
acids H.CuCl2 and H2.CuCl8 are formed. The chloridion of the hy-
drogen chloride must indeed tend to repress the ionization of the dis-
solved part of the cuprous chloride, so that a smaller concentration of
Cu* remains. But the complex negative ion CuCL/ (or CuCl8") which is
formed, is very little dissociated, and gives a still smaller concentration
of Cu* (CuCl/ £± Cu* + 2C1'). Thus this complex ion is formed at the
expense of the Cu* of the insoluble cuprous chloride, and the latter
goes into solution progressively in the effort to restore the balance :
CuCl (solid) <=; CuCl (diss'd) <=± Cl' + Cu* )
2HC1 <=±2H* + 2C^ V
622 INORGANIC CHEMISTRY
The same exact laws of equilibrium used in discussing the dissolving
of salts by acids with a different anion (p. 599) may be applied to the
whole procedure.
Similar behavior is shown by the cyanides of copper, silver, iron,
etc. (q-v.), of which many complex compounds are known.
The dissolving of cuprous chloride by the free ammonia of ammo-
nium hydroxide is explained in the same way. The only difference
is that here the copper is in the complex positive ion. The ion
Cu(NH8)2* gives little Cu* — less than does cuprous chloride, in spite of
the insolubility of the latter. Hence the salt passes into solution
until the ion-product [Cu*] x [Cl'], with continually increasing [Cl'],
reaches its normal value or until the solid is exhausted.
The deep-blue colored ion Cu(NH3)4" given by cupric chloride and
other cupric salts is also very little ionized. Hence ammonium hy-
droxide dissolves all the insoluble cupric compounds save only cupric
sulphide, which is the most insoluble of all — - that is, the one giving
the smallest concentration of dicuprion. Conversely, the sulphide is
the only insoluble compound of copper which can be precipitated from
ammoniacal solution. Zinc and other more active metals, however,
slowly precipitate metallic copper, thereby showing that some dicup-
rion is present.
Cuprous Oxide. — This oxide is red in color, and natural specimens
show octahedral forms. It is produced by oxidation of finely divided
copper at a gentle heat, or by the addition of bases to cuprous chloride,
and is best made by the action of glucose on cupric hydroxide. The
latter is reduced by the former, and the resulting hydrated cuprous
oxide forms a pale-brown precipitate which quickly becomes bright
red. The simple hydrate, CuOH, is unknown, but the above mentioned
precipitate has approximately the composition 4Cu20,H20, and yields
Cu2O when heated.
Cuprous oxide is acted upon by hydrochloric acid, giving cuprous
chloride, or rather HCuCl^. It also dissolves in ammonium hydroxide,
giving, probably, Cu(KH3)2.OH, which is colorless. With dilute
oxygen acids part of it is oxidized, giving the cupric salt, and part is
reduced to metallic copper :
Cu20 + H2S04 -> CuS04 + Cu + H20.
Cupric Oxide and Hydroxide. — Cupric oxide is a black sub-
stance formed by heating copper in a stream of oxygen or by igniting
COPPER, SILVER, GOLD 623
the nitrate, carbonate, or hydroxide. It absorbs moisture from the air,
although it is not soluble in water. When heated strongly it loses some
oxygen, and is partly reduced to cuprous oxide. Its chief use is in the
analysis of compounds of carbon. When heated with the latter, it
Dxidizes the hydrogen to water, and the carbon to carbon dioxide. The
operation is performed in a tube through which passes a stream of
oxygen, and the products are caxight in glass vessels containing calcium
chloride and potassium hydroxide, respectively.
Cupric hydroxide is precipitated as a gelatinous substance by
addition of sodium or potassium hydroxide to a solution of a cupric
salt (Cu- + 20H' -> Cu(OH)2). When the mixture is boiled, the
hydroxide loses water and forms a black hydrated cupric oxide
(Cu(OH)2, 2CuO ?). The hydroxide is soluble in ammonium hydroxide,
with formation of the compound Cu(NH3)4.(OH)2, which imparts a
deep-blue color to the solution. Various forms of cellulose, such as
filter paper and cotton, dissolve in this solution, and are reprecipitated
when the ammonium hydroxide is neutralized with acids. Cupric
hydroxide dissolves also in a solution of sodium tartrate (Na2.C4H2O4
(OH)2), giving a deep-blue liquid (practically "Fehling's solution").
In this action, it enters into the negative ion, as is shown by electrolysis,
interacting ' apparently with the hydroxyl groups of the tartranion.
In this condition it is reduced by sugars (see above) with especial ease,
and is in this form used as a test for them.
Cupric Nitrate. — The nitrate is made by treating cupric oxide
or copper with nitric acid (p. 446), and is obtained from the solution
as a deliquescent, crystalline hydrate. The hexahydrate is secured
at temperatures below 24.5°, its transition point (p. 594), and the tri-
kydrate from 24.5° up to 114.5° (its transition point; see under Man-
»anous sulphate). When dehydrated at 65° the salt is partly
hydrolyzed, and a basic nitrate Cu4(OH)6(N08)2 remains.
Carbonate of Copper. — No normal carbonate (CuC08) can be
obtained. A basic carbonate (malachite) is found in nature, and is
precipitated by adding soluble carbonates to cupric salts :
2CuS04 + 2Na2C03 + H20 -» Cu2(OH),C03 + 2^80, + C02.
Presumably, the carbonate, if formed, would be hydrolyzed by water.
624 INORGANIC CHEMISTRY
Cyanides of Copper. — When potassium cyanide is added to a
solution of a cupric salt, cupric cyanide is precipitated. This is not
stable, however, and gives off cyanogen, leaving cuprous cyanide :
2Cu(CN)2 -» 2CuCN + C2N2.
Cuprous cyanide is insoluble in water, but interacts with an excess of
potassium cyanide solution, producing a colorless liquid, from which
KCN,CuCN, or K.Cu(CN)2, potassium cuprocyanide, may be obtained
in colorless crystals. The complex anion Cu(CN)/ is so little ionized
to Cu* and 2CN' that all insoluble copper compounds, including cupric
sulphide, are dissolved by potassium cyanide ; and none of them can
be precipitated from the solution. Zinc is actually unable to displace
copper from such a solution. The cause of the solution of the salts
is the same as when the complex ions Cu(NH8)2*, Cu(NH8)4", and
are formed (p. 621).
Cupric Acetate. — By the oxidation of plates of copper, sep-
arated by cloths saturated with acetic acid (vinegar), a basic acetate
of copper (verdigris) is obtained :
6Cu + 8HC2H302 + 3O2 -> 2Cu3(OH)2(C2H302)4 + 2H2O.
It is used in manufacturing green paint, is insoluble in water, and is un-
affected by light. It dissolves in acetic acid, and green crystals of the
normal acetate Cu(C2HsO2)2,H20 are obtained from th£ solution. The
basic acetate is used in preparing parts green. A hot solution of
arsenious acid (H3As08) is mixed with a paste of verdigris and a little
acetic acid and boiled. A precipitate of paris green Cu(C2H302)2,(<u3
As206, which has a unique light-green color, is thrown down. On ac-
count of their poisonous nature, this compound and Scheele's green
(CuHAsO3) are little used as pigments. The former is chiefly made
for use in the extermination of potato-beetles and other insects and
employment in the destruction of parasitic fungi.
Cupric Sulphate. — This salt is obtained by heating copper in a
furnace with sulphur, and admitting air to oxidize the cuprous sul-
phide. The mixture of cupric sulphate and cupric oxide which is
formed is treated with sulphuric acid. The salt is also made by allow-
ing dilute sulphuric acid to trickle over granulated copper while air
has free access to the material (2Cu + 2H2S04 -j- O2 —> 2CuS.04 +
COPPER, SILVER, GOLD 625
2H20). When concentrated and at a high temperature, sulphuric
acid will itself act as the oxidizing agent (cf. p. 379).
Cupric sulphate crystallizes as pentahydrate CuS04,5H20 in blue
asymmetric crystals (Fig. 41, p. 120), and in this form is called blue-stone
or blue vitriol. The dissociation of this hydrate has been discussed
on page 122. The aqueous solution has an acid reaction (p. 344).
The anhydrous salt is white, and can be crystallized in thin needles
(rhombic system ?) from solution in hot, concentrated sulphuric acid
(cf. pp. 120-123). Cupric sulphate is employed for making other com-
pounds of copper, in copper-plating (p. 618), in batteries, as a mordant
in dyeing (q.v.) and calico-printing, and, as a germicide and insecti-
cide, for spraying plants.
When ammonium hydroxide is added to cupric sulphate solution,
a pale-green basic salt (Cu4(OH)6S04 ?) is first precipitated. With
excess of the hydroxide the blue Cu(NH3)4" ion (p. 620) is formed, and
crystals of ammonio-cupric sulphate Cu(NH3)4. S04,H20 can be obtained
from the solution. This compound easily loses water and ammonia
(by stages), leaving successively CuS04, 2NH3 and CuS04, NH3. Cupric
sulphate also combines with potassium and ammonium sulphates,
giving double salts of the form CuS04,K2S04,6H20, which are de-
posited in large, inonosymmetric crystals from the mixed solutions
(see Zinc sulphate).
The Sulphides of Copper. — Cuprous sulphide Cu2S occurs in
nature in rhombic crystals of a gray, metallic appearance. It is made
by heating cupric sulphide, a stream of hydrogen gas being used to
assist the removal of the excess of sulphur.
Cupric sulphide is deposited as a black precipitate when hydrogen
sulphide is led through a solution of a cupric salt. Made in this way,
it is always partly decomposed into Cu2S + S. By cautiously treating
copper with excess of sulphur at 114° it may be obtained as a blue
crystalline solid. At higher temperatures it gives off sulphur.
Analytical Reactions of Compounds of Copper. — The ion of
ordinary cupric salts, dicuprion Cu", is blue, and that of cuprous
salts, monocuprion Cu*, is colorless. Cuprous solutions, however, are
easily oxidized by the air and become blue. In solutions containing
dicuprion, hydrogen sulphide precipitates cupric sulphide, even in
presence of acids (p. 600). Bases throw down the blue hydroxide,
and carbonates precipitate a green basic salt (p. 623). Potassium fer-
626 INORGANIC CHEMISTRY
rocyanide gives the brown, gelatinous cupric ferrocyanide (2Cu.S04 +
K4.Fe(CN)6 <± Cu,.Fe(CN)6 j. + 2K2S04). A very characteristic test is
the formation of the deep-blue Cu(NH3)4"ion with excess of ammo-
nium hydroxide. This solution itself gives a precipitate with hydro-
gen sulphide only. Solutions of complex cuprous and cupric cyanides
such as K.Cu(CISr)2 and K2.Cu(CN)4 are colorless, and do not respond to
any of the above tests. With microcosmic salt or borax (pp. 468, 528),
copper compounds form a bead which is green in the oxidizing part of
the flame and becomes red and opaque (liberation of copper) in the
reducing flame.
SILVER.
Chemical Relations of the Element. — This element presents a
curious assortment of chemical properties. It differs from copper in
having a strongly basic oxide, in giving salts with active acids which
are not hydrolyzed by water, and in forming neutral rather than basic
salts. In these respects it approaches the metals of the alkalies and
alkaline earths. It resembles copper in entering into complex com-
pounds, and in giving insoluble halides like the cuprous halides. It
differs from both copper and the metals of the alkalies, and resembles
gold and platinum, in that its oxide is easily decomposed by heat,
with formation of the free metal, and in the low position it occupies in
the electromotive series and the consequent slight chemical activity of
the free metal.
The salts are always represented by the simplest formula, AgCl,
etc., although in organic solvents greater tendencies to polymerization
are observed than in the case of the cuprous compounds (p. 615).
Occurrence. — Native silver, sometimes found in large masses, al-
though more usually scattered through a rocky matrix, contains vary-
ing amounts of gold and copper. Native copper always contains
dissolved silver. Sulphide of silver (Ag2S) occurs alone and dissolved
in galenite (PbS), with which it is isomorphous. Smaller amounts of
the metal are obtained from pyrargyrite Ag3SbS3 and proustite Ag3AsS8,
which are silver sulphantimonite and sulpharsenite respectively, and
from horn-silver AgCl.
Metallurgy. — The silver contained free, or as sulphide, in ores
of copper and lead, is found in the free state dissolved in the metals
extracted from these ores, and is secured by refining them. In the elec-
trolytic refining of copper, silver is obtained from the mud deposited in
COPPER, SILVER, GOLD 627
the baths (p. 617). The proportion present in lead is usually small.
Formerly the Pattinson desilverizing process was largely employed.
In it the metallic lead is melted in iron vessels, and the crystals of
lead, deposited as the metal slowly loses heat, are raked out. These
consist at first of pure lead (cf. p. 294). When the remaining liquid
becomes saturated with silver it begins to deposit lead and silver to-
gether. At this point the residue is placed in a hollow, lined with
bone-ash, forming part of a reverberatory furnace (Fig. 100, p. 572), and
heated strongly while a blast of air passes over its surface. In this
process, called " cupellation," the lead is converted into litharge (PbO),
which, driven by the air, flows in molten condition over the edge of
the cupel. When the last trace of lead is gone, the shining surface of
the pure silver "flashes" into view (cf. p. 547). Parke's process,
which has superseded the above, takes advantage of the fact that mol-
ten zinc and lead are practically insoluble in one another, while silver
is much more soluble in zinc than in lead. Lead dissolves 1.6 per
cent of zinc, and zinc 1.2 per cent of lead. The principle is the same
as in the removal of iodine from water by ether (p. 155). The lead is
melted and thoroughly mixed by machinery with a small proportion of
zinc. After a short time the zinc floats to the top, carrying with it in
solution almost all of the silver, and solidifies at a temperature at
which the lead is still molten. The zinc-silver alloy is skimmed off,
and heated moderately in a furnace to permit the adhering lead to
drain away. The zinc is finally distilled off in clay retorts, and the
lead remaining with the silver is removed by cupellation.
Ores of silver which do not contain much or any lead are often
smelted with lead ores, and the product is treated as described above, but
many other processes are in use. Thus, sulphide ores are sometimes
roasted until the iron and part of the copper are converted into oxide
while the rest of the copper and all the silver remain as sulphate.
The metal is secured by extracting the mass with water and precipitat-
ing the silver by means of copper (p. 362) . Some ores are roasted with
salt, and the resulting chloride of silver is dissolved out with sodium
thiosulphate, or even strong brine. In Mexico the "patio" process has
been in use since 1557. The sulphide is converted into chloride by
the action of cupric chloride. Metallic mercury displaces the silver
(AgCl + Hg — » HgCl + Ag), and, being present in excess, dissolves
it. The treatment occupies several weeks, and much mercury is con-
sumed. The amalgam is finally secured by " washing," and the mercury
is separated from the silver by distillation.
628 INORGANIC CHEMISTRY
In 1899 the production of silver in the United States was 2915 tons,
in Mexico 1418 tons, in Europe 1138 tons. During the first half of
the nineteenth century the total world's output averaged only 643 tons
per year. Up to 1870 a gram of gold could buy 15.5 g. of silver. Now
that the production has reached 6000 tons, the same amount of gold
purchases about 35 g.
Physical Properties, — Pure silver is almost perfectly white. It
melts at 960°. Its ductility is so great that wires can b.e drawn of
such fineness that 2 kilometers of the finest wire weigh only about 1 g.
In the molten condition it absorbs mechanically about twenty-two
times its own volume of oxygen, but gives up almost all of this as
it solidifies. Fantastically irregular masses result from- the " sprout-
ing " or " spitting " which accompanies the escape of the gas.
By addition of ferrous citrate to silver nitrate, a red solution and
lilac precipitate of free silver can be made. The latter, after washing
with ammonium nitrate solution, gives a red solution in water. Other
solutions of colloidal (cf. p. 523) silver showing a variety of colors
have been prepared by Cary Lea. Such colloidal solutions of metals
are formed also by passing an electrical discharge between wires of sil-
ver, gold, or platinum held under water.
Silver is alloyed with copper to render it harder. The silver coin-
age of the United States and the continent of Europe has a " fine-
ness of 900 " (900 parts of silver in 1000), and that of Great Britain
925. Silver ornaments have a fineness of 800 or more. A superficial
layer of almost pure-white silver is produced by heating the object in
the air and dissolving out the cupric oxide thus formed with dilute
sulphuric acid. The surface of the products, if not subsequently bur-
nished, is "frosted." "Oxidized silver" is made by dipping objects
made of the metal in a solution of potassium hydrogen sulphide,
whereby a thin film of silver sulphide is produced.
Chemical Properties. — Silver does not combine with oxygen,
either in the cold or when heated. It does not ordinarily displace
hydrogen from aqueous solutions of acids, but its tendency to form the
sulphide is so great that it decomposes hydrogen sulphide and alkali
sulphides (cf. p. 391). It also displaces hydrogen when boiled with
concentrated hydriodic acid, giving Agl.HI. Silver interacts with
cold nitric acid and with hot, concentrated sulphuric acid, giving the
nitrate or sulphate of silver and oxides of nitrogen or of sulphur (p. 446).
COPPER, SILVER, GOLD 629
Since its hydroxide has ho tendency to behave as an acid, alkalies,
whether in solution or fused, have no action upon silver. Hence alka-
line substances are heated in vessels of this metal or of iron, rather
than in vessels of platinum (q.v.~), because platinum is attacked by
alkaline materials.
The Halides of Silver. — The chloride, bromide, and iodide are
formed as curdy precipitates when a salt of silver is added to a solu-
tion containing the appropriate halide ion. The first is white, and
melts at about 457°. The second and third are very pale-yellow and
yellow respectively. The insolubility in water, which is very great,
increases in the above order. The iodide, after melting, solidifies and
forms quadratic crystals, which, as they cool, pass at 146° into a
different physical variety (hexagonal) with evolution of heat (cf.
pp. 368, 565).-
When exposed to light, the chloride becomes first violet and finally
brown, chlorine being liberated. The bromide and iodide behave simi-
larly. It is believed that a sub-chloride and sub-bromide Ag2Cl and
Ag2Br are formed in the early stages of the action (see Photography,
below). Solid silver chloride absorbs ammonia, forming first 2AgCl,
3NH3, and then AgCl,3NH3, the former with a tension of 93 mm., and
the latter with a tension of about one atmosphere of ammonia at 20°
(cf. p. 123). The bromide forms no compound in this way, but the
iodide yields 2AgI,NH3.
In consequence of the progressive insolubility, a cold solution of a
bromide will slowly convert the precipitate of silver chloride into
bromide, and a soluble iodide will similarly transform the bromide or
the chloride into iodide (cf. p. 585). Chlorine gas, however, displaces
bromine and iodine from the dry compounds (cf. p. 361). This illus-
trates well the absence of any relation between the electromotive
series and double decomposition (p. 362).
Silver fluoride may be made by treating the oxide or carbonate
with hydrofluoric acid (H2F2 + Ag20 -» 2AgF -f H20). The salt is
very soluble and deliquescent.
Complex Compounds of Silver. — Silver chloride dissolves easily
in excess of ammonium hydroxide, giving the complex cation Ag(N"H3)2'.
Under certain conditions octahedral crystals (Fig. 48, p. 138) of AgCl are
deposited from the solution, and, under other conditions, crystals of the
composition 2AgCl, 3NH3. The bromide, which is less readily soluble,
630 INORGANIC CHEMISTKY
gives the same complex ion. The iodide is hardly soluble at all. Am-
monio-argention Ag(NH3)2", in solutions of concentrations such as are
commonly used (.1 N to N), gives about the same concentration of argen-
tion Ag* as does the bromide, and much more than the highly insoluble
iodide (cf. p. 544). Hence the latter is almost insoluble in ammonium
hydroxide, and can be precipitated in ammoniacal solution. All three
of the insoluble halides dissolve in solutions of potassium cyanide and
of sodium thiosulphate, as do also all the other insoluble silver salts.
Usually an equivalent amount of the cyanide or thiosulphate suffices,
but for solution of the sulphide an excess is required. With the
cyanide, double decomposition gives first the insoluble silver cyanide
(AgCN) which then dissolves, forming the soluble potassium argenti-
cyanide K.Ag(CN)2. The thiosulphate gives a solution from which
crystals of a complex salt 2NaAgS2O3, Na2S203 are obtained. The com-
plex anion in the solution appears to be Ag(S203)2"'. Since the iodide
dissolves in the thiosulphate with considerable difficulty, we should
infer that the complex thiosulphate anion gives about the same concen-
tration of argention as does the iodide. An independent method of
measuring the concentrations of argention in all the solutions, places the
compounds in the order of diminishing ability to give argention thus,
AgCl, Ag(NH,V, AgBr, Ag(S203)2'", Agl, Ag(CN)/, Ag2S, and con-
firms the above inferences (see Concentration cells). The more active
metals, like zinc and copper, displace silver from all solutions, whether
the solutions contain simple or complex salts.
Oxides of Silver. — When sodium or potassium hydroxide is
added to a solution of a salt of silver, a pale-brown precipitate is
obtained, which, after being freed from water, is found to be Ag20. We
should expect to obtain the hydroxide (AgOH) in this fashion, but it
appears to be unstable. The aqueous solution of argentic oxide, how-
ever, is distinctly alkaline, and presumably therefore does contain the
hydroxide : 2AgOH <=± Ag20 -f- H2O. Silver oxide is formed by boil-
ing silver chloride with caustic potash. Since the oxide is much more
soluble than the chloride (p. 544), we should expect the reverse of the
above action to be the normal one. Here, however, the excess of potas-
sium hydroxide (hydroxidion) represses the ionization of the silver hy-
droxide and reverses the relations in regard to solubility (cf. p. 585).
Argentic oxide melts and gives off its oxygen at 250-270°. It is an
active basic oxide, and all the salts of silver are derived from it, the two
other oxides having no corresponding salts. When moist, it absorbs
COPPER, SILVER, GOLD 631
carbon dioxide from the air. Its solutions are said to show concentra-
tions of hydroxidion much smaller, it is true, than eqxiimolar solutions
of the active bases, but considerably greater than similar solutions of
ammonixim hydroxide (p. 331). The oxide dissolves easily in ammo-
nium hydroxide, and the ammonio-argentic hydroxide Ag(NH3)2.OH
which is formed is as active a base as is potassium hydroxide. The
solution, when allowed to evaporate, deposits black crystals of an
explosive substance whose composition has not been determined.
This is " fulminating silver " (not to be confused with fulminate of
silver Ag.ONC).
Silver peroxide Ag202 (cf. p. 308) is formed by the action of ozone
on silver. In the electrolysis of silver nitrate it is deposited in
shining black crystals on the positive electrode. There is also a
suboxide Ag40.
Silver Nitrate. — This salt is obtained by treating silver with
aqueous nitric acid :
3Ag + 4HNO3 -> 3AgNO3 + NO + 2H20.
From the solution, colorless rhombic crystals (Fig. 7, p. 13) isomorphous
with those of potassium nitrate (Fig. 98, p. 557) are deposited. These
melt at 218°. In. the form of thin sticks made by casting (lunar* caus-
tic), the substance is used in medicine, partly because it combines with
albumins to form insoluble compounds. When commercial silver, con-
taining copper, is used to make silver nitrate, the solution is evapo-
rated to dryness and heated at 250° until the nitrate of copper has all
been decomposed. At this temperature the silver salt is unaffected,
and when cool can be separated from the insoluble cupric oxide by ex-
traction with water.
The aqueous solution is neutral. The pure salt is not affected by
light, but when deposited on cloth, on the skin of the fingers, or on the
mouth of the reagent bottle, it is converted into the chloride, and from
this, in turn, silver is liberated. For this reason it is an ingredient in
marking-inks. The dry compound combines with ammonia, giving
AgN03. 3NH3. In its aqueous solution ammonium hydroxide pro-
duces, first a faint precipitation of the oxide, and then the soluble
complex salt Ag(NH8)2.N03.
Other Salts of Silver. — Silver carbonate, the neutral salt Ag2C03,
and not a basic carbonate is precipitated from solutions of salts of
* (Lat.) luna (the moon), the alchemical name for silver.
632 INORGANIC CHEMISTRY
silver by soluble carbonates. It is slightly yellow in color. With water
it gives a faint alkaline reaction, and, like calcium carbonate, is soluble
in excess of carbonic acid (p. 594). When heated, the carbonate clecom-
poses, leaving metallic silver. Other compounds of silver, for example,
the chloride, when heated in a crucible with sodium carbonate give this
salt by double decomposition, and hence are finally reduced to a button
of metallic silver. The sulphate is made by the action of concen-
trated sulphuric acid on the metal. It is not very soluble in water, and
crystallizes in rhombic prisms isomorphous with anhydrous sodium
sulphate. When it is mixed with a solution of aluminium sulphate
(q-v.~), octahedral crystals of silver-alum Ag2SO4, Al^SO^g, 24H20 are
obtained. Silver sulphide is precipitated by hydrogen sulphide from
solutions of all silver compounds, whether free acids are present or
not, and irrespective of the form in which the silver is combined.
Excess of potassium cyanide, however, prevents its precipitation from
the argenticyanide. The sulphide is formed by the action of metallic
silver on alkaline hydrosulphides, and this interaction forms the basis
of the " hepar " test for sulphur (p. 391). Silver orthophosphate Ag3PO4
(yellow), arsenate Ag3As04 (brown), and chromate Ag2Cr04 (crimson),
are produced by precipitation, and their distinctive colors enable us to
use silver nitrate in analysis as a reagent for identifying the acid
radicals.
Electroplating. — The process is similar to the electro-deposition
of copper (p. 617). The article to be plated is cleaned with extreme
care and attached to the negative wire. A plate of silver forms the
positive electrode, and since simple salts of silver do not give coherent
deposits, the bath is a solution of potassium argenticyanide. The
kalion (K*) migrates to the negative wire, and since potassium requires
a much greater E.M.F. for its liberation than does silver, silver is
there deposited from the trace of argention given by the complex silver
ions in the neighborhood:
Ag(CN)a' *=5 Ag- + 2CN'
Ag* + 0 -> Ag.
The potassium cyanide remains in solution. At the positive electrode
silver goes into solution in equivalent amount giving argention, and
the above equations are reversed.
Mirrors are silvered through the reduction of silver nitrate by
organic compounds such as potassium-sodium tartrate (Rochelle salt),
COPPER, SILVER, GOLD 633
glycerine, formaldehyde (formol), or sugar. On a small scale, dilute
silver nitrate is mixed with ammonium hydroxide until the solution is
clear, and then a little caustic potash, a few more drops of ammonia,
and finally a very little glycerine, are added. A watch-glass floated
on this mixture quickly acquires a deposit of silver.
Photography. — Bromo-gelatine dry plates are made by preparing
an emulsion of gelatine to which silver nitrate and a slight excess of
ammonium bromide have been added. After the emulsion has been
kept warm until the precipitate of silver bromide has coagulated into
small granules ("ripening"), it is allowed to solidify. It is then cut
up, and the ammonium nitrate is washed out with water. After drying
and remelting, the emulsion is finally applied to plates of glass. The
excess of ammonium bromide and the ripening both increase the
subsequent sensitiveness of the plates.
After exposure, often for only a fraction of a second, there is no
visible alteration in the film. The image is developed. Chemically,
this consists in reducing the silver bromide to metallic silver by means
of reducing agents. While the whole of the halide upon the plate is
reducible, if the reducing agent is kept upon it for a sufficient length
of time, the parts reached by the light are affected first, and with a
speed proportional to the intensity of the illumination undergone by
each part. The reducing agent is poured off when sufficient " contrast "
between the parts variously illuminated has been attained. The un-
reduced silver bromide is then dissolved out with sodium thiosulphate
("hyposulphite of soda" or "hypo"), and the silver image is thus
saved from obliteration by the silver that would be deposited if
the plate were to be brought into the light without this treatment
(fixing). The result is a "negative," as the parts brightest in the
object are now opaque, and the darkest parts of the object are trans-
parent.
According to one view, the exposure reduces certain portions of the
bromide to a sub-bromide, perhaps Ag2Br, which is more easily reduced
than silver bromide, and is consequently first attacked by the developer.
The first particles of free silver then interact with neighboring mole-
cules of AgBr, giving more Ag2Br. Thus the reduction proceeds exten-
sively wherever Ag2Br is found and in proportion to its amount. The
gelatine is the sensitizing substance, and promotes the dissociation of
the silver bromide (2AgBr <=» Ag2Br -f- Br), which is a reversible action,
by combining with the bromine. Potassium bromide, when added to
634 INORGANIC CHEMISTRY
the developer, restrains the development, probably by rendering the
silver bromide less soluble (cf. p. 583).
The simplest developer is potassium-ferrous oxalate K2.Fe(C204)2,
a solution of which may be made by mixing ferrous sulphate and
potassium oxalate. For the sake of simplicity we may regard the
action as a reduction by means of ferrous oxalate, which itself is
oxidized to ferric oxalate (Fe2(C204)3) :
3FeC204 + 3Ag2Br -> Fe2(C204)3 + FeBr3 + 6Ag.
In brief, we have 3Fe" becoming 2Fem + Fem, and this amount of
bivalent iron therefore takes up 3Br, liberating the silver with which
it was combined. Other developers commonly employed are alkaline
solutions of the sodium salts of hydroquinone and pyrogallic acid.
In printing, the light and dark are again reversed, the denser parts
of the negative protecting the compounds on the paper below it from
action, and leaving them white. Either "bromide" papers, which require
only brief exposure and are developed like the plate, are used, or silver
chloride is the sensitive substance, and prolonged exposure to light is
allowed to liberate the proper amount of silver. The operation of
fixing is performed as before. In toning, a solution of sodium chlor-
aurate is employed. A portion of the silver dissolves, displacing gold
(p. 362), which is deposited in its place :
NaAuCl4 + 3Ag -> NaCl 4- 3AgCl + Au.
The thin film of gold gives a richer color to the print. In platinum
toning, potassium chloroplatinite K2PtCl4 is similarly used.
Many other actions are utilized in photography. Thus, ferric
oxalate is reduced by light to ferrous oxalate : Fe2(C204)3 —> 2FeC204
+ 2C02. When paper coated with a solution of the former, or a
mixture of ferric chloride and ammonium oxalate, is used for printing,
the pale-yellow ferric salt loses its color where it has been turned into
the ferrous salt. If the paper is then dipped in a solution of ferri-
cyanide of potassium K3.Fe(CN)6 the ferrous salt precipitates the
insoluble and deep-blue ferrous ferricyanide Fe3[Fe(CN)6]2, while the
unchanged ferric salt simply gives a soluble brown substance, which
can be washed out. For regular blue prints, ammonium-ferric citrate is
employed instead of the oxalate. If the above paper, after printing,
is dipped in. potassium chloroplatinite (or has been coated with this
salt at the same time that it received the ferric oxalate), and is then
COPPER, SILVER, GOLD 035
dipped in potassium oxalate solution, the latter dissolves the insoluble
ferrous oxalate, and the potassium-ferrous oxalate reduces the platinum
compound, giving a platinum print :
6FeC204 + 3K2PtCl4 -> 2Fe2(C2O4)8 + 2FeC]8 + 3Pt + 6KC1.
We have already seen (p. 484) that light of short wave-length —
blue and violet — has the greatest effect upon silver halides. The
time, in seconds, required for equal effects is approximately : violet
15, blue 29, green 37, yellow 330, red 600. Hence objects showing
to the eye a variety of colors are entirely misrepresented, as regards
the relative brightness of their parts, by photography. Now the im-
portant fact, in this connection is, that only that part of the light which
is absorbed in traversing the film, and not that which is scattered or
transmitted, can be used for chemical change. Hence, dipping plates
in solutions of substances capable of absorbing yellow and red radia-
tions causes them to absorb more of the energy of these photographi-
cally weakest radiations, and to give greater chemical action in response
to them. This partially restores the balance. Such plates are called
orthochromatic, and are made with substances like eOsin or cyaniue.
Analytical Reactions of Silver Compounds. — The ion of salts
of silver, argention Ag", is colorless. Many of its compounds are in-
soluble, the precipitation of the chloride, which is insoluble in dilute
acids, being used as a test. Mercurous chloride and lead chloride are
also white and insoluble, but silver chloride dissolves in ammonium
hydroxide, mercurous chloride ('/.#.) turns black, and lead chloride,
which is also soluble in hot water, is not altered in color. With excess
of ammonium hydroxide, silver salts give the complex cation
Ag(NH3)2* and, from solutions containing silver in this form, only the
iodide and sulphide can be precipitated. Sodium thiosulphate and
potassium cyanide dissolve all silver salts, giving salts of complex acids
with silver in the anion (p. 629). Zinc displaces silver from all forms
of combination.
GOLD.
Chemical Relations of the Element. — This element forms
two very incomplete series of compounds corresponding respectively
to aurous and auric oxides, Au20 and Au2Os. The former is a
feebly basic oxide, the latter mainly acid-forming. No simple salts
636 INORGANIC CHEMISTRY
with oxygen acids are stable. All the compounds of gold are easily
decomposed by heat with liberation of the metal. All other common
metals displace gold from solutions of its compounds (p. 362). Mild
reducing agents likewise liberate gold. The element enters into many
complex anions (p. 536).
Occurrence and Metallurgy. — Gold is found chiefly in the free
condition disseminated in veins of quartz, or mixed with alluvial sand.
Small quantities are found also in sulphide ores of iron and copper.
Telluride of gold (sylvanite), in which silver takes the place of part of
the gold [Au,Ag]Te2*, is found in Colorado. This mineral when
heated loses its tellurium, and gold, alloyed with silver, remains.
From the alluvial deposits, gold is usually separated by washing
in a cradle, as in the Klondyke. Quartz veins, which in the Transvaal
Colony reach a thickness of a meter and carry an average of 18 g. of
gold per ton, are mined, and the material is pulverized with stamping
machinery. About 55 per cent of the gold is then separated by allowing
the powdered rock to be carried by a stream of water over copper plates
amalgamated with mercury. The gold dissolves in the latter, and is
secured by removal and distillation of the amalgam. The finer particles
contained in the sludge which runs off ( " tailings " ), are extracted by add-
ing a dilute solution of potassium cyanide (Mac Arthur-Forest process)
and exposing the mixture to the air. Oxidation and simultaneous inter-
action with the cyanide give potassium aurocyanide. ' Hydrogen perox-
ide, which is formed in many oxidations by free oxygen, is produced
also:
2Au + 4KCN 4- 2H204-02 -> 2KAu(CN)2 4- 2KOH 4- H202,
2Au + 4KCN 4- H,02 -» 2KAu(CN)2 + 2KOH.
From this solution the gold is isolated, either by electrolysis, 'in
which a plate of lead forms the cathode (and is subsequently cupelled.
Siemens-Halske process), or in the form of a purple powder by pre-
cipitation with zinc.
Auriferous pyrites is washed, and then treated with chlorine gas.
The chloride of gold which is formed is dissolved out with water. From
* Amongst minerals, mixed crystals of isomorphous salts are so commonly found
that formulae like the above are constantly used by mineralogists. [Au,Ag]Te2
indicates a mixture in varying proportions of the isomorphous tellurides AuTe2 and
AgTea.
COPPER, SILVER, GOLD 637
the solution, the gold is precipitated with ferrous sulphate or oxalic
acid :
2AuCl3 + 6FeS04 -* 2Fe2(S04)3 + 2FeCl3 + 2Au,
2AuCl3 + 3H2C204^ 6HC1 + 6C02 + 2Au.
In the former case a purple powder, and in the latter, if the solution is
heated, a spongy mass (the form used by dentists), is obtained.
The gold separated from ores in the above ways contains silver,
copper, lead, and other metals, and various methods of refining, elec-
trolytic and otherwise, are used. In one of these the gold is melted,
and a stream of chlorine is passed through it. The metals, excepting
gold, are converted into chlorides. The chloride of silver rises as a
liquid to the surface, while chlorides of arsenic and antimony are vola-
tilized. A layer of melted borax prevents loss of silver chloride by
volatilization. The silver chloride, when it has solidified, is placed
between wrought-iron plates and reduced by the action of dilute sul-
phuric acid upon the latter.
The world's production of gold during the first half of the nine-
teenth century averaged 27 tons annually. In 1897 it was 363 tons,
and in 1899, 472.6 tons. In the former year North America, including
Canada, produced 28.5 per cent of the whole, the Transvaal Colony
23.2 per cent, and Australia 21.2 per cent.
Properties of the Metal- — Gold is yellow in color, and is the
most malleable and ductile of all the metals. It melts at 1064°. To
give it greater hardness it is alloyed with copper, the proportion of gold
being defined in " carats." Pure gold is " 24-carat." British sovereigns
are 22-carat and contain ^ of copper. American, French, and German
coins are 21.6-carat, or 90 per cent gold. Silver takes the place of
copper in Australian sovereigns.
Gold is not affected by free oxygen or by hydrogen sulphide. It
does not displace hydrogen from dilute acids, nor does it interact with
nitric or sulphuric acids or any oxygen acids except selenic acid. It
combines, however, with free chlorine, and it therefore interacts with a
mixture of nitric and hydrochloric acids (aqua reyia), which gives
off this gas (p. 448). Chlorauric acid H.AuCl4( = HCl,AuCl,) is
formed, and the action is assisted by the fact that the gold ions are
taken into the little-dissociated anion AuCl/. Gold is the least active
of the familiar metals.
638 INORGANIC CHEMISTRY
Compounds with the Halogens. — Chlorauric acid, formed as
above, is deposited in yellow, deliquescent crystals of H. AuCl4, 4H20.
The yellow sodium chloraurate NaAuCl4, 2H20, obtained by neutraliza-
tion of the acid, is used in photography (p. 634). The acid gives up
hydrogen chloride when heated very gently, leaving the red, crystalline
auric chloride AuClg. The tendency to form complex compounds is
such, however, that when dissolved in water free from hydrochloric
acid, this salt gives ILj.AuCLjO. Red crystals of H2AuCl30, 2H20 are
deposited by the solution. When auric chloride is heated to 180°
aurous chloride AuCl and chlorine are formed. This salt is a white
powder. It is insoluble in water, but in boiling water is converted
quickly into auric chloride and free gold : 3AuCl — > AuCl3 + 2Au.
When potassium iodide is added to a solution of chlorauric acid, or to
sodium chloraurate, the yellow aurous iodide is precipitated :
NaAuCl4 + SKI -4 NaCl + Aul + I2 + 3KC1.
The action is like that on cupric salts (p. 621), and for a similar reason,
namely, that auric iodide is not stable.
Other Compounds. — When caustic alkalies are added to chlor-
auric acid, or to sodium chloraurate, auric hydroxide Au(OH)8 is pre-
cipitated. This substance is an acid, and interacts with excess of the
base, forming aurates. These are derived from met-auric acid (Au(OH)3
— H20 = HAu02), as, for example, potassium aurate K.Au02, 3H20.
This salt interacts by double decomposition, giving, for instance, with
silver nitrate, the insoluble silver salt AgAu02. Its solution is alkaline
in reaction, a fact which shows that auric acid is a weak acid (cf. p.
344).
Auric oxide A^Og is a brown, and aurous oxide A^O is a violet
powder. With hydrochloric acid the latter gives chlorauric acid and
free gold.
On account of its reducing action, hydrogen sulphide precipitates
from chlorauric acid a dark-brown mixture containing much aurous
sulphide AugS and free sulphur, as well as some auric sulphide Au2S3.
The sulphides interact with alkali sulphides, giving complex sulphaurites
and sulphaurates, such as K8AuS2(= 3K2S, Aii2S) and KAuS2(= K2S,
A^Sg), which are soluble (p. 537, and see Tin, Arsenic, and Antimony).
The aurocyanides, like K.Au(CN)2(= KCN, AuCN), and the auri-
cyanides like K.Au(ClSr)4( = KCN, Au(dSr)8), are formed by the action
COPPER, SILVER, GOLD 639
af potassium cyanide on aurous and auric compounds respectively. They
are colorless and soluble. Their solutions are used as baths, in con-
junction with a gold anode, for electrogilding.
Analytical Reactions of Gold. — The metallic "streak," produced
by rubbing the metal on touchstone (Lydian stone, a black basalt),
is not easily removed by nitric acid of sp. gr. 1.36 (57.5 per cent).
In assaying, the material containing the gold is heated with borax and
lead in a small crucible (cupel) of bone-ash. The lead and copper are
oxidized, and the oxides are absorbed by the cupel, leaving a drop of
molten alloy of gold and silver. The cold button is flattened by ham-
mering and rolling, and treated with nitric acid to remove the silver.
The gold, which remains unattacked, is washed, fused again, and
weighed. The acid will not interact with the silver and remove it
completely if the quantity of gold exceeds 25 per cent. When the
proportion of gold is greater than this, a suitable amount of pure silver
is fused with the alloy (." quartation").
Exercises. — 1. How much copper will be deposited per hour on
each sq. cm. of an electrode immersed in cupric sulphate solution when
the current density is \ ampere per sq. cm. (p. 323) ? How much
copper would be obtained under the same conditions from a cuprous
salt?
2. Write equations for the interactions (a) of salt water and oxygen
with copper (p. 618), (£>) of ferrous oxide and sand (p. 617), (e)of ver-
digris, arsenious acid, and acetic acid (p. 624).
3. Write the formulae of the basic chloride, nitrate, carbonates, and
sulphate of copper as if these substances were composed of the normal
salt, the oxide and water (p. 618).
4. What may be the formula of the compound of cupric hydroxide
and sodium tartrate (p. 623) ?
5. Can you develop any relation between the facts that solutions
of cupric salts are acid in reaction and that they give basic carbonates
by precipitation ?
6. Formulate the action of potassium cyanide in dissolving cupric
hydroxide and cuprous sulphide, assuming that potassium cupro-
cyanide is formed.
7. How should you set about making cupric orthophosphate (in
solution), ammonium cuprocyanide, and lead cuprocyanide ?
640 INORGANIC CHEMISTRY
8. Write the formulae of some of the double salts analogous to
potassium-cupric sulphate (p. 625).
9. What chemical agents are present in a Bunsen flame ? If borax
beads were made in the oxidizing flame with cupric chloride, cuprous
bromide, and cupric sulphate, severally, what actions would take
place ?
10. If the solubility ratio of silver in lead and in zinc were 1000 :
1, and 2 per cent of zinc were used, what proportion of the total silver
would be secured by Parke's method?
11. Which is more stable, silver sulphate or cupric sulphate, silver
nitrate or cupric nitrate? To what salts are the silver compounds in
this respect more closely allied ?
12. Write the equations for the interaction of (a) silver and con-
centrated sulphuric acid, (&) silver chloride and sodium carbonate
when heated strongly, (c) sodium thiosulphate and silver bromide, (d)
potassium ferricyanide and ferrous oxalate.
13. What reagents should you use to precipitate the phosphate,
arsenate, and chromate of silver ?
14. Write the equations for the interactions of (a) gold and selenic
acid, in which selenious acid is formed, (b) potassium hydroxide and
auric hydroxide, (c) potassium cyanide and sodium chloraurate.
15. In what respects are the elements of this family distinctly
metallic, and in what respects are they allied to the non-metals (p. 533) ?
16. Collect all the evidence tending to show that the cuprous com-
pounds are more stable than the cupric.
17. Describe in terms of the categories used by the phase rule the
systems (a) cupric nitrate and water at 24.5° and (&) silver iodide at
146°.
18. Make a classified list of the methods by which cupric com-
pounds are transformed into cuprous, and vice versa.
CHAPTER XXXVII
GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY.
THE RECOGNITION OF CATIONS IN QUALITATIVE
ANALYSIS
The Chemical Relations of the Family* — The remaining ele-
ments of the third column of the periodic table, namely, glucinum or
beryllium (Gl, or Be, at. wt. 9.1), magnesium (Mg, at. wt. 24.36), zinc (Zn,
at. wt. 65.4), cadmium (Cd, at. wt. 112.4), and mercury (Hg, at. wt. 200.0),
although all bivalent, do not form a coherent family. Glucinum and
magnesium resemble zinc and cadmium, and differ from the calcium
family, in that the sulphates are soluble, the hydroxides easily lose
water leaving the oxides, the chlorides are comparatively volatile, and
the metals are not rapidly rusted in the air and do not easily displace
hydrogen from water. They resemble the calcium family, and differ
from zinc and cadmium, in that the sulphides are hydrolyzed by water,
the oxides are not reduced by heating with carbon, complex cations are
not formed with ammonia, and the metals do not enter into complex
anions. But glucinum differs from magnesium and resembles zinc in
that its hydroxide is acidic as well as basic. This is not unnatural,
since in the periodic system it lies between lithium, a metal, and
boron, a non-metal. Mercury is the only member of the group that forms
two series of compounds. These are derived (p. 278) from the oxides
HgO and Hg2O. Mercury approaches the noble metals in the ease
with which its oxide is decomposed by heating, and in the position of
the free element in the electromotive series.
The vapor densities of zinc, cadmium, and mercury show the
vapors of these three metals to be monatomic.
The compounds of the metals of this family give no color to the
borax bead.
GLUCINUM.
Chemical Relations of the Element. — Glucinum (or beryllium)
is bivalent in all its compounds. Its oxide and hydroxide are basic,
and are also feebly acidic towards active bases (see Zinc hydroxide).
641
642 INORGANIC CHEMISTRY
On account of this fact and the extreme ease with which its carbonate
gives up carbon dioxide, in both of which respects it resembles alu-
minium, it was first thought to be trivalent. This made its atomic
weight 13.6, the amount combining with one chemical unit of chlorine
being 4.55. In the periodic system, however, there was a space for a
bivalent element with the atomic weight 9.1 (= 2 x 4.55) between
lithium and boron, and none for a trivalent element. Later (1884)
Nilson and Pettersson determined the vapor density of the chloride and
of certain organic compounds of the element, and found only 9.1 parts
of glucimim in the molar weights of the compounds. The element
derives its name from the sweet taste of its salts (Gk. yAwv's, sweet).
The Metal and its Compounds. — Glucinum occurs in beryl, a
metasilicate of glucinum and aluminium Al2Gl2(Si03)6. Specimens of
beryl tinted green by the presence of a little silicate of chromium are
known as emeralds. The metal may be obtained by electrolysis of
the easily fusible double fluoride G1F2, 2KF. In powdered form it
burns when heated in the air. It displaces hydrogen from cold, dilute
acids, and also, when heated, from caustic potash :
Gl + 2KOH -+ K2G102 + H2.
The oxide interacts with acids and with strong bases. The salts give
no color to the Buiisen flame.
MAGNESIUM.
Chemical Relations of the Element. — Magnesium is bivalent
in all its compounds. The oxide and hydroxide are basic exclusively.
The element does not enter into complex cations or anions.
Occurrence. — • Magnesium carbonate occurs alone as magnesite,
and in a double salt with calcium carbonate MgCO3, CaCO3 as dolomite.
The sulphate and chloride are found as hydrates and as constituents
of double salts (see below) in the Stassfurt deposits. Silicates are
also common. Olivine is the orthosilicate Mg2Si04. Serpentine is a
hydrated disilicate, [Mg,Fe]3,Si207, 2H20, as is also meerschaum, and
asbestos is an anhydrous silicate. The element derives its name from
Magnesia, a town in Asia Minor.
The Metal. — Magnesium is manufactured by electrolysis of de-
hydrated and fused carnallite MgCl2,KCl, 6H2O. The iron crucible
GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY 643
in which the material is melted forms the cathode, and a rod of carbon
the anode. The metal is silver- white, and when heated can be pressed
into wire and rolled into ribbon. Commercial specimens of the latter
often contain zinc.
Chemically the metal is less active than are the metals of the alka-
line earths. It slowly becomes coated with a layer of the oxide. It
displaces hydrogen from boiling water and, of course, from cold, dilute
acids. Magnesium burns in air with a white light, rich in rays of
short wave-length such as act upon photographic plates (p. 458). The
ash contains the nitride Mg3N2, as well as the oxide. The presence
of the former may be shown by the evolution of ammonia when the
white powder is boiled with water (p. 417). When the metal is heated
with the oxides of boron, of silicon, and of many of the metals, it com-
bines with the oxygen and liberates the other element.
Powdered magnesium is used in pyrotechny, and, with potassium
chlorate (10 : 17), in making flash-light powder for use. in photography.
Magnesium Chloride. — This salt occurs in salt deposits as the
hexahydrate MgCl^GHjO, a highly deliquescent compound obtained also
by evaporating an aqueous solution, and as carnallite MgCl^ KC1, 6H20.
The latter is an important source of potassium chloride (p. 551), and
almost all the magnesium chloride combined with it is thrown away.
When the hexahydrate is heated, a part of the chloride is hydrolyzed,
some magnesium oxide remaining, and some hydrogen chloride being
given off. Sea-water cannot be used in ships' boilers because of the
hydrochloric acid liberated by the magnesium chloride which the water
contains. The salt forms a double chloride with ammonium chloride
MgClj, NH4C1, 6H20 which is isomorphous with carnallite, and this salt
can be dehydrated without hydrolysis of the chloride. Afterwards
the ammonium chloride can be volatilized (p. 421). To utilize natural
magnesium chloride, the manufacture of chlorine from it, by passing
air and steam over the salt at a high temperature, has been attempted :
4MgCl2 + 2H20 + O2 -> 4MgO + 4HC1 + 2CL,.
The Oxide and Hydroxide. — Magnesium oxide is made by
heating the carbonate, and is known as " calcined magnesia." It is a
white, highly infusible powder, and is used for lining electric furnaces
and making crucibles. It combines slowly with water to form the
hydroxide.
644 INORGANIC CHEMISTRY
The hydroxide is found in nature as brucite. It is also precipi-
tated from solutions of magnesium salts by alkalies. It is very
slightly soluble in water, much less so than calcium hydroxide, but
more so than are the hydroxides of zinc and the other heavy metals.
The solution has a barely perceptible alkaline reaction.
Magnesium hydroxide is not precipitated by ammonium hydroxide
when aminouium salts are present also. The ammonium salts, being
highly ionized and giving a high concentration of ammonion NH4*, re-
press the ionization of the feebly ionized ammonium hydroxide, and so
reduce the concentration of hydroxidioii which it furnishes. With the
ordinary concentration of Mg", therefore, the amount of hydroxidion
existing in presence of excess of a salt of ammonium is too small to
bring the solubility product [Mg"] x [OH']2 up to the value required
for precipitation. Conversely, magnesium hydroxide interacts with
solutions of ammonium salts and passes into solution :
Mg(OH)2 (solid) <r» Mg(OH)2 (diss'd) ^ Mg"+ 20H' ) __
2NH4C1 <± 2C1' + 2NH4- } *
In presence of excess of ammonium, chloride, the OH' combines with
NH*4to form molecular ammonium hydroxide, and the equilibria in the
upper line are displaced forwards to generate a further supply of the
former. With sufficiently great concentration of the ammonium chloride,
all the magnesium hydroxide may thus dissolve ; with only a small excess
a condition of equilibrium with solid magnesium hydroxide is reached.
The whole c*ase is analogous to the interaction of acids with insoluble
salts (p. 598). Magnesium oxide also dissolves in salts of ammonium.
It gives first the hydroxide by interaction with the water.
Magnesium Carbonate. — The normal carbonate is found in nature.
Only hydra.ted basic carbonates are formed by precipitation, and their
composition varies with the conditions. The carbonate manufactured in
large amounts and sold as magnesia alba is approximately Mg4(OH)2
(C03)3. 3H20. The carbonates are not precipitated in the presence of
ammonium salts, and interact with such salts in the same way as does
the hydroxide.
*
Magnesium Sulphate. — The common heptahydrate MgS04, 7H2O
crystallizes from cold water in rhombic prisms, and is called Epsom salts.
At 0° a dodecahydrate appears. The heptahydrate is efflorescent, and
loses its water by stages and with decreasing aqueous tension, The inono-
ULUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY 645
hydrate, found in the salt layers as kieserite MgS04,H20, has a very
low aqueous tension, and is not rapidly dehydrated except above 200°.
The hepta- and inonohydrates present a striking case of difference in
solubility in two forms of one salt, the former giving at 15° a solution
containing 33.8 g. of the sulphate in 100 g. of water, while the latter
is almost insoluble. Magnesium sulphate is used in the manufacture
of sodium and potassium sulphates, and is employed also for " loading "
cotton goods, and as a purgative.
Magnesium Sulphide. — The sulphide may be formed by heating
the metal with sulphur. It is insoluble in water, but is decomposed
and gives, finally, hydrogen sulphide and magnesium hydroxide :
2MgS -H 2H20 s^Mg(SH)2 + Mg(OH)2,
Mg(SH)2 + 2H20 ^ Mg(OH)2 1 + 2H2S.
The hydrolysis is more complete than in the case of calcium sulphide,
and eliminates all the hydrogen sulphide, because magnesium hydrox-
ide is much more insoluble than calcium hydroxide, and so there is
little reverse interaction tending to reproduce the soluble hydro-
sulphide Mg(SH)2:
Phosphates of Magnesium. — The only phosphate of importance
is ammonium-magnesium orthophosphate NH4MgP04, 6H20, which
appears as a crystalline precipitate when sodium phosphate and
ammonium hydroxide are mixed with a solution of a magnesium salt.
This compound is insoluble in water containing ammonium hydroxide,
and is used, in quantitative analysis for estimating both magnesium,
and phosphoric acid. Before being weighed the precipitate is ignited,
and is thus converted into the anhydrous pyrophosphate of magnesium
Mg0P,07. The salt NH4MgAs04,6H20 has similar properties, and is
used for estimating arsenic acid.
Analytical Reactions of Magnesium Compounds. — The mag-
nesium ion is colorless and bivalent. It does not enter into complex
ions. Soluble carbonates precipitate basic carbonates of magnesium,
but not when ammonium salts are present. The latter limitation dis-
tingiiishes compounds of magnesium from those of the calcium family.
Potassium hydroxide precipitates the hydroxide of magnesium, except
when salts of ammonium are present. The mixed -phosphate of ammo-
nium and magnesium, in presence of ammonium hydroxide, is the
least soluble salt,
646 INOKGANIC CHEMISTRY
ZINC.
Chemical Relations of the Element. — Zinc is bivalent in all its
compounds. Of these there are two sets, — the more numerous and
important one in which zinc is the positive radical (Zn.S04, Zn.Cl2,
etc.), and a less numerous set, the zincates, in which zinc is in the
negative radical (Na2.Zn02, etc.). Both sets of salts are hydrolyzed by
water, as the hydroxide is feeble whether it is considered as an acid or
as a base. The element also enters into complex cations and anions.
The salts are all poisonous.
Occurrence and Extraction from the Ores. — The chief sources
of zinc are calamine or smithsonite ZnC03, zinc-blende (Ger. blenden,
to dazzle) or sphalerite ZnS, franklinite Zn(FeO2)2, and zincite ZnO.
The red color of the last is due to the presence of manganese.
The ores are first converted into oxide — the carbonate by ignition,
and the sulphide by roasting. The sulphur dioxide is used to make
sulphuric acid. A mixture of the oxide with coal is then distilled in
earthenware retorts at 1300-1400°, the zinc condensing in earthenware
receivers, while carbon monoxide burns at a small opening :
2ZnS + 302 -* 2ZnO + 2S02,
ZnO + C -» CO + Zn.
At first zinc dust, a mixture of zinc and zinc oxide, collects in the
receiver, and afterwards liquid zinc. The product, which is cast in
blocks, is called spelter. It contains small amounts of lead, arsenic,
iron, and cadmium, because the sulphides of these metals are almost
invariably present in zinc-blende.
Properties and Uses of the Metal. — Zinc is a bluish-white
crystalline metal. When cold it is brittle, but at 120-150° it can be
rolled into sheets between heated rollers and then retains its pliability
when cold. At 200-300° the metal becomes once more brittle, at 433°
it melts, and at 920° it boils. The vapor density at 1740° is 2.64, and
the molecular weight, therefore, 2.64 x 28.955 (p. 215) or 76.4.
The gas is thus monatomic.
The metal burns in air with a bluish flame, giving zinc oxide.
When cold it is not affected by dry air, but in moist air it is oxidized,
and becomes covered with a firmly adhering layer of basic carbonate
which protects it from further action. The metal displaces hydrogen
from dilute acids, but with pure specimens the action almost ceases in
GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY 647
consequence of the formation of a layer of condensed hydrogen on the
surface. Contact with a less electro-positive metal, such as lead, iron,
copper, or platinum, enables the action to go on, because the hydrogen
is then liberated at the surface of the other metal (see Electromotive
chemistry). Crude zinc contains lead and iron and is therefore more
active than pure zinc. Zinc also attacks boiling alkalies, giving the
soluble zincate (see below) : 2KOH -f Zn — »K2Zn02 + H2. The action
on ammonium hydroxide is slower and different in nature :
Zn + 2NH4OH + 2NH8 -> Zn(NH8)4.(OH)a-f H2,
a complex cation being formed.
Sheet zinc, in consequence of its lightness (sp. gr. 7), is used in
preference to lead (sp. gr. 11.5) for roofs, gutters, and architectural
ornaments. Galvanized iron is made by dipping cleaned sheet iron in
molten zinc. The latter, being more active (p. 362), is rusted instead
of the iron. Zinc is used also in batteries and for making alloys
(p. 619). It mixes in all proportions with tin, copper, and antimony,
but with lead (p. 627) and with bismuth separation into two layers
occurs, each metal dissolving only a little of the other. The two
different modes of behavior resemble those of alcohol and water
(p. 136) and ether and water (p. 147) respectively.
Zinc Chloride. — This salt is usually manufactured by treating
zinc with excess of hydrochloric acid, evaporating the solution to dry-
ness, and fusing the residue. When hydrochloric acid is thus present,
the chloride ZnC^ is obtained. Evaporation of the pure aqueous solu-
tion, which is acid in reaction, results in considerable hydrolysis and
formation of much of the basic chloride
H20 <r± HC1 + Zn(OH)Cl, (1)
2Zn(OH)Cl -> ZiLpCL, + H2O. (2)
The salt is used in solid form as a caustic and, by injection of a solu-
tion into wood (e.g., railway sleepers), as a poison to prevent the
growth of organisms which promote decay. In both cases the salt com-
bines with albumins, forming solid products. The aqueous solution,
being acid, is employed also for dissolving the oxides from surfaces
which are to be soldered. The acid is reproduced by hydrolysis as fast
as it is used, and finally the oxychloride remains (equation 1 above).
648 INORGANIC CHEMISTRY
Zinc Oxide and Hydroxide and the Zincates. — The oxide is ob-
tained as a white powder by burning zinc or by heating the precipitated
basic carbonates. It turns yellow when heated, recovering its white-
ness when cold, in the same way that mercuric oxide is brown whilst
hot and bright red when cold. It is employed in making a paint —
zinc-white or Chinese white — which is not darkened by hydrogen
sulphide. For filling teeth, dentists sometimes use a paste made by
mixing the oxide with a strong solution of zinc chloride. It quickly
sets to a hard mass of oxychloride.
The lurdroxide of zinc appears as a white, flocculent solid when
alkalies are added to solutions of zinc salts. It interacts as a basic
hydroxide with acids, giving salts of zinc :
Zn(OH)2 + H2S04 <=± Zn.S04 + 2H2O.
It also interacts with excess of the alkali employed to precipitate it,
giving a soluble zincate :
H2ZnO2 1 + 2KOH +± K2.ZnO2 + H2O.
Both actions are reversible, and the second requires a considerable
excess of alkali for its completion : in fact, most of the zinc hydroxide
seems to be simply in colloidal solution. From a consideration of these
facts it is evident that zinc hydroxide when in solution is ionized both
as an acid and as a base :
2H' + Zn02"<=± Zn(OH)2 (diss'd) <=» Zn» + 20 H'
IT
Zn(OH)2 (solid)
The ionization as an acid is less than that as a base, but both are
small. Addition of an acid like sulphuric acid, however, furnishes
hydrion ; the hydroxyl ions combine with this to form water, and all
the equilibria are displaced to the right. With a base, oo. the other
hand, the hydrion is removed and the basic ionization simultaneously
repressed, so that the equilibria are displaced to the left.
Zinc hydroxide interacts with ammonium hydroxide, giving a soluble
complex compound with ammonia Zn(]S~H3)4.(OH)2. The case is like
those of copper (p. 623) and silver hydroxides (p. 631), and not like
that of magnesium hydroxide (p. 644).
Compounds of zinc, when heated in the Bunsen flame with a salt
pf cobalt, gives a zincate of cobalt (Binmann's green) CoZuO,.
GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY 649
Hydrogen sulphide precipitates zinc sulphide from solutions of
zincates and from solutions containing ammonia, so that some zinc
ions Zn** are present in both.
Carbonate of Zinc. — The normal zinc carbonate may be precipi-
tated by means of sodium bicarbonate :
ZnS04 + 2NaHC03 -* Na^SO, + ZnCO8 + H2O + C02.
The normal carbonate of sodium, however, gives basic carbonates,
which, as in the case of magnesium (p. 644), vary in composition
according to the conditions.
Zinc Sulphate. — This salt is formed when zinc-blende is roasted.
It gives rhombic crystals of the hydrate ZnSO4, 7H20. This, and the
corresponding compounds of magnesium MgS04, 7H2O, of iron FeS04,
7H20, and of several other bivalent metals, are all isomorphous, and are
known as vitriols. The zinc salt is white vitriol. Like Epsom salts,
it is dehydrated by stages, the last molecule of water being difficult to
remove. It is used in cotton-printing and as an eye-wash (\ per cent
solution).
The salt gives double salts with potassium and ammonium sulphate, of
the form ZnS04,K2S04, 6H20, which crystallize in the monosymmetric
system, and are isomorphous with each other, and with double salts
containing copper (p. 625), mercury (Hg11), iron (Fe11), magnesium, arid
other bivalent elements in place of the zinc. These compounds, unlike
the complex cyanides, are almost completely decomposed in dilute
solution (cf. p. 537).
Zinc Sulphide. — This compound is the only familiar sulphide
which is white. The yellow color of zinc-blende is caused by the
presence of sulphide of iron. Zinc sulphide- is more soluble in water
than is sulphide of copper, and hence it interacts with excess of strong
acids, and passes into solution. It is not soluble enough, however, to
be much affected by weak acids like acetic acid. This sort of behavior
is shown also by calcium oxalate (p. 598), and was discussed fully in that
connection. Zinc sulphide is thus capable of being precipitated when
acetic acid is present, or when hydrogen sulphide is led into a solution
of the acetate of zinc :
Zn(C3H8Ojl)a +
650 INORGANIC CHEMISTRY
But when an active acid is present, or is formed during the operation,
the sulphide is precipitated incompletely or not at all, the action being
highly reversible :
ZnS04 + H2S <=> ZnS + H2SO4.
There are thus two ways of obtaining the sulphide by precipita-
tion. A soluble sulphide causes it to be thrown down completely
because no acid is liberated in the action :
ZnCL, + (NH4)2S <=> ZnS J + 2NH4C1.
The other method is to add sodium acetate to the solution of the salt,
and then lead in hydrogen sulphide. The acid which is liberated by the
action upon the salt interacts with the sodium acetate, giving a neutral
salt of sodium and acetic acid, and the zinc sulphide is not affected by
the latter. In terms of the ionic hypothesis, the hydrion, liberated as
the hydrogen sulphide interacts with the zinc salt, combines with
acetanion introduced by the sodium acetate, and gives the little-ion-
ized acetic acid.
Analytical Reactions of Zinc Salts. — Zinc sulphide is precipi-
tated by the addition of ammonium sulphide to solutions of zinc salts
and of zincates. Sodium hydroxide gives the insoluble hydroxide,
which, however, interacts with excess of the alkali, giving the soluble
zincate of sodium. Compounds of zinc, when heated on charcoal with
cobalt nitrate, give Binmann's green (p. 648).
CADMIUM.
Chemical Relations of the Element. — This element is biva-
lent in all its compounds. Its oxide and hydroxide are basic exclu-
sively, and the salts are not hydrolyzed by water. It enters into com-
plex compounds having the ions Cd(NH3)4", Cd(CN")4", and Cdl/'.
The Metal. — Aside from the rare mineral greenockite CdS, cad-
mium is found only in small amounts, as carbonate and sulphide, in the
corresponding ores of zinc. During the reduction, being more volatile
than zinc, it distils over first.
In color the metal resembles tin, and is much more malleable and
ductile than zinc. It melts at 320° and boils at 77°.
It displaces hydrogen from dilute acids but is itself displaced from
solutions of its compounds by zinc, since it is less electro-positive.
GLUOINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY 651
Compounds of Cadmium. — The chloride crystallizes as a dihy-
drate CdCL,, 2H20, which is efflorescent and is not hydrolyzed during
dehydration or in solution. Zinc chloride (p. 647) is deliquescent and
is easily hydrolyzed. The halides are less ionized than are the corre-
sponding compounds of most other metals. The iodide, in particular,
seems to exist in solution as Cd.CdI4, and the complex anion gives
little ionic cadmium. On account of this fact even the sulphide can-
not be precipitated completely from a solution of the iodide. Con-
versely, hydriodic acid dissolves the sulphide to a much greater extent
than do other acids (see below).
The hydroxide is made by precipitation, and interacts, as a basic
oxide, with acids, but not at all with bases. It dissolves in ammonium
hydroxide, however, forming Cd(N"H3)4.(OH)2. The oxide is a brown
powder, obtained by heating the hydroxide, carbonate, or nitrate, or by
burning the metal.
The sulphate crystallizes from solution as 3CdS04,8H20, and is
not isomorphous with the sulphates of zinc and magnesium. Soluble
carbonates throw down the normal carbonate of cadmium CdC03, and
not a basic carbonate.
Hydrogen sulphide precipitates the yellow sulphide CdS even from
acid solutions of the salts. The substance is used as a pigment.
Since zinc sulphide is not formed under these conditions (p. 650),
the first part of the distillate from the reduction of the ore can be
dissolved in hydrochloric acid and the cadmium precipitated as
sidphide, while the zinc remains in solution. The sulphide of
cadmium, however, is less insoluble in water than are the sulphides
of copper and mercury, and therefore cannot be precipitated from a
strongly acid solution. We have thus a distinct gradation in the
solubility in water of various insoluble sulphides. The order of
increasing insolubility, and consequent difference in behavior towards
acids, is :
Barium, strontium, cal- -\
cium, and magnesium > attacked by water, and all acids.
sulphides : )
Ferrous sulphide : attacked even by acetic acid, but not by water.
Zinc sulphide : attacked by dilute active acids, but not by acetic acid,
attacked by concentrated active acids, but not by dilute
( attacked
Cadmium sulphide : -J . ,
. { attacked by fairly concentrated, actively oxidizing acids,
| like nitric acid, but not much by ordinary active acids.
Mercuric sulphide : hardly attacked at all, even by hot nitric acid.
652 INORGANIC CHEMISTRY
Analytical Reactions of Cadmium Compounds. — The cad-
mium ion Cd** is bivalent and colorless. The yellow cadmium sul-
phide is precipitated by hydrogen sulphide, even from acid solutions
of the salts. It is- also precipitated from solutions containing the
complex cation Cd(!N"H3)4" and the complex anion Cd(CN)4", as, for
example, from a solution made by adding excess of potassium cyanide
to cadmium chloride solution (K2.Cd(C]Sr)4). The latter property
enables cadmium to be separated from copper (p. 624). The white
hydroxide is thrown down by sodium hydroxide, and is not soluble in
excess of this reagent. It is not formed from solutions containing the
Cd(NH8)4** and Cd(CN)4" ions, and dissolves in ammonium hydroxide.
These and other precipitations are not complete when cadmium iodide
Cd.CdI4 is used.
MERCURY.
Chemical Relations of the Element. — Like copper, this element
enters into two series of compounds, the mercurous and the mercuric,
in which it is univalent and bivalent respectively. The mercurous
halides, like the cuprous halides (and the argentic halides), are insoluble
in water and are decomposed by light. There are, however, mercurous
as well as mercuric salts of oxygen acids. Both of the oxides, Hg/) and
HgO, are basic exclusively, but in a feeble degree. The hydroxides, like
silver hydroxide, are not stable, and lose water, giving the oxides. The
salts of both sets are markedly hydrolyzed by water, and basic salts are
therefore common. No carbonate is known. Mercury enters into the
anions of a number of complex salts, such as .HgCl4", HgI4", Hg(CN")4",
etc. It does not give complex cations with ammonia resembling those
of cadmium, copper, and silver (Cd(NH3)4", etc.), from which ammonia
is removed by heating, but instead forms a class of mercur-ammonium
compounds like Hg1JNH2Cl, all of which are insoluble. The mercuric
halides and cyanide show many peculiarities due to their being very
little ionized. Salts as a class are highly ionized bodies, and those of
mercury and, to a less degree, those of cadmium are the only con-
spicuous exceptions.
The mercury salts of volatile acids, like the corresponding salts of
ammonium (p. 421), can all be volatilized completely. Mercury vapor
and all mercury compounds are poisonous, the soluble ones more
markedly so than the insoluble ones.
The mercurous salts, as a rule, are formed when excess of mercury
is employed, and mercuric salts when excess of the oxidizing acid or
GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MEKCURY 653
other substance is present. Reducing agents turn mercuric into mer-
curous salts, and oxidizing agents do just the reverse.
As in the case of the cuprous compounds, it is a question whether
simple or multiple formulae, HgCl or Hg^lg, etc., should be employed
for mercurous salts. Pending the discovery of some basis for a decision,
the simple formulae are used here.
Occurrence and Isolation of the Metal. — Mercury occurs to
some extent native and to a larger extent as red, crystalline cinna-
bar, mercuric sulphide HgS. The chief mines are in Spain, California,
and Austria.
The liberation of the metal is easy, because the sulphide is decom-
posed at a high temperature, and the sulphur forms sulphur dioxide.
The mercury does not unite with oxygen, for the oxide decomposes
(p. 12) at 400-600° :
HgS + 02 -> Hg + S02.
In some places the ore is spread on perforated brick shelves in a verti-
cal furnace, and the gases pass through tortuous flues in which the
vapor of the metal condenses. The product is filtered through chamois-
skin. For separation from metallic impurities, like zinc, arsenic, and
tin, which are dissolved, it must be distilled. In the laboratory, where
mercury finds many applications, it becomes impure with use, and then
adheres to glass, and does not run freely in spherical droplets. For
purification it is placed along with a little diluted nitric acid in a
separatory funnel (Fig. 56, p. 147), and kept in continual agitation by
means of a current of air drawn or blown through the mass. By this
treatment foreign metals, such as sodium or zinc, nearly all of which
are much more active than mercury (cf. p. 362), are converted into
nitrates. Pure, dry mercury can be drawn off, when needed, at the
bottom. If a high degree of purity is required, the product must be
distilled in vacua.
Physical Properties. — Mercury or quicksilver (NL. hydrargy-
rum, from Gk. v8o>p, water, and a/ayupos, silver) is a silver-white liquid.
At — 39.5° it freezes, and at 357° it boils. At ordinary temperatures it
has a measurable vapor tension, at 15° 0.0008 mm. and at 99° 0.26 mm.
The vapor is colorless, does not conduct electricity, and is rnonatomic.
A gold-leaf suspended over mercury becomes amalgamated, since the
solution of gold in mercury has a vapor tension smaller than that of
pure mercury (p. 161).
654 INORGANIC CHEMISTRY
On account of its high specific gravity (13.6, at 0°) and low vapor
tension, the metal is employed for filling barometers and manometers.
Its uniform expansion favors its use in thermometers. The tendency
to form amalgams, which it exhibits towards all the familiar metals with
the exception of iron and platinum (the latter, however, is " wet " by
it), is taken advantage of in various ways. Sodium amalgam (p. 569),
which is solid when the sodium exceeds 2 per cent, behaves like free
sodium, but with moderated activity. A layer of mercury on the zinc
plates of batteries reduces the action of the acid on the zinc, while the
cells are not in use. Mixtures of mercury with powdered cadmium
and a small proportion of copper, quickly form solid amalgams, and are
used by dentists. The employment of mercury in the extraction of gold
by washing has already been mentioned (p. 636). Mirrors backed with
a tin-mercury amalgam have been almost completely displaced by silvered
mirrors (p. 632).
Chemical Properties. — When kept at a temperature near to its
boiling-point, mercury combines slowly with oxygen. Its inactivity
towards oxygen when cold places it next to the noble metals. On
account of its general inactivity, it is used in the laboratory for confin-
ing gases. It does interact with hydrogen sulphide and hydrogen
iodide, however (cf. Silver, p. 628). Mercury does not displace hydro-
gen from dilute acids (p. 362), but with oxidizing acids like nitric acid
and hot concentrated sulphuric acid, the nitrates and sulphates are
formed. With excess of mercury, the mercurous salts, and with ex-
cess of the hot acid, the mercuric salts, are produced. When triturated,
for example with milk-sugar, mercury is divided into minute droplets
with relatively large surface. In this form it is used in medicine
("blue pills"), and shows an activity which is entirely wanting in
larger masses.
Mercnrous Chloride. — This salt (calomel) is obtained as a white
powder by precipitation from solutions of mercurous salts. It is man-
ufactured by subliming mercuric chloride with mercury :
HgCL, + Hg -> 2HgCl,
or more usually by subliming a mixture of mercuric sulphate, made as
described above, with mercury and common salt. It is deposited on the
cool part of the vessel as a fibrous crystalline mass, or, when the vapor is
led into a large chamber, as a fine powder. It is slowly affected by light
GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY 655
just as silver chloride is. Here, however, the chlorine which is released
combines with another molecule of the salt to form mercuric chloride.
Since the vapor pressure of calomel reaches 760 mm. before the tem-
perature has risen to the melting-point, the compound sublimes at
atmospheric pressure without melting. Its vapor density corresponds
to the formula HgCl, but the vapor can be shown to contain more or
less Hg + HgCL2. Since this change does not alter the average
molecular weight, and the action is reversed when the temperature
falls (ef. Ammonium chloride, p. 421), it has not been found possible
to determine whether little or much of the calomel is thus affected.
The substance is used in medicine on account of its tendency to stim-
ulate all organs producing secretions.
Mercuric Chloride. — By direct union with chlorine the mercuric
salt, corrosive sublimate HgCl^, is formed. It is usually manufac-
tured by subliming mercuric sulphate with common salt, and crystal-
lizes in white, rhombic prisms. It melts at 265° and boils at 307°. At
20° one hundred parts of water dissolve only 7.4 parts of the salt, and
at 100°, 54 parts. It is more soluble in alcohol and in ether. The
aqueous solution is slightly acid in reaction. The salt is easily reduced
to mercurous chloride. When excess of stannous chloride is added to
the solution, the white precipitate of calomel, first formed, passes into
a heavy gray precipitate of finely divided mercury :
2HgCL, + SnCL, -> SnCl4 + 2HgCl,
2HgCl + SnCl, -». SnCl4 + 2Hg.
The halides of mercury are very little ionized in solution, the
bromide and iodide even less so than the chloride. Hence these salts
are little affected by sulphuric acid or nitric acid. For example, the chlo-
rine and nitrosyl chloride which hydrochloric acid forms with the nitric
acid are not observed when mercuric chloride is added to this acid. On
this account, too, the hydrolysis of the chloride is much less than that
of the nitrate. There is a tendency also to the formation of complex
salts, so that the addition of sodium chloride increases the solubility
in water and renders the solution neutral, Nad, HgCl2 or NaHgCl beino-
formed. The complex salts, like K.HgCl3,H2O, H2.HgCl4, 7H,0, and
NH^HgClgjHjO, are easily made by crystallization from solution. The
anions are relatively highly ionized, however, and the behavior is inter-
mediate between that of complex salts and double salts (p. 537).
656 INORGANIC CHEMISTRY
Corrosive sublimate, when taken internally, is extremely poisonous.
A very dilute solution is used in surgery to destroy lower organisms
and thus prevent infection of wounds. The pharmaceutical tabloids
of mercuric chloride contain sodium chloride, because, although the
latter diminishes the activity of the compound, it also does away with
the formation of insoluble chlorides and hastens solution. Mercuric
chloride acts also as a preservative of zoological materials, forming in-
soluble compounds with albumins, and preventing their decay. For
the same reason, albumin (white of an egg) is given as an antidote in
cases of sublimate poisoning.
The Iodides of Mercury. — Mercurous iodide is formed by rub-
bing iodine with excess of mercury. It also appears as a greenish-
yellow precipitate when potassium iodide is added to a solution of a
mercurous salt. The compound decomposes spontaneously into mer-
cury and mercuric iodide. The decomposition is much hastened by the
use of excess, of potassium iodide, which combines with and removes the
mercuric iodide (see below) :
Mercuric iodide is obtained by direct union of mercury with excess
of iodine, or by addition of potassium iodide to a solution of a mercuric
salt. It is a scarlet powder, insoluble in water, but soluble in alcohol
and ether. It interacts with excess of potassium iodide, forming the
soluble, colorless potassium mercuri-iodide K2.HgI4 with which many
precipitants fail to give mercury compounds. When heated above
116.5° it turns slowly into a yellow modification, and at 223° this new
form melts. On being cooled, it freezes first in the tetragonal yellow
form, and below 116.5°, especially if touched with a glass rod, it
changes into the red, monoclinic variety with evolution of heat. Sul-
phur (p. 368) and ammonium nitrate (p. 565) show similar transition
points. When the vapor of the compound is cooled, it first forms thin
scales of the yellow form, which is the unstable one at low tempera-
tures, and these turn red when touched. Similarly, precipitation gives
first the yellow variety, which presently becomes red (cf. Transforma-
tion by steps, p. 453).
The Oxides. — When bases are added to solutions of mercurous
salts, a brownish-black powder is thrown down, which, when dried, is
found to be mercurous oxide Hg20. The hydroxide is doubtless
GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY 657
formed transitorily and then loses water (cf. Silver oxide, p. 630).
Under the influence of light or gentle heat (100°), this compound re-
solves itself into mercuric oxide and mercury.
Mercuric oxide is formed as a red, crystalline powder, when mer-
cury is heated in air near to 357°, but is usually made by decomposing
the nitrate. Commercial specimens, incompletely decomposed, thus
frequently give some nitric oxide when heated. It is formed also as a
yellow powder by adding bases to solutions of mercuric salts. It is
contended by some chemists that the difference in activity between the
red and yellow forms is due solely to the finer state of division of the
latter, and by others it is maintained that the substance is dimorphous
(p. 369) and that two distirict varieties exist.
The Nitrates. — The mercurous salt is formed by the action of
cold, diluted nitric acid upon excess of mercury. It forms monoclinic
crystals of a hydrate HgN03,H20. It is hydrolyzed, slowly by cold,
and rapidly by warm water, giving a basic nitrate :
2HgN03 + H20 «=* HN03 + Hg2(OH)N03 J .
On this account a clear solution can be made only when some nitric
acid is added. Free mercury is also kept in the solution to reduce
mercuric nitrate, which is formed by atmospheric oxidation :
Hg(NO.)f + Hg->2HgNO, or Hg"+ Hg -> 2Hg\
Mercuric nitrate is produced by using excess of warm, concen-
trated nitric acid. It forms rhombic tables of Hg(N03)2, 8H20. The
aqueous solution is strongly acid, and deposits a yellowish, crystalline,
basic nitrate Hg3(OH)20(NO3)2. The action is reversed by adding
nitric acid.
Sulphides of Mercury. — Mercurous sulphide Hg2S is formed
by precipitation from mercurous salts, but is stable only below —10°.
Above this temperature it decomposes into mercury and mercuric
sulphide.
Crystallized mercuric sulphide occurs as cinnabar, and is red.
When formed by precipitation with hydrogen sulphide, or by rubbing
together mercury and sulphur, it is black and amorphous. By subli-
mation, in the course of which it dissociates, the black form gives the
red, crystalline one. When allowed to stand under a solution of so-
dium sulphide, the black form is slowly transformed into the red. This
658 INORGANIC CHEMISTRY
shows that, as we should expect, the red form is the more stable,
possesses less energy, and is less soluble at ordinary temperatures.
The change is effected by intermediate formation of a complex sul-
phide, the solution, when saturated toward the less stable black sul-
phide, being supersaturated toward the more stable red one. A white,
crystalline sodium mercuri-sulphide NagHgSo, 8H20 can, in fact, be ob-
tained from the solution.
The black and the red varieties do not interact with concentrated
acids, and are even unaffected by boiling nitric acid, which oxidizes
most sulphides readily. They are, therefore, less soluble in water
than cupric sulphide (pp. 600, 651), and much less so than sulphide
of cadmium. They are attacked, however, 'by aqua regia.
The red form of the sulphide is used in making paint (vermilion).
The color is more permanent than that of red lead (Pb304), because re-
ducing gases (e.g. SO2), acids (e.y. H2S04), and hydrogen sulphide,
which are present in the air, do not affect it. It is not stable, how-
ever, when applied to metals, since iron, zinc, etc., all displace mer-
cury from combination, and in these cases, therefore, red lead is pre-
ferred.
Mercuric Cyanide. — This salt is made by treating precipitated
mercuric oxide with hydrocyanic acid, and is obtained in square-pris-
matic crystals. When heated it gives off cyanogen : Hg(CN)2 — >
Hg -f- C2N2, and is a convenient source of this gas (p. 507). The com-
pound is soluble in alcohol, ether, and water. In solution in water
it is so little ionized that the freezing-point of the solution is normal
(p. 292), and many reagents fail to show the presence of either ion.
Thus, with silver nitrate no silver cyanide is precipitated, and with
a base no mercuric oxide. With potassium cyanide it forms a complex
cyanide K2.Hg(CN)4. Hydrogen sulphide throws down the sulphide
from both the simple and the complex cyanides.
The Fulminate and Thiocyanate- — Mercuric fulminate
Hg(ONC)2 is obtained as a white precipitate when mercury is treated
with nitric acid, and alcohol is added to the solution. It decomposes
suddenly when struck, and is used in making percussion caps. The
thiocyanate Hg(SCN)2 is precipitated when potassium thiocyanate
K(SCN) is added to a solution of a mercuric salt. When formed into
little balls and burned in the air, the substance leaves a curiously
voluminous ash ("Pharaoh's serpents").
GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY 659
Mer >cur '-ammonium Compounds. — When ammonium hydrox-
ide is added to a solution of a mercuric salt, a white substance, of a type
which we have not previously encountered, is thrown down. Mercuric
chloride gives HgNH2Cl, commonly called " infusible white precipi-
tate," and the nitrate gives HgNH2NO8 :
HgCL, -+- 2NH3 -> HgNH.Cl + NH4C1.
The substance may be regarded as being derived from ammonium
chloride (p. 421) by the displacement of 2H by Hg :
/H
Hg:N-H.
\C1
It would thus be named mercuf -ammonium chloride.
When the mercuric chloride is added to a boiling solution of
ammonium chloride containing ammonium hydroxide, mercur-diam-
monium chloride Hg(NH3Cl)2, " fusible white precipitate," appears.
The addition of ammonium hydroxide to a solution of potassium
mercuri-iodide K2HgI4 gives rise to a third type of compound,
dimercur-ammonium iodide Hg2NI,H2O, which appears as a brown
precipitate : .
2HgI2 + 4NH3 -» Hg2NI + 3NHJ.
A solution of potassium mercuri-iodide containing potassium hydroxide,
Nessler'a reagent, becomes distinctly yellow with traces of ammonia,
and brown with larger amounts, and is, therefore, a valuable reagent
for detecting traces of this base.
When calomel is treated with ammonium hydroxide, it turns into a
black, insoluble body. This appears to be a mixture of free mercury,
to which it owes its dark color, and " infusible white precipitate,"
Hg + HgN"H,Cl. To this reaction calomel owes its name (Gk.
KuAo/Ae'Aa?, beautiful black). Mercurqus nitrate gives a black, insoluble
mixture, Hg + HgNH2N03. Calomel, when dry. absorbs ammonia gas,
forming a molecular compound of the common type HgCl, NH3. This
substance loses the ammonia again when the pressure is reduced. The
other compounds described above, on the other hand, do not contain
nitrogen and hydrogen in the proportions necessary to form ammonia
and are stable. Hence they are necessarily to be regarded as belong-
ing to a different type.
Analytical Reactions of Mercury Compotiuds. — The two ionic
forms of the element, monomercurion Hg* and dimercurion Hg ", are
660 INORGANIC CHEMISTRY
both colorless, but their chemical behavior is entirely different. Both
give the black sulphide HgS, which is insoluble in acids and other
solvents of mercury salts. Monomercurion gives the insoluble, -white
chloride, the black oxide, and a black mixture with ammonium
hydroxide. Dimercurion gives a soluble chloride, a yellow, insoluble
oxide, and a white precipitate with ammonium hydroxide. The be-
havior with stannous chloride (p. 655) is characteristic of mercuric salts.
With potassium iodide the behavior of the two ions is quite different
(p. 656). The more active metals displace mercury from all compounds.
This displacement is employed as a test for compounds of mercury.
Copper is used by preference as the displacing metal because the mer-
cury is easily seen on its surface.
Salts of mercury are volatile. When heated in a tube with sodium
carbonate, they also give a sublimate of metallic mercury.
THE RECOGNITION OF CATIONS IN QUALITATIVE ANALYSIS.
" Wet-way " analysis consists in recognizing the various positive and
negative ions present in a solution (p. 343). In discussing hydrogen
sulphide (p. 375), it was stated that the sulphides might be divided
into three classes according to their behavior towards water and acids.
Now these differences in behavior furnish us with a basis for distin-
guishing the cations present in a solution. Since the properties of a
number of sulphides and other compounds of the metals have been
studied in recent chapters, it is possible to make a more complete
statement.
The following plan, taken in conjunction with the statements in
the context, shows how a single cation may be identified, and how,
when several cations are present, a separation preparatory to identifi-
cation may be effected. What will be said applies only to the case of
a solution containing salts like the chlorides, nitrates, or sulphates of
one or more cations, and leaves the oxalates (p. 601), phosphates,
cyanides, and some other salts, out of consideration.
Group 1. — It is usual to add, first, hydrochloric acid, to find out
whether cations giving insoluble chlorides are present. Argentic,
mercurous, and plumbic salts give the white AgCl, HgCl, and PbCl2
respectively. (For the further recognition of each, see p. 635.) Fil-
tration eliminates the precipitate, if there is any.
Group 2. — A free, active acid being now present, hydrogen sulphide
is led into the solution. The sulphides insoluble in active acids,
RECOGNITION OF CATIONS 661
namely, HgS, CuS, PbS, Bi2S3, CdS, As2S3, Sb2S3, SnS, are therefore
thrown down. The first four are black or brown, the next two are
yellow, and the last two are orange and brown respectively. A dark-
colored substance will naturally obscure one of lighter color, if more
than one is present. If too much acid is used, the precipitation of
several of the sulphides will be incomplete (p. 651) ; if too little, zinc
sulphide may come down (p. 649). Filtration again eliminates the
precipitate.
This group is easily subdivided. Any or all of the last three sul-
phides will pass into solution when warmed with yellow ammonium
sulphide, for they give soluble complex sulphides similar to potassium
sulphaurate (p. 638). The first five sulphides, or any of them, will be
unaffected. On the other hand, these five sulphides, with the exception
of HgS, will interact with hot nitric acid (p. 658). Other reactions
described in the context are then used to distinguish between, or, if
there is a mixture, to separate, the members of the sub-groups.
Group 3. — The solution (filtrate) is now neutralized with ammonium
hydroxide, and ammonium sulphide is added. Some ammonium chlo-
ride is also used, to prevent the precipitation of magnesium hydroxide
(p. 644), which, in any event, would be incomplete. The sulphides
which are insoluble in water, and are not hydrolyzed by it, now appear.
They are FeS, CoS, NiS, all black, MnS, H2O, and ZnS, which are pink
and white respectively. There are precipitated also the hydroxides of
chromium and of aluminium, Cr(OH)3 arid A1(OH)3, because, although
their sulphides are hydrolyzed by water, the hydroxides are formed by
the hydroxidion in the ammonium sulphide solution. They are too
insoluble to behave like magnesium hydroxide (p. 646) by dissolving
in salts of ammonium. They also form no complex metal-ammonia
ion, as does zinc (p. 648). The sulphides of nickel and cobalt re-
semble the sulphide of zinc in being precipitated by hydrogen sulphide
when acetic acid is the only acid present. The other sulphides inter-
act even with acetic acid (p. 651).
Another plan is to oxidize the iron, if present, and use ammonium
chloride and ammonium hydroxide instead of ammonium sulphide.
The hydroxides of the trivalent elements, Fe(OH)3, Cr(OH)3, A1(OH)3,
can be precipitated by excess of ammonium hydroxide even when salts
of ammonium are present. Those of the bivalent metals, Mn(OH)2,
Fe(OH)2, Zn(OH)2, Ni(OH)2, Co(OH)2, resemble magnesium hydroxide
(p. 644), and, of these, the last three resemble also zinc hydroxide
(p. 648), and so cannot be precipitated. After filtration, ammonium
662 INORGANIC CHEMISTRY
sulphide now throws down the sulphides of the five bivalent metals
(for a third plan, see Chemical relations of aluminium).
Group 4. — After filtration from members of the iron group, if any
were present, ammonium carbonate is added, and precipitates the re-
maining metals whose carbonates are insoluble, BaCO3, SrCO3, CaCO3>
with the exception of magnesium (p. 644).
By addition of sodium phosphate to a portion of the filtrate, mag-
nesium, if present, now conies out in the form NH4MgPO4. There re-
main in solution only salts of potassium, sodium, and ammonium.
Since only ammonium compounds, and other substances which can be
volatilized have been added, evaporation and ignition of the residue
leaves the salts of the two metals. If no other metallic elements have
been shown to be present, it saves time to examine a fresh portion of the
original material. Salts of ammonium must also be sought in a fresh
sample by the usual test (p. 421).
The following simple compounds are soluble, but are so little
ionized that their solutions do not show all the reactions of both of
the ions: NH4OH, H2S, HNC, H2CO8, HgCl2, Hg(CN)2, Fe(CNS)3,
Fe(C2H302)2. With a number of others, for example CdI2, the actions
are incomplete for the same reason. Complex compounds, as we have
seen, give complex ions, and the latter are usually so little resolved
into simpler ions that the latter cannot be discovered by all the usual
tests. Thus : K.Ag(CN)2 gives K* and Ag(CN)/, but very little Ag*
and CW (p. 629) ; Cu(KH3)2.Cl2 gives much Cu(NH3)2", and Cl' but
very little Cu". The individual cases are fully described in the
context.
Exercises. — 1. What is the numerical value (a) of the solubility
product of magnesium hydroxide, (b) of the concentration of hydrox-
idion given by it and by normal ammonium hydroxide respectively
(p. 544) ? Will normal concentration of ammonium chloride suffice to
reduce the latter below the former ?
2. Why should we, perhaps, expect ammonium sulphide solution to
precipitate magnesium hydroxide, and why does it not do so ?
3. What volume of air is required to oxidize one formula-weight
of zinc sulphide to ZnO and SO,, and what volume of sulphur dioxide
is produced? Is the product more or less diluted with nitrogen than
when pure sulphur is burned, and by how much ?
4. Make equations showing (a) the effect of heating zinc chloride
with cobalt nitrate (Co(N03)2) in the Bunsen flame (p. 648), (&) the
RECOGNITION OF CATIONS 663
action of hydrogen sulphide on sodium zincate, (c) the actions of con-
centrated nitric acid and of concentrated sulphuric acid on mercury.
5. What is the distinction between a solid isomorphous mixture of
two salts and a double salt ?
6. What kind of salts might take the place of sodium acetate in
the precipitation of zinc sulphide (p. 650) ? Give examples.
7. Compare the amalgamation of a gold-leaf by mercury vapor with
the phenomenon of deliquescence (p. 162).
8. If the scheme for the recognition of cations (p. 660) were ap-
plied to solutions prepared from materials containing (a) calcium oxa-
late and (#) potassium argenticyanide, at what stage and how would
the presence of each of these substances affect the normal order ?
9. Why do none of the salts of the elements in this family give
recognizable effects with the borax bead ?
ELECTROMOTIVE CHEMISTRY
WE have seen that chemical changes which proceed spontaneously
liberate some form of energy. As a rule, heat is developed ; but with
special arrangement of the apparatus (p. 20) so that it takes the
form of a battery -cell, an equivalent amount of electricity is obtained
instead. To avoid suggesting that this energy comes from nothing,
we say that the original system contained a certain amount of free, or
available, chemical energy and that this has been transformed, con-
comitantly with the chemical change, into an equivalent amount of
heat, or of electrical energy, as the case may be. Nor is the alterna-
tive of producing electrical energy available only when the action
resembles the displacement of hydrogen by zinc, as in our illustration.
Most changes between ionogens, including oxidations and double
decompositions, may be adapted so as to deliver this form of energy.
It need hardly be added that, since the transformation of chemically
equivalent amounts of different sets of substances produces very dif-
ferent quantities of heat, so it produces also correspondingly different
amounts of electrical energy. Thus the original free chemical energy
may, theoretically, be measured by either method. In practice,
however, the thermochemical plan fails entirely in many cases (cf.
pp. 27, 79), and the electrical is, as we shall see, often much more
instructive. The study of what, to parody the phraseology of
thermochemistry, we might call " exoelectrical" actions, thus resolves
itself into constructing experimental battery-cells involving all kinds
of chemical changes, and studying the electric currents which are set
in motion by the progress of the changes. We have therefore named
this branch of the science electromotive chemistry.
In addition to its significance theoretically, electromotive chem-
istry has recently acquired great commercial importance because of the
rapid multiplication of electro-chemical industries. It is true that
the majority of the actions used in these industries are electrolytic
(eudoelectrical), and that this sort of change is the precise inverse of
the other, since in it electricity is consumed instead of set in motion,
664
ELECTROMOTIVE CHEMISTRY 665
but it is also true that neither variety can be understood without a
study of both.
Factors and Units of Electrical Energy. — On account of the
close relation between electromotive chemistry and electrolysis, parts
of the former subject were anticipated when the latter was discussed
(pp. 321-325). These pages should now be re-read attentively. In
particular, it must be recalled that a quantity of electrical energy is
expressed by two factors. One is called the quantity of electricity, and
is measured in coulombs. The other is called the electromotive force
in the case of a current, or, when a current is not flowing or is not be-
ing considered, the difference in potential, and is expressed in volts.
Just as in electrolysis chemically equivalent quantities of elements or
ions, in being liberated from solutions of different substances, use up
equal quantities of electricity (p. 317), so in a battery-cell the inter-
action of chemically equivalent amounts of different sets of substances
produces equal quantities of electricity (p. 321). Likewise, just as in
the former case different amounts of electrical energy (p. 323), and
therefore different electromotive forces, are required to produce in
different solutions equivalent amounts of chemical change (p. 324), so
in the latter case different amounts of electrical energy are generated
by the complete interaction of chemical equivalents of different sets of
substances, and therefore diverse differences in potential are created
and currents of different electromotive force are produced. The electri-
cal energy used in the former case or produced in the latter is expressed
by the product of the factors :
No. of coulombs x No. of volts = Quant, of elect, energy (in joules, p. 25).
If we consider the time occupied by either process, and wish to express
the rate at which the energy is consumed or produced, we regard 1
coulomb per second (1 ampere] as the unit.
Hence :
No. of amperes x No. ) _ ( Rate of production or consumption of elec-
of volts ) ( trical energy (in joules per sec. = watts).
The erg (p. 25) is so small as a unit of energy or work, that the
joule (= 10,000,000 ergs) and the kilojoule (1000 joules) are more often
employed. Similarly, the rate at which the energy is delivered or used
(the power) is expressed by the watt ( = 10,000,000 ergs per sec.) or
the kilowatt (1000 watts). The horse-power is 736 watts. Thus
666 INORGANIC CHEMISTRY
the electrolysis of one formula-weight of hydrochloric acid (36.5 g.)
requires 96,540 coulombs (p. 323) and an E.M.F. of 1.41 volts (p. 324).
The electrical energy needed to perform this work is therefore 96,540
X 1.41 = 136, 121 joules.
Displacement Cells. — While various kinds of chemical changes
may be arranged to deliver electricity, we shall confine our attention
almost entirely to actions in which one free substance displaces
another from combination. Amongst the metals the order of displa-
cing power, is as follows (p. 362) :
Mg-Zn-Cd-Fe-Sn-Pb-H-Cu-Hg-Ag-Pt-Au.
We have already noted many of the single facts expressed by this list.
Thus, metallic copper precipitates metallic mercury, silver, platinum,
and gold from solutions which contain salts of these elements, but it
is in general entirely inactive in solutions containing salts of the
elements which precede it in the list. Conversely, all the metals pre-
ceding copper will severally precipitate metallic copper from a solution
of a cupric salt, but the metals following copper do not affect the
solution at all. Furthermore, the speed with which the precipitation
is effected and the amount of heat evolved are greater when copper
precipitates. gold than when it precipitates silver : they are also greater
when magnesium precipitates copper than when tin does so. The
same statements apply, mutatis mutandis, to each of the other ele-
ments. This is the behavior when the experiment is made in the
ordinary way.
One preliminary condition is indispensable when electricity is to
be obtained. When the zinc was placed in sulphuric acid and con-
nected with a platinum wire (p. 20), there was nothing to prevent the
metal from interacting directly with the acid, and generating heat. In
fact, it did so act. Hydrogen appeared on the zinc itself, as well as on
the platinum, and both the heat and the electricity would have had to
be counted if a measurement or study of the relation between the
amount of zinc and acid used and the quantity of energy produced had
been in question. Since this is precisely the present problem, all
the energy must now be secured as electricity, and the only way to
accomplish this is to prevent the mingling of the interacting materials.
Paradoxical as it may seem, it is easily possible to get the electricity
and yet fulfil this essential condition. The plan in all battery-cells is
to place the one substance in or around one pole, and the other substance
ELECTROMOTIVE CHEMISTRY 667
around or in the other pole, and to separate the substances by a porous
wall or equivalent arrangement. The figure (Fig. 102) shows dia-
grammatically the plan that will serve for the generation of electricity
by the use of any interaction.
Suppose the action is a displacement, and that, for example, the
metal is zinc and the salt cupric sulphate in solution. The pole on the
left is metallic zinc ; the solution to the right of the porous partition
contains the cupric sulphate. To complete the arrangement, a pole
made of some conductor is needed
on the right, and a conducting solu-
tion on the left. For these, any
substances may be chosen, provided
only that they do not interact with
the adjacent material. If they do,
part of the two main substances
will be used up, with generation
of heat instead of electricity. For
the pole we may take copper itself
or any metal following it in the
series ; and copper is less expensive
than any of these. For the solution
we may take a salt of zinc or of any
metal preceding zinc in the series, provided that the salt chosen will
not interact with the cupric sulphate which it meets inside the porous
wall. Zinc sulphate or sodium chloride will serve the purpose. The
function of the porous partition is to permit the migration of ions, but
prevent a general mixing of the materials. It is evident that, when
the corresponding conditions are observed, any other pair of elements
from the series may be chosen for study. In practice, it is common to
place the solution of the substance corresponding to the cupric sulphate
along with its electrode in an outer jar, and a rod of the more active
metal along with its conducting fluid in an inner jar made of porous
earthenware. The following are the phenomena observed :
1. The pole on the left becomes charged negatively on account of the
departure of positively charged ions from its surface, for this metal
(zinc, for example) goes into solution. The pole on the right becomes
positively charged on account of the discharge of positive ions, and de-
position of the metal (copper, for example) on its surface. A current,
therefore, flows from + to — through the wire, and passes in the form
of migrating ions from -— to + through the liquids. The negative
FIG. 102.
668 INORGANIC CHEMISTRY
ions simultaneously drift away from the pole where the solution
is losing metal, towards the pole where it is gaining metal, and so
the balance of the two sorts of ions is preserved in every part of the
liquids. The pole on the left is the anode, that on the right the cath-
ode. This whole state of affairs continues until either all the copper
is precipitated, or all the zinc is consumed.
2. The difference in potential (E.M.F.) is not affected by changes in
the size or shape of the poles or in the amounts of the solutions, provided
the materials are not changed. It is very noticeably affected, however,
by alterations in the concentrations of the solutions. In particular,
an increase in the concentration of the ions round the cathode (the
Cu" ions, for example) increases the difference in potential. It seems
to assist the action by helping to force the ions out of the solution.
3. In general, the magnitude of the electromotive force is greater
the farther the chosen metals are removed from one another in the
series. Zinc with cupric sulphate gives a greater difference in poten-
tial than cadmium with the same salt, but a smaller one than zinc with
silver nitrate. The sum of the differences in potential of the pairs
Zn-Cd", Cd-Cu", Cu-Ag*, is exactly equal to the difference of the pair
Zn-Ag*, provided solutions of equivalent ionic concentration are
employed.
4. The quantity of electricity produced depends on the number of
equivalents of material consumed. The strength of the current (the
rate at which the electricity is developed) depends on the surface of
the poles and other circumstances which influence the rate at which
the materials are able to undergo chemical change. One gram-equiva-
lent of the active metal (say, 32.7 g. of zinc) in becoming ionized will
set in motion 96,540 coulombs, and deposit one gram-equivalent of the
less active metal (say, 31.8 g. of copper). If this amount were to be
ionized during every thirty minutes ( = 1800 seconds) — and plates of
great surface would be required to make this possible — the current
strength would be 96,540/1800 = 53.6 amperes.
Potential Differences Produced by the Metals Singly. — The
facts just discussed may be stated in various ways. For example, we
may say that, since every metal (save the last of the series) displaces
some other metal or metals, each metal has a distinct tendency to
become ionic. We may picture this as a pressure or tension driving
the particles into solution, and actually resulting in ionization when
the necessary electricity is available. This pressure must evidently
ELECTROMOTIVE CHEMISTRY 669
be greatest with magnesium, and least with gold. On the other hand,
we may suppose a converse tendency of the ions to force themselves
out of solution and to deposit themselves on the cathode, this tendency
actually becoming continuously operative when means of disposing of
the electric charges is provided. This tendency is evidently greatest
with gold, and least with magnesium. Now the use of cells like those
described above does not enable us to verify these hypotheses, because
in each cell two operations are going on simultaneously, and the electri-
cal effects are the joint result of both. Fortunately it has been found
possible to observe and measure the electrical effects at each end of
the cell separately. It is found, in fact, that the difference in poten-
tial between the poles is for the most part made up of the differ-
ences in potential between each pole and the liquid in which it is
immersed.
After the cell has been in operation for some time, the cathode
system consists of a pole covered with deposited metal (say copper) in
contact with a solution containing ions of this same element. At this
end the ions are diminishing in number, and the amount of free metal
is increasing. The anode system likewise consists of a metal (say
zinc) in contact with a solution containing ions of the same element.
Here the number of ions is increasing, and the metal is wearing away.
The ions of the foreign salt, if such a salt was introduced at first, need
not be considered. Thus, for the purpose of making a strict compari-
son, we may immerse each metal of the series in a solution of one of
its own salts, taking the latter of such strength that there is normal
concentration of the metallion, and measure the difference in potential
between the metal and the solution.
The following table (p. 670) contains data obtained in just this way.
The number opposite each of the twenty-six metals (including hydrogen)
represents the potential difference. The sign preceding the number
indicates the nature of the electrical charge borne by the solution.
Two or three examples will make the meaning of the table clearer.
Opposite Mg, we find + 1.21. This means that when the metal mag-
nesium is immersed in a solution of a magnesium salt containing a
normal concentration of Mg", the solution becomes positively charged
(the metal, negatively) and that the difference in potential between
metal and solution is 1.21 volts. Similarly in the case of zinc, the
solution is positive, but the potential difference is smaller. With
copper the solution is negative (and the metal positive). With silver
the solution is negative, and more strongly so than with copper. Cobalt
670
INORGANIC CHEMISTRY
is near the border line, the normal solution of cobaltion being very
slightly negative.
POTENTIAL DIFFERENCES IN VOLTS FOR NORMAL SOLUTIONS
OF CATIONS.
K
(+2.9 )
Ni
-0.04
Na
( + 2.54)
Sn (Sn")
- 0.08 ?
Ba
( + 2.54)
Pb
- 0.13
Sr
(+2.49)
H
-0.28
Ca
(+ 2.28)
Cu (Cu")
- 0.61
Mg
+ 1.21
As
- 0.62 ?
Al
+ 1.00
Bi
- 0.67 ?
Mn
+ 0.80
Sb
-0.74?
Zn
+ 0.49
Hg (Hg")
- 1.03
Cd
+ 0.14
Ag
- 1.06
Fe (Fe")
+ 0.06
Pd
- 1.07?
Tl
+ 0.04
Pt
-1.14?
Co
- 0.04
Au
- 1.35 ?
For a hydrogen pole a piece of palladium saturated with hydrogen
gas is used. The values for the metals which decompose water with
ease cannot be ascertained by direct observation. The numbers in
parentheses are calculated from the heats of ionization and serve
simply to indicate the order of these elements. An interrogation
point indicates that the value is uncertain.
These facts enable us to state in more definite terms the formula-
tive hypothesis foreshadowed above. It was first put forward by
Kernst.
Every metal has a certain solution tension or pressure tending to
drive it into solution (in ionic form,* of course, since it is not soluble
otherwise). The value of this pressure becomes rapidly less as we
pass through the series from magnesium to gold. If the ions of the
same metal are already present, they tend to give up their electrical
charges and deposit themselves upon the metal. These two
tendencies oppose one another just as solution pressure and
osmotic pressure oppose one another in the ordinary process
* The idea of the charge of electricity is apt to interfere with the ready accept-
ance of this hypothesis. If it is remembered that the ionic form of an element is
simply an allotropic modification (Ostwald) with a different amount of available
energy, the difficulty disappears. In the ionic allotrope the free energy is some-
times greater (cobalt to platinum) and sometimes less (potassium to thallium) than
in the free element.
ELECTROMOTIVE CHEMISTRY 671
of dissolving any substance (p. 152). When the tendency 'of
the ions to deposit themselves is the greater of the two, a very minute
excess of deposition over solution occurs, and thus the solution has, as
a whole, a negative charge (having lost some positive ions), and the
metal has a positive charge (having acquired it from the deposit of a
few ions). This is the case with gold and the metals as far up the
list as cobalt. When, on the other hand, the solution pressure of the
metal is the greater of the two, the solution acquires a very slight excess
of positive ions, and is, therefore, positively charged when compared
with the metal. This is the case from potassium down to thallium.
The measure of the " tendency of the ions to deposit themselves "
is simply the osmotic pressure of the metallions. We perceive this to
be the case, for, when we take a stronger solution of the salt and there-
fore an increased osmotic pressure of the ions, an instant effect is pro-
duced. The solution becomes less positive, or more negative, as the
case may be. Evidently the solution pressure of each of the metals
near to cobalt is almost exactly balanced by the osmotic pressure of a
normal solution of the ions composed of the same metal. This pressure,
for a univalent metal, is 22.4 atmospheres (p. 289), and for a bivalent
metal 11.2 atmospheres. The metals above cobalt have solution press-
ures higher and higher above this norm ; those below cobalt have solu-
tion pressures farther and farther below it. The effect of changing the"
osmotic pressure is independent of the particular substances used, and
depends only on the valence. When the concentration of the metallion
becomes WN, 0.058 volts must be subtracted (algebraically) from the
potential (see above table) of the liquid, if the metallion is univalent.
If it is w-valent, 0.058/w must be subtracted. When the solution is 0.1
N, 0.058/w volt must be added; when it is 0.01 JV, 2 x 0.058 must be
added, and so forth. Thus, zinc with decinormal zincion gives + 0.49
-f 0.058/2 = .52 volts, approximately ; silver with centinormal argent-
ion gives —1.05 + (2 x 0.058) = —.93 volts, approximately. And, in
general, if c be the equivalent concentration of the metallion in the
liquid under consideration, and irc the electrical potential of that liquid,
while Trjyr is the potential of the liquid containing N metallion,
, 0.058 . 1
^--nH--^ ^g-.
Application to Cells. — When, now, a cell with two poles and two
metallions is set up, we can tell from the above table what the differ-
672
INORGANIC CHEMISTRY
ence in potential between the two poles will be. We may regard the
two systems — the anodic and cathodic — as working against each other.
Each metal tends to project its ions into the solution and to generate
a positive current in the liquid and a negative one in the wire. If
both solutions are normal, or, in general, of equal equivalent concen-
tration, the relative solution pressures of the metals decide the direc-
tion of the resultant current, and its magnitude will be the difference
of the two effects. Thus, the values for the following pairs will be :
Zn-Cd", +0.49- ( + 0.14) = +0.35,
Cd-Cu", +0.14- (-0.61) =+0.75,
Zn-Cu", +0.49- (-0.61)= +1.1,
Zinc the negative pole.
Cadmium the negative pole.
Zinc the negative pole.
The Daniell or gravity cell (Fig. 103) represents the last of these
three combinations. The copper pole is at the bottom, and the zinc
plate is suspended above it.
The cell is charged with a dilute
solution of sodium chloride, and
blue vitriol crystals are thrown
in and dissolve. So long as
the contents are not disturbed,
the solutions require no porous
septum to keep them apart. It
is true that, when the current
is not being used, and the cell
is not working, the cupric sul-
phate diffuses upwards. During
the time that the circuit is
closed, however, the effects of
diffusion are nullified by the
migration of the cuprion away
FlQ- 103< from the zinc and towards the
positive pole. The actual elec-
tromotive force of the current delivered by this cell is a little over
1 volt, and accords, therefore, with the value calculated from the poten-
tial difference observed at each of the two poles.
The cell Zn-H" (0.77 volts) works without a septum, provided the
direct action of the zinc on the acid is minimized by adequate amalga-
mation with mercury. It gives a very inconstant electromotive force,
however, because the platinum plate used as the cathode becomes cov-
ered with bubbles of hydrogen, and so the internal resistance of the cell
ELECTROMOTIVE CHEMISTRY 673
is greatly increased. The polarization (p. 324) also diminishes the
electromotive force. These difficulties are remedied, and, in fact, a
great increase in the E.M.F. of the cell is effected, by surrounding the
cathode with an oxidizing agent which shall convert the hydrogen into
water. The energy obtainable is thus that of a strong oxidizing agent
on zinc, and not merely that of an acid. In the Bunsen cell the cath-
ode is a carbon block surrounded by concentrated nitric acid. In the
dichromate battery it is a carbon block with chromic acid. Each of
these cells gives an E.M.F. of 1.9 volts. In the Leclanche cell the cath-
ode is a mixture of carbon and manganese dioxide, and the fluid is a so-
lution of ammonium chloride from which the zinc displaces hydrogen.
The dioxide, being solid, oxidizes the hydrogen slowly, and the cell can
be used for only a few minutes at a time without becoming polarized.
The E.M.F. is 1.48 volts. Dry cells are of the same nature, but con-
tain a porous solid which holds the liquid by capillary forces (for
Accumulators, see under Lead).
A cell is thus an engine for the direct transformation of chemical
into electrical energy, just as a steam-engine transforms chemical
energy, by several stages, it is true, into mechanical energy. . Accord-
ing to our hypothesis the cell is driven by pressure-differences in the
materials in and around the two poles.
Other Applications : Couples : Concentration Cells. — We are
now in a position to understand the effect of a couple (p. 96). Zinc,
in contact with platinum or copper and immersed in acid, is practically
a short-circuited cell, and it is found that, for some undetermined
reason, hydrogen is liberated more readily from the surface of the
platinum or copper than from that of the zinc. Again, galvanized iron
is also a couple. The zinc is the anode, and, when dilute carbonic acid
rests on a damaged part of the siirface of a sheet of the material, and
is in contact with both metals, the zinc is ionized and passes into
combination as carbonate. The iron is the cathode and is not affected.
On the other hand, a damaged tin-plate (tin on iron) is rapidly rusted.
Here iron is the anode, and goes into combination, while the tin is the
cathode. The ferrous carbonate (q.v.), at first formed, is subsequently
oxidized and gives rust.
When two strips of a metal are immersed together in a solution of a
salt of the same metallic element, — for example, two rods of tin in a
normal solution of stannous chloride, — the pieces of metal show no
difference in potential and no current flows when they are connected
674
INORGANIC CHEMISTRY
FIG. 104.
by a wire. The two poles and their solutions are here alike and
— .07— (—.07) = 0.* But if the solution round one pole is diluted
to iV/10 concentration, the potential at
that pole becomes at once — .07 + .058/2
= — .04 volts, approximately, and a cur-
rent is set up. The positive current
flows from the pole in the stronger solu-
tion through the wire to that in the
dilute solution, and thence through the
liquid. Thus, tin is dissolved from the
latter pole (the,anode), and the concentra-
tion of the solution round it is increased.
On the other hand, tin is deposited in
long needles on the former pole (the cath-
ode), and the concentration in its neigh-
borhood is correspondingly decreased.
The figure (Fig. 104) shows the simplest
arrangement, where the more concen-
trated, denser liquid is below, and one rod
of tin, passing through both layers, fur-
nishes at once the two poles and the con-
nection. The chloridion migrates through the solution, in a direction
opposite to that taken by the tin ions, and thus passes \ipwards into
the dilute solution to balance the fresh tin ions that are continuously
formed. All change ceases when the concentrations have become
equalized. A cell of this kind is called a concentration cell.
The concentration cell is instructive because it shows that the order
of the metals in the electromotive series is not determined by the metal
alone, but also by the concentration of the solution. The order of the
metals in the electromotive series is therefore subject to variation. An
extreme case of this occurred in a recent chapter. Zinc displaces cop-
per from a solution of a cupric (or cuprous) salt, and any but a prodi-
giously dilute solution will show the effect. But a solution containing
cuprocyanion Cu(CN)/ has precisely this very minute concentration of
copper ions which will turn the scale. Hence, zinc will not displace
copper from this solution (p. 624). On the contrary, copper will dis-
place zinc from a solution of a salt of the latter containing excess of
potassium cyanide, and therefore the complex salt K.Zn(CN)3.
* In copper refining (p. 617), the state of equilibrium is destroyed and the
scale turned by the introduction of a current of low E.M.F. from a dynamo.
ELECTROMOTIVE CHEMISTRY 675
The solubilities of insoluble salts, such as those of silver (p. 630),
have also been measured on the principle of the above tin cell. If a
saturated solution of silver chloride is placed round one pole of silver,
and normal argention (say from silver nitrate) round another, a con-
centration cell of high E.M.F. is formed. The magnitude of the
difference in potential enables us at once to calculate the ratio of the
concentrations of argention round each pole, and therefore to get at
the concentration of the silver chloride solution.
Electrolysis : Discharging Potentials. — Another kind of cell
may be made by placing a rod of tin (or any other metal) in an elec-
trolyte on one side of a porous septum (say on the left side of the cell
shown in Fig. 102), and chlorine water on the other. An indifferent
ionogen (say KC1) must be used on the right side also, because chlorine
water is not a good conductor. A platinum plate is likewise required
as a pole at this end. With this arrangement the tin becomes ionized
and renders its pole negative. Simultaneously the chlorine in contact
with the platinum becomes ionic and leaves the other pole positive.
Thus, a current of considerable E.M.F. flows from the platinum pole
to the tin pole through the wire. Tin ions and chlorine ions are
formed, and therefore, potentially, tin chloride. But it will be noted
again that the materials have to be kept apart, otherwise direct union
and mere heat-production will result.
Now it will be recalled (p. 324) that when a salt is electrolyzed,
the materials liberated on the electrodes generate a counter current,
called the polarization current, which it is the business of the electro-
lyzing current to overcome. It will now be clear that the cell just
described is identical with the arrangement which would result from
electrolyzing tin chloride. The amount by which the E.M.F. of the
electrolysing current is cut down (p. 324) by the polarization current
is exactly equal to the potential of a cell like the above containing the
same elements. The minimum difference in potential which will de-
compose an electrolyte is called the discharging potential and is nu-
merically equal to the E.M.F. of the corresponding battery-cell, but of
opposite sign. Furthermore, since the potential of a cell is made up
of two simple potentials, the discharging potential must be made up of
the same two (discharging) potentials. Thus, with the help of a list
of anions and their single potentials, we can calculate at once the total
E.M.F. required to electrolyze any salt. A few of the values are a.^
follows ;
676 INORGANIC CHEMISTRY
POTENTIAL DIFFERENCES FOR ANIONS (IN VOLTS).
I
- 0.80
OH
- 1.96
Br
- 1.27
S04
-2.2
0
- 1.36
HS04
-2.9
Cl
- 1.69
As before, the anode potential is supposed to work against the cathode
potential and is subtracted from it. Thus, the tin-chlorine cell pro-
duces — 0.08- (-1.69) = 1.61 volts, and this E.M.F. will just suffice
to electrolyze tin chloride. Similarly, hydrochloric acid will require
at least — 0.28 -(- 1.69) = 1.41 volts, zinc sulphate 0.49 -(— 2.2)
= 2.69 volts.
Oxygen acids like sulphuric acid show a trace of decomposition at
1.08 volts ( = — 0.28— ( — 1.36) ), and a noticeable but still small decom-
position at 1.68 volts (= -0.28- (-1.96)), due to the H* and 0" and
the H* and OH' respectively. But it is only when the E.M.F. reaches
the values for H* and SO/', and H* and HSO/, namely, 1.92 and
2.62 volts, that rapid electrolysis begins. This observation answers,
incidentally, the question whether in the so-called " electrolysis of
water," when dilute sulphuric acid is used, it is the water or the acid
that is decomposed. The H" and OH' decomposition at 1.68 volts is
very slight, because of the small concentration of the OH', and a lens
is required for its recognition. The more vigorous action resulting
from the discharge of S04" and HSO/ by the use of 2-3 volts is
therefore the one invariably used.
In view of the foregoing facts, it is probably most correct to say
that when dilute sulphuric acid is electrolyzed, e.g. as a lecture
experiment, the oxygen liberated at the anode comes mainly from a
secondary interaction of the discharged material of the anions with
the water (p. 95). A minute proportion of the oxygen in such an
experiment does arise from primary electrolysis of the water, but this
effect of the current is in itself too slight to be visible at a distance.
When, on the other hand, the solution electrolyzed contains a salt of
sodium, and hydrogen is liberated at the cathode, this gas must prob-
ably be regarded as coming chiefly from primary electrolysis of the
water. The discharging potential for sodium chloride should be + 2.54
— ( — 1.69) = 4.23 volts, and with the help of a mercury cathode a
sodium amalgam is easily obtained (p. 554). But it will be found
that, with platinum electrodes, hydrogen and chlorine are liberated
freely from a solution of salt by a current of little more than half the
above mentioned E.M.F. The positive electricity is carried in the
ELECTROMOTIVE CHEMISTRY 677
liquid mainly by the very numerous sodium ions. But, apparently,
when these ions reach the cathode, the potential difference, being
insufficient to discharge the natrion, liberates the hydrion of the water
instead. Thus the accumulating hydroxidion of the water, and the
natrion arriving by migration, together constitute the sodium hydroxide
which is another product of this electrolysis. With higher E.M.F.
and sufficient current density, natrion is doubtless actually discharged,
and in that case a part of the hydrogen liberated is furnished by the
interaction of the metal with the water.
The ordinary chemical behavior of the halogens accords with the
order of their potential differences. Bromine displaces iodine, and
chlorine displaces both (p. 361). Chlorine, however, does not displace
easily perceptible amounts of oxygen from water, because of the small
concentration of the 0" (obtained by secondary ionization of the
OH'). The oxygen freely liberated in sunlight comes from the de-
composition of the HC10 (p. 269). Fluorine, however, which would
probably show a potential difference below that of the much more
plentiful OH', displaces oxygen vigorously by discharging this ion.
In electrolyzing a mixed solution with a moderate current, the
cations and anions with the lowest discharging potentials are first
liberated. Thus, silver (— 1.05) appears before copper (— 0.6).
Hence, in copper refining (p. 617), the copper, of which there is a
continuous supply, is deposited, and the more active metals remain
combined. Indeed, the E.M.F. used is not sufficient in any case to
discharge them.
The Factors of Energy. — We have seen that the amount of a
given supply of electrical energy is described by two factors, the
E.M.F. and the quantity of electricity, and that the weight of material,
which, by its influence, undergoes a given chemical change, is propor-
tional solely to this second factor. On the other hand, the question
whether the supply of energy can initiate the change at all depends on
the magnitude of the first factor alone (p. 324). The total amount of
available energy does not influence the result if the E.M.F. is not above
a certain minimum, which differs from case to case. Now the same is
true of other kinds of energy. The quantity of each may be expressed
as the product of an intensity factor and a capacity factor. The mag-
nitude of the former determines whether the energy can be trans-
ferred or transformed or not. Heat energy, no matter how much of it
is at hand, can neither flow nor be transformed into work unless the
678 INORGANIC CHEMISTRY
source is at a higher temperature than the surroundings. A head of
water will do work only when it is connected with a receptacle at a
lower level. It is the pressure of the water that determines its avail-
ability. The E.M.F. is the corresponding factor of electrical energy.
Now we may presume that chemical energy can be expressed by
two factors. One of these, the capacity factor, must be proportional
to the quantity of material, in other words, to the number of chemical
equivalents. The other is the chemical potential. A chemical change
which does not take place on a small scale will not take place when
more material is used, provided the relative amounts of the interact-
ing substances and the conditions remain unchanged.* We have, in
fact, been assuming all along that this, the capacity factor, is not
the most significant one. But we have devoted ourselves to noting
such things as these : that chlorine will displace bromine, and there-
fore has the higher potential of chemical energy ; that magnesium
reduces sand, while hydrogen does not, and that magnesium is therefore
a more active reducing agent ; and that hypochlorous acid will oxidize
indigo, while free oxygen will not, and is therefore a more powerful oxi-
dizing agent. When we were comparing degrees of activity, therefore,
we were really trying to describe the relative potential of the chemical
energy in all sorts of substances. At present the state of the science
permits this to be done in most cases in a rough fashion only.
Since the capacity factor of chemical energy is proportional to the
number of equivalent weights transformed, and the capacity factor of
electrical energy is proportional to the same thing (Faraday's law), it
follows that the intensity factor of the chemical energy (the chemical
potential) in a given substance undergoing a given change, must be
proportional to the corresponding factor (the E.M.F.) of the electrical
energy produced when the same change takes place in a suitable cell.
The potential differences described above are therefore often much more
significant than are the results of thermochemical measurements, for
the latter attempt to give only the gross quantity of chemical energy
(in terms of the equivalent amount of heat energy), and not the values
of the factors. The potential differences 'come nearer, therefore, to
giving us absolute values for chemical activity than do any other data
we possess.
As we have noted before Q>. 78), in spite of the enormous range
of temperature at our disposal, extending to a point far above 2500° in
* Change in concentration, however, does affect activity, and therefore modi-
fies the chemical potential.
ELECTROMOTIVE CHEMISTRY 679
the electric furnace, there are many substances for whose decomposition
a sufficient potential of heat energy is not available. On the other
hand, amongst substances that are capable of furnishing an electro-
lyte, when dissolved in a suitable solvent or when fused, there are few
that are not decomposable by a current with an E.M.F. of less than 10
volts. Hence even the elements which give the most stable compounds
and are the most difficult to isolate, such as calcium and aluminium,
are liberated by electrical methods with extreme ease.
Methods of Measuring Chemical Activity. — The following is
a summary of the methods of measuring chemical activity.
The thermochemical method (p. 79) can be used in every chemi-
cal change. But the heats of reaction represent the free energy, and
therefore the affinity, only when the heat capacity of the products is
equal to that of the factors and no changes in concentration arise.
For measuring the activity of acids in dilute solution, several
methods have been mentioned : The speed of interaction of different
acids with the same metal (p. 347) ; the acceleration of the speed of
hydrolysis of ethyl acetate (p. 504) and of cane-sugar (p. 500) by
different acids ; the amounts of insoluble salts, such as calcium oxa-
late (p. 598), or zinc sulphide, which, when the system has reached
equilibrium, are found to have been decomposed by different acids under
like conditions ; the relative extents of the hydrolysis of salts of differ-
ent weak acids (p. 344) ; the electrical conductivity (p. 325) and the
freezing- and boiling-points of solutions of acids (pp. 292, 293). These
last measure by physical methods the same thing that the others de-
termine by chemical means, namely, the tendency to ionization on
which the activity of acids depends (p. 347. See also p. 356).
For measuring the activity of bases, we have : The relative speeds
of saponification of esters by different bases (p. 505) ; the relative ex-
tents of the hydrolysis of salts of different weak bases (p. 344) ; the
conductivity and the freezing- and boiling-point methods, which meas-
ure by physical means the tendency to ionization.
For measuring the relative activities of metals and non-metals,
we have : The single potential differences (pp. 670, 676) ; and, for the
former, the speed of interaction of different metals with the same acid
(p. 111).
For measuring the relative activity in non-reversible actions, we
have : The speed with which the actions take place iinder like con-
ditions (p. 250).
680 INORGANIC CHEMISTRY
For measuring the relative activities of the opposed actions in
reversible changes, we have : The concentrations of the materials re-
maining when equilibrium has been reached (p. 254). The relative
activities in different reversible changes may also be ascertained by
comparing the concentrations in one, at equilibrium, with those in
another (cf, p. 257).
For measuring the relative activities 6f oxidizing and reducing
agents, we have : The potential differences in cells arranged after the
manner of the Bunsen and Leclanche cells (p. 673).
If we consider the whole mass of phenomena, it must be admitted
that the scientific study of the quantities of material has reached a far
higher level of exactness, and has very much more nearly enveloped
the whole field covered by the science, than has the study of relative
activity. Yet it is evident that within the past few years substantial
advances have been made in this direction also.
Exercises. — 1. What will be the E.M.F. of each of the following
cells when each of the metallions is present in normal concentration :
Mn — Cu", Cd — Pb"?
2. What will be the E.M.F. of a concentration cell in which the
poles are of lead and the plumbion is one hundred times more concen-
trated round one pole than round the other ?
3. What will be the discharging potential for solutions of the
following substances, if we assume that the concentration of the ions is
normal : manganous chloride, hydrogen bromide ?
4. What weight of aluminium must become ionized every hour in
a cell in order that a current of five amperes strength may be produced ?
What would be the E.M.F. of the current if an acid with normal con-
centration of hydrion surrounded the cathode and a solution of normal
.aluminiou the anode ? How would this E.M.F. be affected if the
uluminion was only one-hundredth normal ?
CHAPTER XXXIX
ALUMINIUM AND THE METALS OF THE EARTHS
THE fourth column of the periodic table (p. 411) contains boron and
aluminium along with a number of rare elements. The chief members
of the family are : boron (B, at. wt. 11), aluminium (Al, at. wt. 27. 1),
gallium (Ga, at. wt. 70), indium (In, at. wt. 115), thallium (Tl, at. wt.
204.1), all on the right side of the column, and scandium (Sc, at. wt.
44.1), yttrium ( Yt, at. wt. 89), lanthanum (La, at. wt. 138.9), samarium
(Sa, at. wt. 150.3), and ytterbium (Yb, at. wt. 173) on the left side.
These elements are all trivalent.
The Rare Elements of this Family. — The oxide and hydroxide
of boron are acidic (p. 527). Those of aluminium (A1(OH)3), gallium
(Ga(OH)8), indium (In(OH)3), and thallium (T10.0H) are basic, but
behave also as acids towards strong bases.
Gallium and indium occur occasionally in zinc-blende, and were
discovered by the use of the spectroscope. The former takes its
name from the country (France) in which the discovery was made, and
the latter from two blue lines shown by its spectrum. Indium gives
a complete series of compounds in which it is trivalent, and the
chlorides InCl and InC^ are also known.
Thallium is found in some specimens of pyrite and blende. It was
discovered by Crookes by means of the spectroscope in the seleniferous
deposit from the flues of a sulphuric acid factory. It received its
name from the prominent green line in its spectrum (Gk. 0aAAds, a
green twig). It gives two complete series of compounds. In those in
which it is trivalent (thallic salts), it resembles aluminium (q.v.\ Thus,
the salts of this series are more or less hydrolyzed by water. Univalent
thallium recalls both sodium and silver. Thallous hydroxide (T10H)
is soluble, and gives a strongly alkaline solution. The chloride is
insoluble in cold water. The solutions of the thallous salts are neutral.
The metal is displaced from its salts by zinc.
Of the elements on the left side of the column, scandium, whose
existence and properties were predicted by Mendelejeff (p. 412), is the
681
682 INORGANIC CHP:MISTRY
best known. The metals of the rare earths, of which it is one, are
found in rare minerals such as euxenite, gadolinite, orthite, and mona-
zite, which occur in Sweden, Greenland, and the United States.
Cerium (Ce, at. wt. 140.25), neodymium (Nd, at. wt. 143.6), arid pras-
eodymium (Pr, at. wt. 140.5), occur along with lanthanum in cerite, a
silicate of these four elements. These four are included amongst the
metals of the rare earths. The compounds of many of these rare ele-
ments behave so much alike that separation is difficult. It is certain,
however, that there are several with atomic weights near to that of
lanthanum for which accommodation cannot easily be found in the
periodic table, and some of these are probably mixtures of still more
closely related elements. Ostwald has compared them to a group of
minor planets such as in the solar system takes the place of one large
planet.
ALUMINIUM.
The Chemical Relations of the Element. — Aluminium is triva-
lent exclusively. Its hydroxide, like that of zinc (p. 648), is feebly
acidic as well as basic, and hence the metal forms two sets of compounds
of the types Nag.AlOg and A12.(S04)8. The salts of both series are more
or less hydrolyzed by water, the former very conspicuously so. It is
worth noting that the hydroxides of the trivalent metals, or metals in
the trivalent condition, such as A1(OH)8, Cr(OH)3, Fe(OH)3, are all
distinctly less basic than are those of the bivalent metals such as
Zn(OH)2, Cd(OH)2, Fe(OH)2, Mn(OH)2. This fact is used in analysis
(cf. also p. 661) in separating the two sets. When a solution of the
chlorides is shaken with precipitated barium carbonate, the free acid
from the more highly hydrolyzed salts of Al'", Cr'" and IV" inter-
acts with this substance, the hydrolysis is promoted :
A1C18 + 3H2O <=> A1(OH)8 + 3HC1,
and eventually the hydroxides A1(OH)3, Cr(OH)s, and Fe(OH)8 are
completely precipitated. The chlorides of the bivalent metals remain
in the solution. Aluminium does not enter into complex anions or
cations, and is too feebly base-forming to give salts like the carbonate
or sulphite.
Occurrence. — Aluminium is found very plentifully in combina-
tion, coming next to oxygen and silicon in this respect. The feldspars
(such as KAlSi308), the micas (such as KAlSiO.,), ami kaolin (clay
ALUMINIUM AND THE METALS OF THE EARTHS 683
H,Al2(Si04)2,H20), are the commonest minerals containing it. Gar-
nets, which are found in metamorphic rocks, are mainly an orthosili-
cate of calcium and aluminium Ca8Al2(Si04)3. Turquoise is a hydrated
phosphate A1.,(OH),PO4,H,O, and cryolite a double fluoride 3NaF,AlF8.
Various forms of the oxide and hydroxide (see below) are also not
uncommon minerals.
Preparation and Physical Properties. — The metal is now
made on a large scale by electrolysis of the oxide (A120S) dissolved in
a bath of molten cryolite. The operation is conducted in cells, the
carbon linings of whicli form the cathodes. The anodes are rods of
carbon which combine with the oxygen a's it is liberated. The metal
sinks to the bottom of the cell and is drawn off periodically, while fresh
portions of the oxide are added from time to time. A current density
of 5 amperes per sq. cm. of cathode area and an E.M.F. of 5-6 volts
maintain the temperature of the molten materials, and cause the
decomposition.
The metal melts at 600-700°, but is not mobile enough to make
castings. It is exceedingly light (sp. gr. 2.6), and in hardness and
tensile strength is the equal of any of the other metals, with the
exception of steel. It has a silvery luster, and does not tarnish, the
firmly adhering film of oxide first formed protecting its surface.
Although, comparing cross-sections, it is not so good a conductor of
electricity as is copper, yet weight for weiyht it conducts better. It is
difficult to work on the lathe or to polish, because it sticks to the tools,
but the alloy with magnesium (6-30 per cent) called magnalium has
admirable qualities in these respects. Aluminium bronze (5-12 per
cent aluminium) is easily fusible, has a magnificent golden luster, and
possesses mechanical and chemical resistance exceeding that of any
other bronze.
The metal and its alloys are used for making cameras, opera-
glasses, cooking utensils, and other articles requiring lightness and
strength. The powdered metal, mixed with oil, is used in making a
silvery paint.
Chemical Properties. — The metal displaces hydrogen from
hydrochloric acid very easily. In sulphuric and nitric acid, however,
it receives a coating of the hydroxide, formed by hydrolysis of the
salt, and the action is slow in the former case, and almost nil in
684 INORGANIC CHEMISTRY
the latter. It displaces hydrogen also from boiling solutions of
the alkalies, forming aluminates :
2A1 + 6NaOH -4 2Na3A103 + 3H2.
In consequence of its very great affinity for oxygen, aluminium dis-
places all the metals, save magnesium, from their oxides. Thus, when
a mixture of aluminium powder and ferric oxide is placed in a crucible
and ignited by means of a piece of burning magnesium ribbon,
aluminium oxide and iron are formed :
Fe203 + 2A1 -> AL,03 + 2Fe.
The very high temperature (about 3000°) produced by the action is
sufficient to melt both the iron (m.-p. 1700°) and the oxide of alumin-
ium. The products, not being miscible, separate into two layers. This
very simple method of making pure specimens of metals like chromium,
uranixim, and manganese, whose oxides are otherwise hard to reduce, is
called by Goldschmidt, the inventor, " alumiuothermy." The sul-
phides, such as pyrite, are reduced with equal vigor by aluminium.
- Aluminium Chloride. — If the metal or the hydroxide is treated
with hydrochloric acid, and the solution is allowed to evaporate, crys-
tals of AlClgjGHgO are formed. When heated, this hydrate is com-
pletely hydrolyzed, hydrochloric acid is given off, and only the oxide
remains. The anhydrous chloride is much used as a catalytic agent
for caiising combination in organic chemistry. It is made by passing
dry chlorine over aluminium, or by heating the oxide with carbon in a
stream of chlorine. Just as in the case of silicon dioxide (p. 520),
neither carbon nor chlorine alone will act upon the oxide.
Aluminium chloride gives a vapor pressure of 760 mm. at 183°, and
sublimes, as a white crystalline solid, without melting. Under pres-
sure, it melts at 193°. In the mode of preparation described above, it
is, therefore, vaporized, and condenses in a cool part of the tube. It
fumes when exposed to moist air on account of the hydrogen chloride
produced by hydrolysis, and only with excess of hydrochloric acid does
it give a clear solution free from basic salts.
Aluminium Hydroxide and the Aluminates. — When an alkali
is added to a solution of a salt of aluminium, the hydroxide A1(OH)S is
precipitated in gelatinous form. It is a white hydrogele (p. 523), and
loses water gradually when dried, without forming any intermediate
hydroxides, until A1208 alone remains. It interacts both with acids
ALUMINIUM AND THE METALS OF THE EARTHS 685
and with bases, and is, therefore, like zinc hydroxide (p. 648), ionized
both as a base and as an acid. It interacts only slightly with ammonium
hydroxide, because this substance is too feebly basic, but, from the solu-
tion in the active alkalies, the aluminates Na3.AlO3, Na.A102, and K.A102,
can be obtained in solid form. Natural forms of this substance are hydrar-
gyllite A1(OH)3(= A1203,3H20), bauxite A1,0(OH)4 (= A1203,2H2O),
which always contains ferric oxide, and diaspore A10.0H (= A1203,
H20).
Commercially, the hydroxide is made by heating bauxite with
sodium carbonate. The ferric oxide, having no tendency to form a
carbonate or to interact with a base, remains unchanged. The sodium
aluminate which is formed can be extracted with water :
A120(OH)4 + Ka2C03 -+ 2NaA102 + C02 + 2H2O.
The hydroxide is then precipitated by passing carbon dioxide through
the solution :
2NaAlO2 + CO3 + 3H2O -> Na2CO8 + 2A1(OH)8.
The aluminates are largely hydrolyzed by water ;
NaA102 + 2H2O <=> NaOH + A1(OH)8.
Hence an excess of sodium hydroxide is required for the complete
solution of aluminium hydroxide by the reversal of this action. Sodium
aluminate is used as a mordant in dyeing (see below), on account of the
ease with which the solution gives up aluminium hydroxide when any
material is present which can combine with the free portion of the
hydroxide and so cause forward displacement of the above equilibrium.
When calcium chloride is added to a solution of sodium aluminate,
the insoluble calcium metaluminate is deposited :
2NaA102 + CaCl2 -> Ca(A102)2 + 2NaCl.
The relations of these various substances are shown by the following
formulae : n
• I W \~s
Al
0 — H 0 — Na
Al-O-H Al-O-Na Al^ Al^ ^ Ca.
686 INORGANIC CHEMISTRY
A number of insoluble metaluminates are found in nature. They
crystallize in the regular system, and are known as spinelles. They
contain bivalent metals in place of the calcium in the last-named
compound. Thus we have spinelle proper Mg(A102)2, and gahnite
Zn(A102)2. Corresponding and isomorphous derivatives of chromic
and ferric hydroxides are chromite Fe(CrO2)2 and magnetite Fe(FeO2),.
Aluminium Oxide* — The oxide (alumina) is found in nature in
pure form as corundum. This mineral is only one degree less hard
than the diamond. Emery is a common variety, contaminated with
ferric oxide, and is widely used as an abrasive. The ruby is pure alu-
minium oxide tinted by a trace of a compound of chromium, while
the sapphire is the same material colored with aluminate of cobalt.
Both can be made artificially by adding a little of the oxide of chro-
mium or of cobalt when aluminium oxide is produced by Goldschmidt's
method (p. 684). The alumina made by gently heating the hydroxide
interacts easily with acids, but after being strongly heated it re-
sembles natural alumina in being very slowly affected by them. Min-
erals containing insoluble compounds of aluminium are attacked when
heated strongly with potassium bisulphate (cf. p. 560), the sulphate of
aluminium being formed.
Aluminium, Sulphate : The Alums. — The sulphate is prepared
by treating either the hydroxide or pure clay (kaolin) with sulphuric
acid. In the latter case the insoluble residue of silicic acid is removed
by nitration :
H2Al2(Si04)2 + 3H,S04 -» A1,(S04)3 + 2H,Si03 + 2H,0.
The salt crystallizes from water as A12(S04)3, 18H20, forming aggre-
gates of leaflets which are very soluble. The solution is acid in
reaction. This compound is used as a mordant under the name
of "concentrated alum." It is employed also in sizing cheaper
grades of paper, an operation required to prevent the absorption and
consequent spreading of the ink. For writing-paper, gelatine solution
is employed. In making printing-papers, rosin soap (made by dissolv-
ing rosin in caustic soda) is mixed with the pulp, and aluminium
sulphate is added. The rosin and aluminium hydroxide are precipi-
tated, perhaps in feeble combination, and pressing between hot rollers
afterwards melts the former and gives a surface to the paper.
When sulphate of potassiiim is added to a strong solution of alu-
ALUMINIUM AND THE METALS OF THE EARTHS 687
minium sulphate, octahedral crystals of alum (potash alum, K2S04,A12
(SO4)3,24H20) are deposited. This is a double salt, and is one of a
large number known as the alums. These have the general formula
M2IS04,M2II1(S04)8, 24H20, and may be made as above by us ing a sulphate
of a univalent metal with one of a trivalent metal. Thus, for M1 we
may use K, NH,, Rb, Cs, and Tl1, and for Mm, Al, Fem, Crni, Mnm,
and Tl111. We may even employ selenates, such as K2Se04. All of
the resulting double salts are isomorphous, and a crystal of one will
continue to grow in a solution of another, acquiring, of course, an outer
layer of different composition but of the same crystallographic orien-
tation.
Potassium- Aluminium Sulphate. — Ordinary alum K2S04,
A12(SO4)3, 24H2O is made from aluminium sulphate obtained from clay
(see above). It is also prepared by heating alunite, a basic alum
found near Rome and in Hungary, and extracting the product with
hot water. The alunite, having the composition KA13(OH)6(SO4)2,
leaves an insoluble residue of the hydroxide, mixed with ferric oxide
which is present as an impurity :
2KA13(OH)6(S04)2 -> K2S04,A12(S04)3 + 4A1(OH)3.
The aqueous solution of alum contains, at 10°, 9 parts of the anhy-
drous salt in 100 parts of water, and at 100° 422 parts in 100 of water.
The hydrated salt melts at 90°. An aqueous solution of this salt or of
sodium phosphate (p. 577) is used for fire-proofing draperies, because the
crystals deposited in the fabric melt easily, and the fused material
protects the fibers from access of oxygen. When heated more strongly
alum loses its water of hydration together with some sulphur trioxide,
and leaves a slightly basic, anhydrous salt known as " burnt alum."
A solution of alum dissolves a considerable amount of aluminium hy-
droxide, giving " neutral alum," a basic salt K2S04, A14(OH)6(S04)S used
as a mordant. The substance is usually prepared by adding sodium
carbonate to the solution of alum as long as the aluminium hydroxide,
formed locally, continues to redissolve.
Aluminium, Sulphide. — This compound is most easily obtained
by mixing pyrite with aluminium powder and igniting with magnesium
ribbon (p. 684) :
4A1 -> 2A12S3 + 3Fe.
688 INORGANIC CHEMISTRY
It forms a grayish-black solid, and is decomposed by water like mag-
nesium sulphide, giving the hydroxide and hydrogen sulphide.
Aluminium Acetate. — This salt is used by dyers, because, being
a salt of a weak base and a weak acid, it is much hydrolyzed by water,
especially at 100°. In mordanting, it thus gives aluminium hydroxide
very easily. It is made by treating lead or barium acetate with
aluminium sulphate, and filtering and crystallizing the solution :
A12(S04)3 + 3Ba(C2H802)2^3BaSo4 + 2A1(C2H3O2)3.
Dyeing: Mordanting. — The problem of the dyer is to confer
the desired color upon a fabric made, usually, of cotton, linen, wool,
or silk, and to do this in such a way that the dye is fast to (i.e., is not
removed or destroyed by) rubbing, and often, also, to washing with
soap. To understand the means by which this is achieved, it must be
noted that cotton and linen consist of hollow fibers of the composition
of cellulose (CgH^Og),,.. Wool is made of hollow fibers, also, and silk of
rods, but the material is entirely different. It contains 17 per cent
of nitrogen in the case of wool, and 20 per cent in the case of silk, and
the nitrogen compounds of which the material is composed are much
more active chemically than is cellulose, and combine incomparably more
easily and firmly with the many kinds of organic compounds which
are used as dyes. Hence, stains on wool and silk are much less often
removable by washing than are those on cotton.
We have space to mention only three kinds of dyes and to describe
in mere outline their use. These are:
1. Insoluble colored bodies which are formed by precipitation
within the fibers and may be applied to any fabric, for their retention
is due to mechanical and not to chemical causes. If cotton is boiled
in a solution of lead acetate (or, better still, sodium plumbite, q.v.),
and is then soaked in boiling potassium chromate solution, it is dyed
a brilliant and permanent yellow. Lead chromate is the colored body :
Pb(C2H802)2 + K2Cr04 <=> 2KC2H302 + PbCrO, |.
In indigo dyeing the fabric is saturated with a solution of indigo-white
(an acid substance) in caustic soda, and is then exposed to the air.
Indigo-blue is formed by oxidation, and, being insoluble, is precipitated
in the fibers :
2CJaHuN1Ot + 02 -» 2C16H10N202J + 2H30.
ALUMINIUM AND THE METALS OF THE EARTHS 689
2. We have direct or substantive dyes, which are withdrawn from
a solution by the goods which are being dyed, and confer upon the latter
a depth of color depending on the strength of the solution and the affinity
of the material for the dye. The union is due in some cases to chemical
combination, and in others to the fact that the dye is more soluble in
the material being dyed than in water, and gives a solid solution in the
former (cf. pp. 146, 235). When the case is simply one of extraction of
the dye, due to the solvent power of the goods, the dye must necessarily be
removed again by washing with sufficient water. Only a small minority
of direct dyes are taken up by cotton or linen in such a way that they
cannot be washed out. Congo red, C^H^NgSgOgNag, is soluble in water,
and is used in dyeing both cotton and wool. The dye is much faster
on the latter than on the former, however.
3. The last class comprises the mordant or adjective dyes. They
work on the principle that the cloth is first impregnated with a substance
capable of attaching itself both to the cloth and, subsequently, to the
dye also, and is then immersed in the dye itself. Substances of this
kind are tannic acid (for basic dyes) and colloidal hydroxides (for acid
dyes) like those of aluminium, tin, iron, and chromium. They are
called mordants (Lat. mordere, to bite). When aluminium hydroxide
is to be used, the cloth is first treated with a hot solution of neutral
alum, aluminium sulphate, aluminium acetate, or sodium aluminate,
and thereby acquires, either by adsorption or feeble combination, a
certain amount of the hydroxide. The fabric is then boiled in water
with the dye. If, for example, alizarine (madder) is used, the cloth is
dyed Turkey red. Alizarine is an orange-yellow, very slightly
soluble acid of the composition C14H804. Since the color is that
of the compound of the dye withjthe mordant, different mordants
give different colors, or shades of color, with the same dye.
This may be illustrated by preparing three solutions, one of
ferric chloride (q.v.) containing ferric hydroxide in solution, one of
aluminium acetate, and one of chromium acetate. When a drop
or two of an alcoholic solution of alizarine is added to each, a precipi-
tate at once appears, which in the first case is violet, in the second
bright-red, and in the third claret-red. These insoluble compounds of
dyes with mordants are identical with the coloring matters produced
in the cloth, and are called lakes (Fr. laque, lac).
Kaolin and Clay : Earthenware and Porcelain* — By the
action of water and carbon dioxide upon granite and other rocks con-
690 INORGANIC CHEMISTRY
taining feldspar KAlSi3O8, the potash is slowly removed, and the com-
pound changed largely into a hydrated orthosilicate H2Al2(SiO4)2,H20.
When this remains in situ, it forms kaolin or china clay, a white,
crumbly material. It usually contains particles of mica and free silica.
When washed away and redeposited, it acquires compounds of iron, and
more or less of the carbonates of calcium and magnesium, becoming com-
mon clay. Ocher, umber, and sienna are clays colored with oxides of
iron and manganese. Fuller's earth is a purer variety.
On account of its plasticity when moist, and its tendency to become
hard, but not to melt, when heated strongly, clay is used in making
bricks, pottery, and porcelain. The presence of calcium and magne-
sium carbonates makes the clay more fusible, that of silica less so. Iron
compounds cause it to turn red during firing. The impure varieties
are formed into bricks and tiles, and are fired at a low temperature.
The efflorescence which often appears on the surface of the bricks
("niter") is generally due to sodium sulphate or sodium chloride
present originally in the clay. For earthenware, glazing must be
applied to make the vessels water-tight. This is often done by throw-
ing salt into the kiln. The hot steam hydrolyzes the salt to sodium
hydroxide and hydrochloric acid, and the former combines with the
clay, giving a fusible silicate which fills the pores of the surface. For
porcelain, very pure clay, free from iron, is employed, and it is mixed
with feldspar and quartz. The feldspar melts and fills the pores so that
a continuous, semi-transparent material results. For china painting,
powdered enamels (p. 607) and metallic oxides which combine with
the clay, giving colored silicates, are used.
Porcelain, if made with sufficient silica, is very infusible. It is
attacked by aqueous and by fused alkalies, however, giving soluble sili-
cates.
Ultramarine was formerly obtained by pulverizing natural lapis
lazuli. Artificial ultramarine of similar composition and more beauti-
ful color is now manufactured. It is made by heating together kaolin,
sodium carbonate, sulphur, and charcoal. The resulting green mass is
then powdered, mixed with sulphur, and heated again until it acquires
the desired shade. The product is used largely in making wall-papers,
water-color paints, and laundry blue. Note-paper is often colored
with it. Its composition is approximately 4NaAlSi04, Na2S2, but the
cause of the brilliant color is not understood. Hydrochloric acid lib-
erates sulphur and hydrogen sulphide and destroys the color of ultra-
marine, decomposing, apparently, the sulphide of sodium.
ALUMINIUM AND THE METALS OF THE EARTHS 691
Analytical Reactions of Aluminium, Compounds- — The
alkalies, and alkaline solutions like that of ammonium sulphide, pre-
cipitate the white hydroxide. The product is soluble in excess of the
active 'alkalies. Soluble carbonates also throw down the hydroxide.
Aluminium compounds, when heated strongly in the flame with cobalt
salts, give a blue aluminate of cobalt.
Exercises. — 1. What are the differences between zinc and alu-
minium, and their corresponding compounds ?
2. Construct equations showing (a) the hydrolysis of aluminium
sulphate (p. 686), (&) the interaction of aluminium sulphate and cobalt
nitrate in the Bunseii flame.
3. Formulate the ionization of aluminium hydroxide (pp. 648, 684).
CHAPTER XL
GERMANIUM, TIN, LEAD
THE elements of the fifth column of the periodic table, aside from
carbon and silicon, are germanium (Ge, at. wt. 72.5), tin (Sn, at. wt.
119), and lead (Pb, at. wt. 206.9). These are on the right side, while
titanium (Ti, at. wt. 48.1), zirconium (Zr, at. wt. 90.6), cerium (Ce,
at. wt. 140.25), and thorium (Th, at. wt. 232.5) occupy the left side.
The Chemical Relations of the Family. — All of these ele-
ments show a maximum valence of four. Germanium, tin, and lead
are also bivalent. In this respect they resemble carbon and differ
from silicon, which is more closely allied to the elements on the left
side of the column. The oxides and hydroxides in which these three
elements are bivalent become more basic, and the elements themselves
more metallic in chemical relations, with increase in atomic weight.
In this they resemble the potassium, calcium, and gallium families.
Curiously enough, the same three hydroxides are also acidic. They
are more strongly acidic than is zinc hydroxide, for the salts they form
by interaction with bases are less hydrolyzed than are the zincates.
This acidic character likewise increases in the order in which the
elements are named above.
GERMANIUM.
Germanium (cf. p. 412) may be described as a transition element
between carbon and tin. It forms two oxides GeO and Ge02 corre-
sponding to those of carbon and of tin. Germanious oxide is not very
definitely basic or acidic, and the sulphide is the only other well-de-
fined compound of this set. Germanic oxide and hydroxide are acidic
entirely. The resemblance to carbon is shown in the formation of an
unstable compound with hydrogen, and of germanium chloroform
GeHCl3. Like carbon, tin, and silicon, germanium gives a volatile
chloride GeCl4 (b.-p. 87°). Like tin and gold (p. 638), it forms
complex sulphides derived from germanic sulphide, such as K,GeS3.
The element was discovered (in 1886) in argyrodite, a complex sul-
phide 4Ag2S,GeSa.
692
GERMANIUM, TIN, LEAD 693
TIN.
The Chemical Relations of the Element. — Tin is both biva-
lent and qiiadrivalent. Each of the oxides and hydroxides SnO and
Sn(OH)2, Sn02 and SnO(OH)2 (or Sn(OH)4), is both basic and acidic,
so that there are really four series of compounds. Still, stannous
hydroxide is mainly a base, of a feeble sort, while stannic hydroxide is
mainly an acid. ^Thus we have stannous chloride, sulphate, and
nitrate, which are stable, although they are all more or less hydrolyzed
by water, and sodium stannite Na2.Sn02 which is unstable. On the
other hand, stannic nitrate, sulphate, and chloride are completely
hydrolyzed by water, while sodium stannate Na2Sn03 is comparatively
stable. The dioxide Sn02 is an infusible solid, and resembles, there-
fore, silicon dioxide. Tin has a tendency to give complex acids and
salts, like H2SnCl6, (NH4)2.SnCl6, H2SnI6, K2SnF6, but these are
ionized also to a small extent after the manner of double salts, giving
ions of Sn"". Tin forms no compounds with hydrogen and no salts
with weak acids, like carbonic acid.
Occurrence and, Extraction. — The chief ore of tin is tin-stone,
or cassiterite Sn02, which consists of tetragonal crystals whose dark
color is due to the presence of iron compounds. The mineral occurs
in Cornwall and the East Indies. The ore is roughly pulverized
and washed, to remove granite or slate with which it is mixed, and
is then roasted, to oxidize the sulphides of iron and copper, and drive
off the arsenic which it contains. After renewed washing to eliminate
sulphate of copper and oxide of iron, it is reduced with coal in a
reverberatory furnace. The tin is afterwards remelted at a gentle
heat, and the pure metal flows away from compounds of iron and
arsenic. In 1900 the production was 4100 tons and 63,700 tons in
England and in the East Indies, respectively. These quantities
together constitute 83 per cent of the total world's output.
Physical and Chemical Properties. — Tin is a silver-white,
crystalline metal of low tenacity but great malleability (tinfoil). Its
specific gravity is 7.3, and its melting-point about 233°. Tin some-
times changes into a gray, pulverulent, specifically lighter modification
(sp. gr. 5.85) when it is kept at a low temperature. The transition
point is 20° (cf. Sulphur, p. 368), and ordinary tin, although it can be
kept almost indefinitely, is therefore really in a metastable condition
(p. 159) below this temperature.
694 INORGANIC CHEMISTRY
Tin-plate (cf. p. 673) is made by dipping carefully cleaned sheets
of mild steel into molten tin. Vessels of copper are also coated,
internally, with tin, to prevent the formation of the basic carbonate
(p. 618). For this purpose they are cleaned with ammonium chlo-
ride, sprinkled with rosin (to reduce the oxide), and heated to 230°.
Molten tin is then spread on the surface with a piece of tow. Com-
mon pins are made of brass wire, and are coated with tin by being
shaken in a solution containing a salt of this metal. The zinc in the
alloy displaces some of the tin, and this is deposited on the surface of
the brass. Alloys of tin, such as bronze (p. 619), soft solder (50 per
cent lead), pewter (25 per cent lead), and britannia metal (10 per
cent antimony and some copper), are much used in the arts.
Tin, although it displaces hydrogen from dilute acids, is not tar-
nished by moist air. With warm hydrochloric acid it gives stannous
chloride SnCLj and hydrogen. Hot, concentrated sulphuric acid forms
stannous sulphate SnS04 and sulphur dioxide. Nitric acid, when cold
and dilute, interacts with it, giving stannous nitrate, and a portion of
the nitric acid is reduced to ammonia (cf. p. 446) :
4Sn + 10HN08 -> 4Sn(N03)2 + 3H20 + NH4N08.
With concentrated nitric acid, stannic nitrate is formed, but most of
this salt is hydrolyzed by the water at the high temperature of the
action (cf, p. 657), and metastannic acid (H2Sn03)5 remains. The final
result is therefore shown by the equation (simplified):
Sn + 4HN08 -> H2Sn03 + 4N02 + H20.
The white, insoluble product continues to give nitric acid during pro-
longed washing, and seems therefore to contain some nitrate, or basic
nitrate. Tin also displaces hydrogen from caustic alkalies, giving a
metastannate, such as K2Sn08.
Stannous Chloride. — This salt is made by the interaction of tin
and hydrochloric acid. Evaporation of the solution gives the colorless
SnCl2, 2H20. When the crystals are heated, or when a strong aqueous
solution is diluted, the salt is partially hydrolyzed. In the latter case
the basic chloride Sn(OH)Cl is deposited. By presence of excess of
hydrochloric acid, the hydrolysis is prevented. The solution is used
as a mordant (p. 688).
GEEMANIUM, TIN, LEAD 695
Stannous chloride tends to pass into stannic chloride SnCl4, and is
therefore an active reducing agent. Thus, it reduces the chlorides of
mercury (p. 655) and of the noble metals, liberating the free metals.
The action is of the form Hg" + Sn" — >Hg + Sn"". Stannous
chloride reduces cupric and ferric chlorides to the cuprous and ferrous
conditions in like manner :
2FeCl3 + SnCLj -> 2FeCL, + SnCl4 or 2Fe*" + Sn" -> 2Fe" + Sn"".
It also reduces free oxygen, or, what is the same thing, is oxidized by
the air. In this case, stannic chloride is formed in the acid solution
and the liquid remains clear ; in the neutral solution a precipitate of
the basic chloride is formed as well :
GSnCL, + 2H20 + 02 -> 4Sn(OH)Cl + 2SnCl4.
Powdered tin, if placed in the bottle along with the acid solution, will
undo the effects of this action by reducing the stannic salt to the
stannous condition once more.
Stannic Chloride. — When chlorine acts upon tin, or upon stan-
nous chloride (either solid or dissolved), stannic chloride is formed.
The compound is a colorless liquid (b.-p. 114°) which fumes very
strongly in moist air, giving hydrochloric acid and stannic acid. It
was formerly known, after its discoverer (1605), as spiritus fumans
Libavii. The aqueous solution, when freshly made, has almost no
conductivity, and the compound is therefore very slightly ionized.
As hydrolysis proceeds, the conductivity increases, but the hydro-
chloric acid is the conducting substance. After a time hydrolysis
becomes almost complete. The stannic acid which is formed is not
precipitated, however, but remains dissolved in colloidal (p. 523)
form :
SnCl4 + 4H20 <z> 4HC1 + Sn(OH)4.
The chloride, with small amounts of water, gives crystalline hydrates
SnCl4,3H20, SnCl4,5H20, and SnCl4,8H2O, of which the second is used
as a mordant under the name " oxymuriate " of tin. The double (or per-
haps complex) salts SnCl4,2MICl are easily made. Ammonium-stan-
nic chloride or " pink-salt " (NH4)2SnCl6, which is an example of this
class of compounds, is used as a mordant on cotton. It gives a red
lake with alizarine (p. 689).
Stannic bromide SnBr4 melts at 30°, boils at 201°, and is soluble in
water.
696 INORGANIC CHEMISTRY
a-Stannic Acid and its Salts. — When a solution of stannic
chloride is treated with ammonium hydroxide, a white, gelatinous pre-
cipitate is formed. To this the formula H2Sn03 is generally assigned :
SnCl4 + 4NH4OH -> 4NH4C1 -I- H2Sn03 + H20.
It is, however, in reality, a hydrogele, and loses water gradually until
the dioxide remains. Thus, neither Sn(OH)4 nor SnO(OH)2 is obtain-
able as a definite compound. When stannic oxide is fused with
caustic soda, the metastannate, or a-stannate, is formed :
Sn02 + 2NaOH -* Na2Sn03 + H20.
This compound is obtainable as Na^nO,,, 3H20, and is used as a
mordant under the name of " preparing salt." When its solution is
acidified, the above mentioned a-stannic acid is formed by double decom-
position. This a-stannic acid interacts readily with acids and alkalies,
and the chloride obtained from it is identical with stannic chloride
described above.
The a-stannates of the metals, aside from those of potassium and
sodium, like the silicates and carbonates which they much resemble,
are -all insoluble in water, and may be made by double decomposition.
^-Stannic Acid., or Metastannic Acid. — The product of the
action of nitric acid upon tin is a hydrated stannic oxide like the fore-
going substance, but is not identical with it. It is not easily soluble
in alkalies. By boiling it with caustic soda, .however, and then ex-
tracting with pure water, a soluble sodium /?-stannate, K^S^On, is
obtained. /3-stannic acid is also very slowly attacked by acids, and
the chloride secured from it is not identical with the ordinary chloride.
For these reasons it is supposed to be a hydrate of a polymer of
stannic oxide (Sn02)5. When fused with caustic soda, it gives the
same a-stannate as does the dioxide itself.
The difference between the properties of the two stannic acids
was noticed by Berzelius (1811), and was the first case in which iden-
tity in composition was found not to be accompanied by identity in
properties (cf. Isomers, p. 488).
The Oxides of Tin. — When stannous oxalate is heated in absence
of air, stannous oxide remains : SnC204 — > SnO -f- C02 + CO. It is a
black powder which burns in the air, giving the dioxide. The corre-
sponding hydroxide Sn20(OH)2 is formed by adding sodium carbonate
GERMANIUM, TIN, LEAD 697
to stannous chloride solution. It is a white powder, easily dehydrated,
and interacts with alkalies to give a soluble stannite, such as N"a2Sn02.
When the solution is boiled, tin is deposited, and sodium stanuate is
formed, the behavior resembling that of cuprous oxide when heated
with acids (p. 622). With acids, the hydroxide gives stannous salts.
Stannic oxide is found in nature (p. 693), and may be made in pure
form by igniting /3-stannic acid. When heated, it becomes yellow, but
recovers its whiteness when cooled (cf. Zinc oxide, p. 648). Prepared
at a low temperature, it interacts easily with acids, but after strong
ignition, is affected by them very slowly.
The Sulphides of Tin. — Stannous sulphide is obtained as a dark-
brown precipitate when hydrogen sulphide is led into a solution of a
stannous salt.
Stannic sulphide is formed likewise by precipitation, and is yellow
in color. It is made also by heating together tin filings, mercury, sul-
phur, and ammonium chloride. The mercury and ammonium chloride
are ultimately volatilized, and the stannic sulphide remains in the form
of yellow, crystalline scales ("mosaic gold" or "bronze powder").
Stannic sulphide loses sulphur when strongly heated, and leaves stan-
nous sulphide. It is not much affected by dilute acids, but interacts
with solutions of ammonium sulphide or sodium sulphide, giving a
soluble complex sulphide, the sulphostannate :
SnS2 + (NH4)2S -» (NH4)2.SnS3.
The corresponding sodium salt is easily crystallized in the form
Na^SnSg, 2H20. Stannous sulphide is not affected by sulphides, but
poly sulphides, such as yellow ammonium sulphide, give with it the
above mentioned sulphostannates :
SnS + (NH4)2S + S -> (NH4)2.SnS3.
With acids the sulphostannates undergo double decomposition, but
the free acid H2.SnS3 thus produced is unstable and breaks iip, giving
off hydrogen sulphide, and depositing stannic sulphide.*
Analytical -Reactions of Salts of Tin. — The two ionic forms
of tin, Sn", and Sn'*", are both colorless. Their behavior is different.
* These and similar compounds are often called thiostannates, orthothian-
timonates, etc. The prefix sulpho- gives more euphonious words, however, and
is used here for all excepting the thiocyanates.
698 INORGANIC CHEMISTRY
They give a brown and a yellow sulphide, respectively, with hydrogen
sulphide. The solubility of these sulphides in yellow ammonium
sulphide distinguishes them (cf. p. 661) from those of cadmium, copper,
and other metals whose sulphides are similarly insoluble in dilute acids.
The sulphides of arsenic, antimony, and gold (q.v. ), however, behave
like those of tin in this respect. The reducing power of distannion
Sn" is very characteristic (p. 695). Zinc displaces tin from solutions
of its salts. The oxides are reduced by charcoal in the reducing part
of the Bunsen flame and the metal is liberated.
LEAD.
The Chemical Relations of the Element. — Lead is both biva-
lent and quadrivalent. The oxides PbO and Pb02, and the correspond-
ing hydrated oxides, are both basic and acidic. Lead monoxide is a
fairly active base, comparable with cupric oxide, and lead dioxide a
feeble one. Both are feebly acidic. The salts of bivalent lead, like
Pb(N03)2, commonly called the plumbic salts, are somewhat hydrolyzed
by water, but less so than are those of tin. The tetrachloride and
other salts of quadrivalent lead are completely hydrolyzed. The
plumbites Nag.PbOg and plumbates Na2.Pb03 are hydrolyzed to a con-
siderable extent. All the compounds in which lead is quadrivalent
give up half of the negative radical readily, and are reduced to the
" plumbic " condition. The metal displaces hydrogen with difficulty,
and is easily displaced by zinc. Lead compounds are all poisonous,
and the effects of repeated, very minute doses- are cumulative — re-
sulting in " lead colic."
Occurrence and Metallurgy. — Commercial lead is almost all
obtained from galena PbS, which crystallizes in cubes. This ore
often contains considerable amounts of silver sulphide Ag2S, which
is isomorphous with it, and it occurs in association with sulphides of
arsenic, antimony, zinc, copper, and iron. Other salts of lead are of
less common occurrence.
The sulphide of lead is first roasted until a sufficient proportion
of it has been converted into the oxide and sulphate. The furnace-
doors are then closed, and the temperature raised in order that these
products may interact with the unchanged part of the sulphide :
PbS-f 2PbO
PbS + PbS04 -> 2Pb + 2S02
GERMANIUM, TIN, LEAD 699
Another plan consists in heating galenite with scrap iron or iron ores
and coal :
PbS + Fe -> Pb + FeS.
The molten ferrous sulphide rises to the top as a matte.
The purification of the lead from the other metals whose sulphides
have been reduced at the same time is often troublesome. In Parke's
process (p. 627) for the extraction of the silver by means of zinc, the
greater part of the foreign metals, with the exception of bismuth,
passes into the zinc scum. About 0.5 per cent of zinc remains in the
lead, and is oxidized by the action of a jet of steam before the lead is
poured into the molds. In one establishment, at Trail, near Rossland,
B.C., the refining is carried out electrolytically. The lead is cast into
plates, and the process is similar to that used for refining copper
(p. 617). The bath is a solution of hydrofluosilicic acid containing
hydrofluosilicate of lead. The less electro-positive metals, Cu, Sb, Bi,
As, Ag, Au, with 10-16 per cent of lead, remain as a sort of skeleton
of the anode, while Zn, Co, Ni, and Fe go into solution and are not
redeposited.
Physical and Chemical Properties. — Metallic lead is gray in
color, very soft, and of small tensile strength. Its specific gravity is
11.4, and its melting-point 326°. While warm, it is formed by hydraulic
pressure into pipes which are used in plumbing and for covering elec-
tric cables. On account of its very slow interaction with most sub-
stances, sheet lead is used in chemical factories, for example, to line
sulphuric-acid chambers. An alloy containing 0.5 per cent of arsenic
is used in making small shot and shrapnel bullets. Type-metal con-
tains 20-25 per cent of antimony (q.u.). In both cases greater hard-
ness (cf. p. 532) is secured by the addition of the foreign metal.
Lead oxidizes very superficially in the air. The suboxide Pb20 is
supposed to be first formed. The final covering is a basic carbonate.
Contact with hard waters confers upon lead a similar coating composed
of the carbonate and the sulphate. These deposits, being insoluble,
inclose the metal and protect the water from contamination with
lead compounds. Pure rain-water, however, since it has no hard-
ness, but contains oxygen in solution, gives the hydroxide Pb(OH)2,
which is noticeably soluble. When heated in the air, lead gives
the monoxide PbO or minium Pb304, according to the temperature
(see below).
700 INORGANIC CHEMISTRY
The metal displaces hydrogen from hydrochloric acid very slowly.
It is hardly affected by concentrated sulphuric acid, although, when
the commercial acid is diluted with water, a slight precipitate of lead
sulphate, acquired from the evaporating-pans, is thrown down. Nitric
acid attacks it readily, giving lead nitrate and oxides of nitrogen (p.
446).
Chlorides and Iodide of Lead. — Plumbic chloride is precipi-
tated when a soluble chloride is added to a solution of a soluble lead salt.
It is slightly soluble in water (1.5 : 100) at 18°, and considerably more
so at 100°. In the saturated solution at 25° about 50 per cent of the
lead is in the form Pb**, 44 per cent as PbCl", and 6 per cent as PbCl2
(cf. p. 346).
Lead tetrachloride is a solid at —15°, and loses chlorine at the
ordinary temperature. It is made by passing chlorine into plumbic
chloride suspended in hydrochloric acid. The solution appears to con-
tain H2PbClg. With ammonium chloride this solution deposits crystals
of a double or complex salt PbCl4,2NH4Cl analogous to pink-salt
(p. 695). When this is thrown into cold, concentrated sulphuric acid,
an oil of the composition PbCl4 settles to the bottom. The oil fumes
in the air, and, in general, closely resembles stannic chloride SnCl4.
When dissolved in little water, it slowly deposits PbCl2 and gives off
chlorine. With much water it is quickly hydrolyzed, and lead dioxide
is thrown down :
PbCl4 + 2H20 -> Pb02 + 4HC1.
Lead iodide PbI2 is yellow in color, and is formed by precipitation.
It is somewhat soluble in boiling water, and crystallizes in yellow
scales from the hot solution.
Plumbic chloride and iodide are both, under some conditions, more
soluble in acids or salts with a common negative ion than they are in
water, and form soluble, but somewhat unstable, complex salts (cf.
p. 621).
Oxides and Hydroxides. — There are five different oxides of
lead, Pb20, PbO, Pb304, Pb203, and Pb02. The suboxide Pb20 is a
dark-gray powder, formed by gently heating the oxalate. Plumbic
oxide, or lead monoxide PbO, is made by cupellation (p. 627) of lead,
and the solidified, crystalline mass of yellowish-red color is sold as
" litharge." The yellow, powdery form is called " massicot/' and may
GERMANIUM, TIN, LEAD 701
be obtained by heating the nitrate or carbonate. All the other oxides
yield this one when they are heated above 600° in the air. Plumbic
oxide takes up carbon dioxide from the air, and therefore usually con-
tains a basic carbonate. It dissolves in warm sodium hydroxide solu-
tion, giving a plumbite Na^.PbOgj a saturated solution redeposits
part of the oxide in crystalline form when it cools. The oxide is used
in glass-making and for preparing salts of lead.
Plumbic hydroxide is formed by precipitation. It gives up water
in three stages with different aqueous tensions (cf. p. 122), the prod-
ucts in the order of decreasing tension being Pb(OH)2, Pb20(OH)2,
Pbs02(OH)2. These substances, as will be seen, are equivalent in
composition to PbO, H20, 2PbO, H20, and 3PbO, H20 respectively. The
hydroxide is observably soluble in water, and gives a solution with a
faintly alkaline reaction. With acids it forms salts of lead. It inter-
acts also with potassium and sodium hydroxides to form the soluble
plumbites, like Na.j.Pb02.
Minium, or red lead, Pb804, gives off oxygen when heated :
2Pb304 «=± 6PbO + 02.
The dissociation pressure varies with the temperature :
Temperature 445° 500° 555° 636°
Pressure in mm 5 60 183 763
Since the partial pressure of oxygen in the air is 150 mm., the sub-
stance decomposes at about 550°. It can be formed in air by reversal
of the action represented above, but only below this temperature
(cf. p. 591). In pure oxygen of one atmosphere pressure it could be
formed at 600°, but not at 650°. On account of unequal heating
during manufacture, commercial red lead is never fully oxidized, and
always contains litharge. Conversely, commercial litharge usually
contains a little minium.
Minium, when heated with warm, dilute nitric acid, is decomposed,
and leaves lead dioxide as an insoluble powder. It is therefore re-
garded as lead orthoplumbate (see below) :
Pb2.Pb04 + 4HNO3 <=± 2Pb(N03)2 + H4PbOv
The double decomposition as a salt that it thus undergoes is followed
by dehydration of the plumbic acid, which is unstable (H4PbO4 — »
PbOa + 2H20), and the dioxide remains. Ked lead is used in glass-
702 INORGANIC CHEMISTRY
making, and, when mixed with oil, gives a red paint which is specially
applicable to iron- work (cf. p. 658).
Lead dioxide may be obtained as described above in the form of a
brown powder. Unlike most oxides, it is a conductor of electricity.
It is usually made by adding bleaching powder to an alkaline solution
of plumbic hydroxide :
Na^PbO, + Ca(OCl)Cl + H2O -» 2NaOH + CaCL, + Pb02 J .
In this action we may regard the free lead hydroxide, formed by
hydrolysis of the plumbite, as being oxidized by the bleaching powder.
This dioxide is an active oxidizing agent. It interacts with, and
sets fire to, a stream of hydrogen sulphide, and it liberates chlorine from
hydrochloric acid. With acids it gives no hydrogen peroxide, and is
not a peroxide in the restricted sense of the term (p. 308). Lead
dioxide interacts with potassium and sodium hydroxides, giving soluble
plumbates. These are derived from metaplumbic acid. The potas-
sium salt K2Pb03,3H20 is analogous to the metastannate K2Sn08,3H20
(p. 696). A mixture of calcium carbonate and lead monoxide absorbs
oxygen when heated in a stream of air, and the yellowish-red calcium
orthoplumbate is formed :
4CaC03 + 2PbO + 02 <=> 2Ca2Pb04 + 4CO2.
The action is reversible, and is at the basis of Kassner's method of
manufacturing oxygen from the air.
The Storage Battery. — In the storage battery the plates con-
sist of leaden gratings, the openings of which are filled with the active
materials, and the fluid is dilute sulphuric acid. When the battery
discharges, the S04" ions migrate towards those plates (usually the
outer ones) which are filled with finely divided lead, and convert this
into a mass of (insoluble) lead sulphate : SO/' + Pb -» PbS04 -4- 2 © .
These plates receive therefore negative charges. Simultaneously the
H" ions pass towards the other plates, and reduce to monoxide the
lead dioxide with which they are filled :
2H' + Pb02 -* ILp + PbO + 2 ©.
These plates acquire positive charges, and, by secondary interaction of
the monoxide with the sulphuric acid, become filled, like the negative
GERMANIUM, TIN, LEAD 703
plates, with lead sulphate. During the discharge the fluid thus loses
much sulphuric acid, and acquires a lower specific gravity, a fact by
means of which the approach of complete discharge may be ascer-
tained. The E.M.F. of the current is about 2 volts.
When a current, opposite in direction to the one which it yields, is
led into the exhausted cell, the negative terminal of the dynamo cir-
cuit being connected with the negative pole of the battery and the
positive with the positive, a new set of changes occurs. The H* ions
of the bath are attracted to the plate which has the negative charge, and
an equivalent number of SO/' ions are formed, so that only metallic
lead remains :
PbS04 + 2H' + 2 0 -> Pb + 2H' + SO/'.
To express this otherwise, the hydrogen here reduces the lead sulphate
(which is white) to metallic lead (a black powder). Simultaneously
the SO/' is attracted to the positively charged plate, and forms lead
persulphate with the lead sulphate there present : SO/' -f- PbSO4 + 2
© — » Pb(S04)2. The lead persulphate, however, in a battery which is
working normally, is at once hydrolyzed, and the filling of the plate is
changed into lead dioxide: Pb(S04)2 + 2H20 -»Pb02 + 2H2S04.
Both plates are thus brought back to the condition in which they were
before the discharge.
The set of changes last described, that involved in the operation of
charging, consumes energy, while the changes connected with the dis-
charging, liberate energy. The whole may be put into a single equa-
tion :
Charge — »
2PbS04 + 2H20 «=» Pb + 2H2S04 + Pb02.
< — Discharge
Lead Nitrate. — This salt may be made by treating lead, lead
monoxide, or lead carbonate with nitric acid. It forms white, anhy-
drous octahedra. The nitrate and acetate (see below) are the salts of
lead which, because of their solubility, are most commonly used. The
solubility of the nitrate is, 48 parts in 100 at 10°, and 153 parts in
100 at 100°. Since the solubility increases with rise in temperature,
we should expect the process of solution to be accompanied by absorp-
tion of heat (p. 260), and this is found to be the case. On account of
hydrolysis, the solution is acid in reaction.
704 INORGANIC CHEMISTRY
Lead Carbonate. — This compound is found in nature in rhombic
crystals isomorphous with those of aragonite. It may be formed as a
precipitate by adding a soluble bicarbqnate to lead nitrate solution.
With normal sodium carbonate, a basic carbonate Pb3(OH)2(C02)2 is de-
posited. This basic salt is identical with white lead, which, on account
of its superior opacity, has better covering power than zinc-white
(p. 648) or permanent white (p. 610). The substance is manufactured
in various ways, all of which involve the oxidation of the lead by the
air, the formation of a basic acetate by the interaction of vinegar or
acetic acid with the oxide, and the subsequent decomposition of the salt
by carbon dioxide. The best quality is obtained by the Dutch method.
In this, gratings of cast lead are placed above a shallow layer of vinegar
in small pots. These pots are buried in manure, which by its decom-
position furnishes the carbon dioxide and the necessary warmth. The
gratings are gradually converted into a white mass of the basic carbonate.
The vapor of acetic acid arising from the vinegar may be regarded as a
catalytic agent (cf. p. 508), since it is used over and over again.
Lead Acetate. — This salt is made by the action of acetic acid on
litharge, and crystallizes in prisms of the composition Pb(C2H302)2,
3H2O. It is easily soluble in water, and, from the sweet taste of the
solution, is named sugar of lead (used in medicine). The basic salt
Pb(OH)(C2H302) is formed by boiling a solution of lead acetate with
excess of litharge. Unlike most basic salts, this basic salt is soluble in
water, and its solution has a faintly alkaline reaction.
Lead Sulphate. — The sulphate occurs in nature as anglesite, and
is isomorphous with heavy spar. Being insoluble in water, it is easily
obtained by precipitation. It is slightly soluble in concentrated sul-
phuric acid (p. 387). It is dissolved to a noticeable extent by nitric
acid, since this acid is more active than sulphuric acid (cf. p. 601).
It is also soluble in concentrated sodium hydroxide solution, on ac-
count of the removal of the Pb" ions which are a factor in its solubility
product and their passage into the Pb02" anion of sodium plumbite
(cf. p. 621). Finally, it dissolves easily in ammonium tartrate, since
lead enters into the complex anion of the tartrates in the same way
that copper does (cf. p. 623). Barium sulphate, which is of the same
order of insolubility as lead sulphate, is somewhat affected by nitric
acid, but not by sodium hydroxide or by tartrates. The element
barium lacks both the characteristics which lead here exhibits.
GERMANIUM, TIN, LEAD 705
Lead Suljthide. — Natural lead sulphide (galena) is black, and its
crystals have a silvery luster. The precipitated salt is black and
amorphous. It is more easily attacked by active acids than is mercuric
sulphide (cf. p. 651 ). Concentrated nitric acid, being an oxidizing agent
as well as an acid, attacks and dissolves it readily.
Analytical Reactions of Lead Compounds. — Hydrogen sul-
phide precipitates the black sulphide, even when dilute acids are
present. Sulphuric acid throws down the sulphate. Potassium hy-
droxide gives the white hydroxide, which dissolves in excess to form
the plumbite. Potassium chromate or dichromate (q.v.} gives a yellow
precipitate of lead chromate PbCr04, which is used as a pigment under
the name of " chrome-yellow."
TITANIUM, ZIRCONIUM, CERIUM, THORIUM.
The metals on the left side of the fifth column of the periodic table
are all quadrivalent, although compounds in which a lower valence ap-
pears are numerous in this family. The first two are feebly base-form-
ing as well as feebly acid-forming ; the last two are base-forming ex-
clusively.
Titanium occurs in rutile TiTi04. Derived from it are a number of
titanates of the form K2Ti03, titanic iron ore (menaccanite) being fer-
rous titanate FeTi08.
Zirconium is found in zircon, the orthosilicate of zirconium
ZrSi04, which occurs in square prismatic crystals isomorphous with
rutile, cassiterite (SnSn04), pyrolusite (MnMn04), and thorite
(ThSi04). The oxide is used in making the incandescent substance in
some forms of gas lamps.
Cerium occurs chiefly in cerite [Ce, La, Nd, Pd]Si04.H20.
Thorium is found in thorite ThSi04 but most of the supply comes
from monazite sand. The nitrate Th(N08)4,6H20 is used in making
Welsbach incandescent mantles (cf. Flame, p. 510). The mantle of
knitted cotton is dipped in a solution of this salt along with one per
cent of cerium nitrate Ce(ISr03)4, and is then ignited. The oxides Th02
(thoria) and Ce02 (ceria) which remain form a fairly coherent mass.
Larger or smaller proportions of cerium oxide diminish the luminosity of
the glowing material. It is supposed that the particles of finely divided
cerium oxide assist the union of the oxygen and illuminating-gas cata-
lytically and hence are themselves raised to a high temperature by the
706 INORGANIC CHEMISTRY
action which is thus concentrated in their neighborhood. The thoria
acts simply as a* nonconducting support for the ceria.
The Nernst lamp is an incandescent electric lighting arrangement in
which a rod of the oxides of several of the rare metals takes the place of
the common carbon filament. The peculiarity of this lamp is that pre-
heating is required before the rod attains a temperature at which it
will conduct the current. When this point has been once reached, the
resistance enables the current to maintain the rod at the temperature
of incandescence. For equal consumptions of electricity, this form of
electric lamp gives a greater yield of light than does the ordinary,
carbon-filament, incandescent bulb.
Exercises. — 1. IH what order should you place the elements dealt
with in this chapter, beginning with the least metallic, and ending
with the most metallic (p. 533) ?
2. Construct equations showing («) the interaction of tin and con-
centrated sulphuric acid, (£>) of water and stannous chloride, (c) of
chlorauric acid and stannous chloride (p. 695), (d} of oxygen and stan-
nous chloride in acid solution, (e) the decomposition of lead oxalate
(p. 700), (/) the interaction of lead monoxide and acetic acid, (<?) and
of lead monoxide and lead acetate.
3. To which class of ionic actions do the reductions by stannous
chloride belong ?
4. What interactions probably occur when lead dioxide liberates
chlorine from hydrochloric acid ?
5. How should you set about preparing (a) lead oxalate (insol-
uble), (6) lead chlorate (soluble) ?
6. Should the formula of the sulphate of quadrivalent lead be
written Pb(S04)2 or Pt>S208, and is it related to persulphuric acid
(H2S208) ?
7. Describe in terms of the categories used in connection with the
phase rule (p. 592) the system furnished by minium at 500°.
ARSENIC, ANTIMONY, BISMUTH
THIS family is very closely related to the elements phosphorus and
nitrogen which precede it in the same column of the periodic table.
In reading this chapter, therefore, constant reference should be made
to the chemistry of the corresponding compounds of phosphorus. A
general comparison of the elements arsenic (As, at. wt. 75), antimony
(Sb, at. wt. 120.2) and bismuth (Bi, at. wt. 208.5) with each other and
with the two already disposed of will be given at the end of this chap-
ter. It is sufficient here to say that arsenic is mainly an acid-forming
element, and is therefore a non-metal, while antimony is both acid-
forming and base-forming, and bismuth is base-forming. Each of the
three elements gives sets of compounds in which it is trivalent, and
others in which it is quinquivalent. None of the free elements dis-
places hydrogen from dilute acids.
ARSENIC.
The Chemical Relations of the Element. — Arsenic forms a
compound with hydrogen AsH3. It gives several halogen derivatives of
the type AsX3 which are completely hydrolyzed by water. Its oxides
and hydroxides are acidic. Salts derived from H3As03 (= As208,3H20)
and HAs02 (= As203,H20), and named orthoarsenites and metarsenites,
are known, as are also salts derived from various arsenic acids :
H3As04 = As205,3H20 = AsO(OH)3,
H4As207 = As205,2H20 = As203(OH)4,
HAs03 = As206,H20 =AsO(OH).
In analogy to the phosphates (p. 465), these are named ortho-, pyro-
and metarsenates respectively. The last column shows a method of
writing the formulae which is often adopted (cf. p. 470). Sulphates,
nitrates, carbonates, and other salts of arsenic are not formed. There
are, however, complex sulphides of the forms Na3.AsS8 (sodium ortho-
tmlpharsenite), ISTag.AsS.j (sodium orthosulpharsenate), etc.
In many natural sulphides, such as pyrite FeS2 and zinc-blende
707
708 INORGANIC CHEMISTRY
ZnS, a part of the sulphur is replaced by arsenic, which must there-
fore be playing the part of a bivalent element in such cases. When
much arsenic is present, the formulae are written thus : Fe[S,As]2 and
Zn[S,As].
Occurrence and Preparation. — Arsenic is found free in
nature. It occurs also in combination with many metals, particularly
in arsenical pyrites FeAsS. Two sulphides of arsenic, orpiment
As2S8 and realgar As2S2, and an oxide, white arsenic As208, are less
common.
The element is obtained either from the native material or by heating
arsenical pyrites : FeAsS — » FeS + As. During the roasting of the
sulphur ores of metals, arsenic trioxide is formed by the oxidation of
the arsenic so frequently present, and collects as a dust in the flues.
Physical Properties. — In its ordinary condition, the free element
is steel-gray in color, metallic in appearance, and crystalline in form.
When the vapor is suddenly cooled, however, a yellow variety is
obtained, which is soluble in carbon disulphide, is phosphorescent in
the air, and in other ways resembles common phosphorus. This, like
yellow phosphorus, is the less stable form.
Elementary arsenic is easily volatilized at 180°, and acquires a vapor
pressure of 760 mm. long before the melting-point (480°, under high
pressure) is reached. The density of the vapor measured at 644°
gives 308.4 as the weight of the G.M.V. (22.4 liters at 0° and 760
mm.). The weight combining with one chemical unit weight (35.45 g.)
of chlorine, is 75 g., and this is the amount of the element found in the
G.M.V. of the vapor of the chloride. It is also the smallest weight
found in the G.M.V. of any volatile compound of arsenic, and is there-
fore accepted as the atomic weight. Since 308.4 is equal approxi-
mately to 4 x 75 (= 300), the formula of the vapor of the simple
substance at 644° is As4. At 1700° dissociation has occurred, and the
formula is As2 (cf. p. 206).
Chemical Properties. — The free element burns in the air, pro-
ducing clouds of the solid trioxide As208. It unites directly with the
halogens, with sulphur, and with many of the metals. When boiled
with nitric acid, chlorine water, and other powerful oxidizing agents,
it is oxidized in the same way as is phosphorus, and yields arsenic
acid H3As04.
ARSENIC, ANTIMONY, BISMUTH 709
Arsine. — This substance corresponds in composition to ammonia
and phosphine, and some of the ways in which it may be formed are
analogous to those used in the case of these substances. Thus, when
arsenic and zinc are melted together in the proportions to form zinc
arsenide Zn8As2, and the product is treated with dilute hydrochloric
acid, the result is similar to the action of water or dilute acids upon
calcium phosphide. Arsine in fairly pure condition is evolved as
a gas:
Zn3As2 + 6HC1 -» 2AsH3
Arsine is formed also by the action of nascent hydrogen (cf. p. 423)
upon soluble compounds of arsenic, such as arsenious chloride AsCl8
or arsenic acid. When a solution of one of these substances is added
to zinc and hydrochloric acid in a generating flask, the disagreeable
odor of the hydrogen evolved shows the presence of the arsine :
AsCl8 + 6H -> AsH3 + 3HC1.
This method, naturally, does not furnish pure arsine, for free hydrogen
predominates in the gas. Pure arsine may be secured by leading the
mixture with hydrogen through a U-tube immersed in liquid air. The
arsine (b.-p. — 40°) condenses as a colorless liquid.
Arsine burns with a bluish flame, producing water and clouds of
arsenic trioxide : 2AsH8 -f 302 — » 3H20 + As208. The combustion of
hydrogen containing arsine produces the same substances. Since
arsine, when heated, is readily dissociated into its constituents
(cf. p. 252), the vapor of free arsenic is present in the interior of the
hydrogen flame. This arsenic may be condensed in the form of a
metallic-looking, brownish stain by interposition of a cold vessel of white
porcelain. Even when only a trace of the compound of arsenic has
been added to the materials in the generator, the stain which is pro-
duced is very conspicuous. This behavior thus furnishes us with an
exceedingly delicate test — Marsh's test — for the presence of arsenic
in any soluble form of combination. The compounds of antimony
alone show a similar phenomenon (see Stibine). In carrying out the
test, a tube of hard glass is attached to the generator, and is heated, by
means of a Bunsen flame, at a point near to the flask. With this
arrangement the arsenic is deposited in the form of a dark, lustrous
ring just beyond the heated part. Zinc of .special purity must be
employed for generating the hydrogen, as all common specimens of
the metal contain a sufficient amount of arsenic, to give the metallic
710 INORGANIC CHEMISTRY
film without any special addition of an arsenic compound, and a blank
experiment must be run, with other portions of the same reagents, to
guard against the possibility of its coming from any of them.
Arsine is exceedingly poisonous, the breathing of small amounts
producing fatal effects. It differs from ammonia more markedly than
does phosphine, for it is not only without action on acids, but does not
unite directly even with the halides of hydrogen.
Halides of Arsenic. — The representatives of this class include a
liquid trifluoride AsF8, a pentafluoride, which is obtained only as
a double compound with potassium fluoride, a liquid trichloride AsCls,
a solid tribromide AsBr3, and a solid tri-iodide AsI8.
The trichloride, which is prepared by passing chlorine gas into a
vessel containing arsenic, is easily formed as the result of a vigorous
action. It is a colorless liquid, boiling at 130°. When mixed with
water it is at once converted into the white, almost insoluble trioxide.
The action is presumably similar to that of water upon the correspond-
ing compound of phosphorus (p. 181), but the arsenious acid for the
most part loses water and forms the insoluble anhydride :
AsCl3 + 3H20 <=± As(OH)8 + 3HC1,
2As(OH)3 <=> As2O3 J + 3H20.
This action, however, differs markedly from the other in that it is
reversible, and arsenic trioxide interacts with aqueous hydrochloric
acid, giving a solution of arsenious chloride. When this solution is
boiled, the volatility of the arsenious chloride causes it to be carried
over with the hydrochloric acid (b.-p. 110°, cf. p. 182), and this method
of separating arsenic from other substances is used in chemical
analysis.
Oxides of Arsenic; — Two oxides are known — the trioxide
As203 and the pentoxide As20B. Arsenic trioxide is produced by
burning arsenic in the air and during the roasbing of arsenical ores
(p. 708), and is known as " white arsenic" or simply "arsenic." It is
purified for commercial purposes by subliming the flue-dust in cylin-
drical pots. The pure trioxide is deposited in the glassy form in
the upper part of the vessel. It passes slowly from this amorphous
condition into the common crystalline variety. Its vapor density in-
dicates that it possesses the molecular weight As406, but the simpler
formula expresses its chemical properties sufficiently well.
ARSENIC, ANTIMONY, BISMUTH 711
When treated with water, the trioxide dissolves to a very slight
extent (0.3 : 100), forming arseiiious acid, by reversal of the second of
the equations given above. As usual, the less stable, amorphous variety
is the more soluble (1 : 100). In boiling water the solubility is eventu-
ally much greater" (11.5 : 100), but a condition of equilibrium is reached
very slowly. With concentrated sulphuric acid the trioxide forms a sul-
phate of rather complex composition, indicating that it has basic proper-
ties, but this sulphate is decomposed into the oxide and sulphuric acid
when treated with water. When heated in a tube with carbon, this ox-
ide is reduced, and the free element, being volatile, is deposited upon
the cold part of the tube just above the flame. It is an active poison,
since it gradually passes into solution, forming arsenious acid.
The pentoxide is a white crystalline substance, formed by heating
arsenic acid :
2H3As04 -> As205 + 3H20.
When raised to a higher temperature, it loses a part of its oxygen,
leaving the trioxide. In consequence of this instability, it cannot be
formed by direct union of oxygen with the trioxide, after the manner
of phosphorus pentoxide.
Acids of Arsenic. — When elementary arsenic or arsenious ox-
ide is treated with concentrated nitric acid, or with chlorine and
water, arsenic acid is produced. The substance, which is a deliques-
cent white solid, possesses a composition similar to that of orthophos-
phoric acid, and when heated loses water in progressive stages, fur-
nishing intermediate acids — pyroarsenic acid and metarsenic acid —
and finally the pentoxide. The relationship of the substances
(cf. p. 465) is shown by the formulae : H3As04 — »• H4As2O7 — > HAs03
— > As2O6. These acids, however, differ from the corresponding com-
pounds of phosphorus in that upon solution in water they immediately
pass back into the ortho-acid. The final elimination of all the water
by simple heating is also impossible with metaphosphoric acid. Many
salts of these acids are known, those of orthoarsenic acid being iso-
morphous with the corresponding salts of phosphoric acid. The red-
dish-brown silver orthoarsenate and the white MgNH4As04 resemble
the corresponding phosphates in being insoluble in water.
Arsenious acid, like sulphurous and carbonic acids, loses water,
and yields the anhydride, arsenic trioxide, when the attempt is made
to obtain it from the aqueous solution. The potassium and sodium
712 INORGANIC CHEMISTRY
arsenites, K3As03 and N~a3As03, are made by treating arsenic trioxide
with caustic alkalies, and are much hydrolyzed by water. The arsen-
ites of the heavy metals are insoluble, and can be made by precipita-
tion. Paris green and Scheele's green (p. 624) are arsenites of
copper. The poisonous effects of wall-paper colored with these com-
pounds seem to be due to volatile organic derivatives of arsine which
are formed by the action of a mold. In cases of poisoning by white
arsenic, freshly precipitated ferric hydroxide or magnesium hydroxide
is administered, since by interaction with the arsenious acid they form
insoluble arsenites. The salts of arsenious acid are readily oxidized,
passing into arsenates. The action of a standard solution of iodine
upon sodium arsenite, for example, is used in volumetric analysis :
Na3AsO3 + H2O + I2 +± Na3As04 + 2HL
Sulphides of Arsenic. — Three sulphides of arsenic are known,
As2S6, As2S3, As2S2. The first, arsenic pentasulphide, is obtained as a
yellow powder by decomposition of the sulpharsenates (see below), and
by leading hydrogen sulphide into a solution of arsenic acid in concen-
trated hydrochloric acid. The latter action seems to show that the
ion As""***, derived from AsCl5, is present in the solution.
Arsenious sulphide As2S3 occurs in nature as orpiment, and was
formerly used as a yellow pigment (auripigmentmn). The word
arsenic is derived from the Greek name for this mineral (dpo-m/cov).
It is obtained as a citron-yellow precipitate when hydrogen sulphide
is led into an aqueous solution of arsenious chloride. The slowness
with which the precipitate appears is due to the temporary formation
of a colloidal solution of the sulphide. When arsenious acid solution
is employed, the precipitation may be delayed for days. A beam of
light falling upon the solution.- is dispersed, and the liquid is thus shown
to be filled with minute particles, which ultimately collect in flocculent
form.
Realgar As2S2 is a natural sulphide of orange-red color, and is also
manufactured by subliming a mixture of arsenical pyrites and pyrite :
2FeAsS + 2FeS2 -> 4FeS + AsaS2J.
It burns in oxygen, forming arsenious oxide and sulphur dioxide, and
is mixed with potassium nitrate and sulphur to make " Bengal lights."
SulpJiarsenites and Sulpharsenates. — The sulphides of ar-
senic interact with solutions of alkali sulphides after the manner of
ARSENIC, ANTIMONY, BISMUTH 713
the sulphides of tin (p. 697), giving soluble, complex sulphides. Arsen-
ious sulphide with colorless ammonium sulphide gives ammonium sul-
pharsenite, and with the yellow sulphide gives ammonium sulphar-
senate :
- 3(NH4)2S + As2S3 -> 2(NH4)3AsS8,
3(NH4)2S + As2S3 + 2S -> 2(NH4)8AsS4.
Proustite (p. 626) is a natural sulpharsenite of silver.
There are also metasulpharsenites and pyro- and metasulphar-
senates of the forms KAsS2, K4As2S7, and KAsS8. In fact, there are
sulpho-compounds corresponding to all the forms of the oxygen acids.
The particular products present in any solution depend on the concen-
trations used in making it. The formation of these soluble compounds
is used in analysis (cf. p. 661).
Since these salts in solution furnish the ions AsS/" and AsS4///)
which with hydrogen ions give a feebly ionized sulpho-acid, they are
decomposed by acids, and give free sulpharsenious or sulpharsenic
acid :
(NH4)3AsS3 + 3HC1 -» H3AsS3 + 3NH4C1,
(NH4)3AsS4 + 3HC1 -> H3AsS4 + 3NH4C1.
These sulpho-acids, however, are unstable, and at once break up, giving
hydrogen sulphide as a gas, and the sulphides of arsenic as yellow
precipitates :
2H3AsS3 -Hi 3H2S + As2S8,
2H8AsS4 -» 3H2S + As2S
ANTIMONY.
The Chemical Relations of the Element. — Antimony resembles
arsenic in forming a hydride SbH3 and halides of the forms SbX3 and
SbX5. The latter are partially hydrolyzed by water with ease, but
complete hydrolysis is difficult to accomplish with cold water. The
oxide Sb208 is basic as well as feebly acidic, and the oxide Sb205 is
acidic. Aside from the salts derived from the oxide Sb208, such as
Sb2(S04)3, the compositions of the compounds are similar to those of
the compounds of arsenic. The element gives complex sulphides like
those of arsenic.
Occurrence and Preparation. — Antimony is found free in
nature to a small extent, The chief supply, however, is furnished
by the black trisulphide Sb2S3, known as stibnite, which is found in
Hungary and Japan, and forms shining, prismatic crystals of thf
rhombic system.
Native stibnite is roasted in the air in order to remove the sulphur,
and the white oxide which remains is mixed with carbon and reduced
by strong heat :
Sb2S3 + 502 -> Sb.2O4 + 3S02,
Sb204 + 40 -» 2Sb + 4CO.
Properties. — Antimony is a white, crystalline metal. It is brit-
tle, and easily powdered. Its vapor at 1640° has a density correspond-
ing to the formula Sb2, while at lower temperatures Sb4 is present.
It is used in making alloys such as type-metal, stereotype-metal, and
britannia metal (q-v.). The alloys of antimony expand during solidifi-
cation, and therefore give exceptionally sharp castings.
The element unites directly with the halogens. It does not rust,
but when heated it burns in the air, forming the trioxide Sb2O3 or a
higher oxide Sb204. When heated with nitric acid, it yields the tri-
oxide or antimonic acid (H3Sb04), according to the concentration of
the nitric acid and the temperature employed. When heated with
concentrated sulphuric acid, it forms the sulphate Sb2(S04)3.
Stibine. — The antirnonide of hydrogen SbH3 is formed by the
action of zinc and hydrochloric acid on a soluble compound of anti-
mony. By the action of dilute, cold hydrochloric acid on an alloy of
antimony and magnesium (1 : 1), a mixture of hydrogen and stibine
containing as much as 11.5 per cent (by volume) of the latter may be
made. It is separated by cooling with liquid air, and gives a liquid
boiling at —18° and freezing at —91.5°. It is more easily dis-
sociated than arsine, and forms a deposit of antimony when a porce-
lain vessel is held in the flame or when the gas passes through a
heated tube. The behavior under this treatment is in all respects
similar to that of arsine (p. 709).
The layers of arsenic or antimony obtained upon white porcelain
in Marsh's test (p. 709) may be distinguished readily in several ways.
The arsenic spots are brownish in color, lustrous, and volatile. The
antimony spots are black, smoky-looking, and involatile at the tempera-
ture of the Bunsen flame. The arsenic spots dissolve in dilute nitric
acid, while those of antimony do not. The arsenic spots dissolve in a
fresh solution of bleaching powder, producing calcium chloride and
ARSENIC, ANTIMONY, BISMUTH 715
arsenic acid, while those of antimony are unaffected. The arsenic
spots are scarcely attacked by a solution of yellow ammonium sulphide,
while those of antimony dissolve readily, forming an ammonium sul-
phantimoniate. Another distinction between arsine and stibine is found
in their action upon a solution of nitrate of silver. Stibine precipi-
tates a silver antimonide Ag3Sb, and none of the antimony remains in
the solution. Arsine, on the contrary, precipitates metallic silver,
while arsenious acid remains in the solution.
Antimony Halides. — The halogen compounds of antimony which
are known include the trichloride, a solid melting at 73° ; the penta-
chloride, a liquid ; the tribromide, tri-iodide, trifluoride, and pentafluo-
ride.
Antimony trichloride SbCl3 is the most familiar of these substances.
It forms large, soft crystals, and used to be named " butter of anti-
mony." When treated with water, it forms a white, opaque, insoluble
basic salt. When little water is used, the product is antimony oxy-
chloride :
SbCl8 + H20 <=» SbOCl J + 2HC1.
With a large amount of water, a greater proportion of the chlorine is
removed, and Sb405Cl2 ( = 2SbOCl,Sb203) remains. With boiling water
the oxide is finally formed. The action is not complete as long as
hydrochloric acid is present. It may therefore be reversed, so that on
addition of hydrochloric acid to the mixture, a clear solution of the tri-
chloride is re-formed. If the concentration of the acid is once more
reduced by dilution with water, the oxychloride is again precipitated.
The pentachloride is formed by leading chlorine over the trichloride.
It is a liquid (b.-p. 140°) which fumes strongly in the air, being hydro-
lyzed by the moisture.
Oxides of Antimony. — Three oxides are known: antimony tri-
oxide Sb203, antimony pentoxide Sb205, and an intermediate oxide
Sb204. The trioxide is obtained by oxidizing antimony with nitric acid,
or by combustion of antimony with a limited supply of oxygen. It
is a white substance, insoluble in water. It is in the main a basic
oxide, interacting with many acids to form salts of antimony. The
pentoxide is a yellow, amorphous substance, obtained by heating anti-
monic acid. It combines with bases to form salts, and is therefore an
acid-forming oxide exclusively. The intermediate compound Sb204 is
716 INORGANIC CHEMISTRY
formed by heating antimony or the trioxide in oxygen. It is the most
stable of the three oxides. It is neither acid- nor base-forming, and
may be antimoniate of antimony (SbSb04).
The hydrated trioxide Sb(OH)3 may be obtained as a white pre-
cipitate by adding dilute sulphuric acid to tartar-emetic (see below).
It is insoluble, and easily loses water, giving the trioxide.
The trioxide interacts with potassium and sodium hydroxides,
forming soluble antimonites, but the latter are much hydrolyzed by
water, and cannot be isolated in solid form.
Salts of Antimony. — The nitrate Sb(N03)3 and the sulphate
Sb2(S04)3 are made by the interaction of the trioxide with nitric and
sulphuric acids. They are hydrolyzed by water, giving basic salts,
such as (SbO)2S04 ( = Sb202S04), which, like SbOCl, are derived from
the hydroxide SbO(OH). When the trioxide is heated with a solution
of potassium bitartrate KHC4H406, a basic salt K(SbO)C4H406, known
as tartar-emetic, is formed. This is a white, crystalline substance
which is soluble in water and is used in medicine. The univalent
group SbO1 is known as antimonyl, and the above mentioned basic
compounds are often called antimonyl sulphate, etc.
Antimonic Acid. — By vigorous oxidation of antimony with nitric
acid, or by decomposing the pentachloride completely with water, a
white, insoluble substance of the approximate composition H3Sb04 is
obtained. This substance interacts with caustic potash and passes into
solution. But the salts which have been made are pyro- and met-
antimoniates. Thus, when antimony is fused with niter, potassium
metantimoniate KSb03 is formed. When dissolved, this salt takes up
water, and forms a solution of the acid pyroantimoniate :
2KSb03 + H20 -> K2H2Sb207.
If this is added 'to a strong solution of a salt of sodium, a sodium
pyroantimoniate is thrown down, Na2H2Sb207. The same insoluble
body, almost the only salt of sodium which deserves this name, is
formed also by direct action of sodium hydroxide upon antimonic
acid.
Sulphides of Antimony. — There are two sulphides, the trisul-
phide Sb2S3 and the pentasulphide Sb2S6. The trisulphide is found in
nature as the black, crystalline stibnite. By the action of hydrogen
AESENIC, ANTIMONY, BISMUTH 717
sulphide upon solutions of salts of antimony, the trisulphide is precipi-
tated as an orange-red powder, which, however, after having been
melted, assumes the appearance of stibnite :
2SbCl3 + 3H2S <=± Sb2S3 J+ 6HC1.
The antimony trisulphide is decomposed, and the above action is re-
versed, by concentrated hydrochloric acid. Like cadmium sulphide,
this .substance is formed only when the acid present is dilute.
The pentasulphide is obtained by the decomposition of sulphanti-
moniates (see below). In appearance it resembles the trisulphide and,
when heated, it decomposes very readily into this substance and free
sulphur. It is used for vulcanizing rubber.
Sulphantimonites and Sulphantimoniates. — The behavior of
the sulphides of antimony towards solutions of the alkali sulphides is
very similar to that of the sulphides of arsenic (p. 712). The tri-
sulphide dissolves in colorless ammonium sulphide with difficulty,
forming an unstable ammonium sulphantimonite :
Sb2S3 + 3(NH4)2S -* 2(NH4)3SbS3.
With the pentasulphide or with yellow ammonium sulphide the action
takes place more readily and ammonium sulphantimoniate is formed :
Sb2S5 + 3(NH4)2S -* 2(NH4)8SbS4,
Sb2S3 + 3(NH4)2S + 2S->2(NH4)3SbS4.
The most familiar substance of this class is Schlippe's salt Na3SbS4,
9H20. Pyrargyrite Ag3SbS3 (p. 626) is a native sulphantimoniate.
When acids are added to solutions of sulphantimoniates, the sul-
phantimonic acid which is liberated decomposes, and antimony penta-
sulphide is thrown down (see under Arsenic, p. 713).
BISMUTH.
The Chemical Relations of the Element. — Bismuth forms no
compound with hydrogen. Its compounds with the halogens are of
the form BiX3 and are hydrolyzed by water giving basic salts. The
oxide Bi203 is basic, and, although an oxide Bi205 is known, it is not
acidic. Bismuth gives a carbonate, nitrate, sulphate, phosphate, and
other salts, in all of which it acts as a trivalent element. It forms no
soluble complex sulphides.
718 INORGANIC CHEMISTRY
Occurrence and Physical Properties. — This element is found
free in nature, and also to some extent as trioxide Bi203 and trisul-
phide BijSg. It is a shining, brittle metal with a reddish tinge. It
melts at about 270°. When converted into vapor, its density at 1600-
1700° is somewhat less than that corresponding to the formula Bi2.
Mixtures of bismuth with other metals of low melting-point fuse
at lower temperatures than do the separate metals. This is another
illustration of the fact that a solution melts at a lower temperature than
the pure solvent (p. 532). Thus, Wood's metal, containing bismuth
(m.-p. 270°) 4 parts, lead (m.-p. 326°) 2 parts, tin (m.-p. 233°) 1 part,
and cadmium (m.-p. 320°) 1 part, melts at 60.5°, considerably below the
boiling-point of water. Similar alloys are used for safety plugs in
steam-boilers, and, in the chemical laboratory, for filling baths in which
a uniform temperature higher than 100° is to be maintained.
Chemical Properties. — Bismuth does not tarnish, but when
heated strongly in the air it burns to form the trioxide. With the
halogens it forms a fluoride, a bromide, and an iodide, in all of which the
element is trivalent. When the metal is treated with oxygen acids, or
the trioxide with any acids, salts are produced.
Oxides. — In addition to the basic trioxide, which is a yellow
powder obtained by direct oxidation of the metal or by ignition of the
nitrate, three other oxides are known — BiO, Bi204, and Bi205. None of
these, however, is acid-forming, or gives corresponding salts when
treated with acids.
Salts of Bismuth. — The salts of bismuth, when dissolved in
water, are decomposed in a manner which recalls the behavior of the
compounds of antimony. The products are insoluble basic salts, and
the actions are reversible, the basic salts being redissolved by addition
of an excess of the acid. In the case of the chloride BiCl3, H20 and the
nitrate Bi(NO3)3,5H2O, the actions taking place are :
BiCl3 + 2H20 <=» Bi(OH)2Cl + 2HC1,
Bi(N03)3 + 2H20 -+ Bi(OH)2N03 + 2HM)3.
The former of these products, when dried, loses a molecule of water,
giving the oxychloride BiOCl. The oxynitrate of bismuth is much
used in medicine under the name of "subnitrate of bismuth."
ARSENIC, ANTIMONY, BISMUTH 719
It will be seen that, although bismuth forms a colorless ion Bi%",
and is in this respect a metal in the chemical sense of the term, yet,
like many other metals, it is related to the non-metals inasmuch as its
salts are at least partially hydrolyzed by water.
Bismuth forms a trisulphide Bi2S8, which may be obtained by
direct union of the elements, or by the action of hydrogen sulphide
upon solutions of bismuth salts :
2BiCls + 3H2S <=> Bi2S8| + 6HC1.
It is a brownish-black, insoluble substance, but on addition of much
acid the above action is reversed. This sulphide is not affected by
solutions of ammonium sulphide or of potassium sulphide. It differs,
therefore, markedly from the sulphides of arsenic and antimony in its
behavior.
THE FAMILY AS A WHOLE.
When we compare the elements of this group, taking nitrogen as
the first of the family in spite of the fact that it is somewhat less
closely related to the other members than they are to one another, we
find an admirable illustration of the general principles which the
periodic system presents.
The elements themselves change progressively in physical proper-
ties as the atomic weight increases. Nitrogen is a gas which with
sufficient cooling yields a white solid, phosphorus an almost white, or
a red solid, and arsenic, antimony, and bismuth are metallic in
appearance. The first combines directly with hydrogen, the next
three give hydrides indirectly, and the last does not unite with
hydrogen at all. The hydride of nitrogen combines with water
to form a base, while the other hydrides show no such tendency.
Ammonia unites with all acids, including those of the halogens, to
form salts; phosphine with the hydrogen halides only; the others do
not combine with acids at all. As regards their metallic properties,
in the chemical sense, nitrogen and phosphorus do not by themselves
form positive ions, and furnish us therefore with no salts whatever.
Arsenic gives a trivalent positive ion, which is found in solutions of
the halides only. It forms no normal sulphates, nitrates, or other
salts. Antimony and bismuth both give trivalent positive ions. The
sulphates, nitrates, etc., of antimony, however, are readily decomposed
by water with precipitation of the hydroxide. The salts of bismuth,
720 INORGANIC CHEMISTRY
on the other hand, do not readily give the pure hydroxide with water,
although they are easily hydrolyzed to basic salts.
The halogen compounds of nitrogen and phosphorus are completely
hydrolyzed by water, and do not exist when any water is present, even
when excess of the halogen acid is used. The halogen compounds of
arsenic are completely hydrolyzed by cold water, but exist in solution
in presence of excess of the acids. The halogen compounds of anti-
mony and bismuth are incompletely hydrolyzed by cold water.
Nitrogen forms five different oxides, a larger number than that
furnished by any of the other elements of this group. The oxides
upon which we naturally fix our attention are the trioxide and pent-
oxide in each case. With nitrogen these are acid-forming, being the
anhydrides of nitric and nitrous acids. With phosphorus the trioxide
and the pentoxide are anhydrides of acids. With arsenic the trioxide
is basic towards the halogen acids, and is the first example of a basic
oxide which we encounter in this group. The pentoxide, however, is
acid-forming. The trioxide of antimony is mainly base-forming,
although it is feebly acid-forming also. The pentoxide is acid-forming.
The trioxide of bismuth is base-forming exclusively, and the pentoxide
has no derivatives.
If the chemical actions which these elements and their compounds
undergo were examined more closely, this progressive change in the
properties with change in atomic weight could be developed much
further. We have said sufficient, however, to show that when the
periodic law is borne in mind it furnishes valuable aid in systematizing
the chemistry of a group like this.
Analytical Reactions of Arsenic, Antimony, and Bismuth. —
The ions which are most frequently encountered are As*", Sb***, Bi*",
AsO/", and As03'". The first three, with hydrogen sulphide, give
colored sulphides which are not affected by dilute acids. The sul-
phides of arsenic and antimony are separable from the sulphide of
bismuth by solution in yellow ammonium sulphide. The ion of the
arsenates AsO/" is identified by its interaction with salts of silver and
the formation of MgNH4As04, while that of the arsenites AsO/" is rec-
ognized by its reducing power. By means of Marsh's test and the facts
described- under it (p. 714), the presence of traces of compounds of
Arsenic and antimony may be recognized and the elements may be
distinguished. Oxygen compounds of arsenic, when heated with car-
bon, give a volatile, metallic-looking deposit of arsenic.
ARSENIC, ANTIMONY, BISMUTH 721
VANADIUM, COLUMBITTM, TANTALUM.
Of these elements, vanadium is less uncommon than the others. It
is found in rather complex compounds. When these are heated with
soda and sodium nitrate, sodium vanadate is formed, and can be ex-
tracted with water. Solid ammonium chloride is added to the solution,
and ammonium metavanadate NH4y08, which is less soluble in solu-
tions of salts of ammonium (of. p. 584) than in water, appears in the
form of yellow crystals. When this salt is heated, vanadic anhydride
V205, a yellowish-red powder, remains. This oxide interacts with bases
giving vanadates, of which the most stable are the metavanadates.
The element forms several chlorides, such as VCLj, VC13, VC14, VOClg,
and five oxides, V20, VO, V203, V02, and V205. The element has very
feeble base-forming properties, and gives only a few unstable salts.
Columbium (or niobium) and tantalum likewise possess feebly base-
forming properties, their chief compounds being the columbates and
tantalates.
Exercises. — 1. How do you account for the fact that the molec-
ular weight of arsenic at 644° is not exactly 300, and why is 308.4
-5- 4 not accepted as the atomic weight ? Could the atomic weight
be found from the determination of the vapor density of the free ele-
ment alone?
2. What should you expect to be the interaction of arsine with
concentrated nitric acid ?
3. Formulate the series of changes involved in the solution of
arsenic trioxide and the interaction of hydrochloric acid with the
arsenious acid so formed (cf. p. 371).
4. What is the full significance of the fact that arsenic penta-
sulphide may be precipitated by hydrogen sulphide from a solution of
arsenic acid in hydrochloric acid ?
5. To what classes of chemical changes do the interactions of
arsenious sulphide and antimony trisulphide with yellow ammonium
sulphide belong?
6. Construct equations showing the interaction of (a) concentrated
sulphuric acid and antimony, (6) arsenic and bleaching-powder solu-
tion, (c) antimony and yellow ammonium sulphide, (d) silver nitrate
and stibine, (e) silver nitrate and arsine, (/) concentrated nitric acid
and antimony, (</) acids and ammonium orthosulphantimoniate.
7. How should you set about making Schlippe's salt ?
CHAPTER XLII
THE CHROMIUM FAMILY. RADIUM
THE chromium (Cr, at. wt. 52.1) family includes molybdenum (Mo,
at. wt. 96), tungsten (W, at. wt. 184), and uranium (U, at. wt. 238.5),
and occupies the seventh column of the periodic table along with the
sulphur and selenium family.
The Chemical Relations of the Family. — The features which
are common to the four elements are also those which affiliate them
most closely with their neighbors on the right side of the column.
They yield oxides of the forms Cr08, Mo08, W08, and U03, which, like
S08, are acid anhydrides, and show the elements to be sexivalent.
They give also acids of the form H2X04, corresponding to sulphuric
acid, whose salts resemble the sulphates. Thus, sodium chromate
Na^OO^ lOH^O is isomorphous with Glauber's salt (p. 576), and potas-
sium chromate KgCrC^ with potassium sulphate.
Aside from the chromates, the first element forms also two basic
hydroxides Cr(OH)2 and Cr(OH)8, from which the numerous chromous
(Cr**) and chromic (Cr"") salts are derived. Uranium forms a dioxide
UO2, to which correspond the uranous salts like U(S04)2, but the most
familiar salts of this metal are basic salts of the oxide U08, and have
the form UO2(N03)2. Molybdenum and tungsten are not base-forming
elements.
CHROMIUM.
The Chemical Relations of the Element. — Chromium gives
four classes of compounds, and most of them are colored substances
(Gk. xpw/xa, color). The chromates are derived from chromic acid
H2CrO4, which, however, is itself unstable, and leaves the anhydride
CrO3 when its solution is evaporated. The oxide and hydroxide in
which the element is trivalent, namely Cr203 and Cr(OH)3, are weakly
basic and still more weakly acidic. Hence we have chromic salts such
as CrClg and Cr2(S04)3 which are somewhat hydrolyzed, but no carbon-
ate, and no sulphide which is stable in water. The compounds in
which the same hydroxide acts as an acid are the chromites, and are
722
THE CHROMIUM FAMILY. RADIUM 723
derived from the less completely hydrated form of the oxide CrO(OH).
Potassium chromite K.CrO2 is more easily hydrolyzed, however, than is
potassium zincate or potassium aluminate. Finally, the chromous salts
such as CrCl2 and CrS04 correspond to chromous hydroxide Cr(OH)2
in which the element is bivalent. This hydroxide is more distinctly
basic than is chromic hydroxide, and forms a carbonate and sulphide
which can be precipitated in aqueous solution. The chromous salts
resemble the stannous and ferrous (q.v.) salts in being easily oxidized
by the air.
Occurrence and Isolation. — Chromium is found chiefly in
ferrous chromite Fe(Cr02)2, which constitutes the mineral chromite,
and in crocoisite PbCr04, which is chromate of lead. The metal may
be made by heating chromic oxide with carbon in the electric furnace,
or, more easily, by reduction of the oxide with aluminium filings by
Goldschmidt's method (p. 540).
Physical and Chemical Properties. — Chromium is steel-gray
in color, very hard, and extremely infusible. It does not tarnish, but
when heated it burns in oxygen, giving the green chromic oxide Cr2O3.
It seems to exist in two states, an active and a passive one, the relations
of which are still somewhat obscure. A fragment which has been made
by the Goldschmidt method, or has been dipped in nitric acid, is passive,
and does not displace hydrogen from hydrochloric acid. When, how-
ever, the specimen is warmed with this acid', it begins to interact, and
thereafter behaves as if it lay between zinc and cadmium in the elec-
tromotive series. If left in the air, it slowly becomes inactive again.
Tin and iron with hydrochloric acid form stannous and ferrous chloride
respectively, because the higher chlorides, if present, would be reduced
by the nascent hydrogen. Here, for the same reason, chromous
chloride and not chromic chloride is formed :
Cr + 2HC1 -+ CrCl2 + H3 or Cr + 2H' -> Cr"+ H2.
DERIVATIVES OF CHKOMIC ACID.
Potassium Chromate. — This and the sodium salt, or rather the
corresponding dichromates (see below), are made directly from chro-
mite, and form the starting-point in the preparation of the other com-
pounds of chromium. The finely powdered mineral is mixed with
potash and limestone, and roasted. The lime is employed chiefly to
724 INORGANIC CHEMISTRY
keep the mass porous and accessible to the oxygen of the air, the
potassium compounds being easily fusible :
4Fe(Cr02)2 + 8K2C08 + 702 -> 2Fe2O3 + 8K2Cr04 + 8C02.
The iron is oxidized to ferric oxide, and the chromium passes from the
state of chromic oxide in the chromite (FeO,Cr2Os) to that of chromic
anhydride in the potassium chromate (K20,Cr03). Thus, more insight
is given into the nature of the action by the equation
4(FeO,Cr208) + 8(K20,C02) + 7O2 ~> 2Fe208 4- 8(K20,CrO,).+ 8C02.
The cinder is treated with hot potassium sulphate solution. This
interacts with the calcium chromate, which is formed at the same
time, giving insoluble calcium sulphate :
CaCr04 + K2S04 <± CaS04 j+ K2O04.
The whole of the potassium chromate goes into solution.
Potassium chromate is pale-yellow in color, rhombic in form (iso-
morphous with potassium sulphate), and is very soluble in water
(61 : 100 at 10°).
Sodium chromate is made by using sodium carbonate in the
process just described.
The Dichromates. — When a solution of potassium sulphate is
mixed with an equivalent amount of sulphuric acid, potassium bisul-
phate is obtainable by evaporation : K2S04 + H2S04 — » 2KHSO4. The
dry acid salt, when heated, loses water (p. 388), giving the pyrosulphate
(or disulphate) : 2KHS04 <=± K2S207 -f- H2O, but the latter, when redis-
solved, returns to the condition of acid sulphate. Now, when an acid
is added to a chromate we should expect the chromic acid H2Cr04, thus
liberated, to interact, giving an acid chromate (say, KHCr04). No acid
chromates are known, however, and instead of them, pyrochromates or
dichromates are produced, with elimination of water. In other words,
the second of the above actions is not appreciably reversible when
chromates are in question :
K2Cr04 H
K2Cr()4 (
- H2S04
+ H2Cr04) -
* (H2Cr04)
-» K2Cr207 +
-f- K2S04
H20
2K2Cr04 + H3S04 -» K2Cra07 + HaO + K2S04
THE CHROMIUM FAMILY. RADIUM 725
In terms of the ionic hypothesis, S207" is unstable in water, and in-
teracts with it, giving hydrion and sulphanion, while Cr207" is stable
in water and is formed from the interaction of hydrion and chromanion :
S207" + H20 ±3 2H' + 2S04",
Cra07" + H20 <=? 2H* + 2Cr04". (2)
The dichromates of potassium and sodium are made by adding
sulphuric acid to the crude solution of the chromate obtained from
chromite (p. 724). They crystallize when the liquid cools, and the
mother-liquor, containing the potassium sulphate and undeposited
dichromate, is used for extracting a fresh portion of cinder. As the
dichromates are much less soluble than the chromates, they crystallize
from less concentrated solutions, and can therefore be obtained in
purer condition. For this reason the extract is always treated for
dichromate.
Potassium dichromate K2Cr207 (or KgCrO^CrOg) crystallizes in
asymmetric tables of orange-red color. Its solubility in water is 8 :
100 at 10° and 12.5 : 100 at 20°. Sodium dichromate Na2Cr207, 2H20
forms red crystals also, and its solubility is 109 : 100 at 15°. This salt
is now cheaper than potassium dichromate, and has largely displaced
the latter for commercial purposes.
By treatment of the chromates with larger amounts of free acid,
other polychromates are formed. Thus, with increasing amounts of
nitric acid, ammonium chromate gives first the dichromate (NH4)2Cr207,
which may be written (NH4)2Cr04,Cr03, then the trichromate (NH4)2
Cr04, 2Cr03, and even the tetrachromate (NH4)2Cr04, 3Cr03, all of which
are red crystalline substances.
Chemical Properties of the Dichromates. — 1. When concen-
trated sulphuric acid is added to a strong solution of a dichromate (or
chromate), chromic anhydride separates in red needles :
Na2O207 + H2S04 -> Na2S04 + H2O + 2CrO8.
2. Although a dichromate lacks the hydrogen, it is essentially of
the nature of an acid salt, just as SbOCl lacks hydroxyl, but is essen-
tially a basic salt. Hence, when potassium hydroxide is added to a
solution of potassium dichromate, potassium chromate is formed :
K2Cr20T + 2KOH -> 2K2Cr04 + H20.
726 INORGANIC CHEMISTRY
The solution changes from red to yellow, and the chromate is obtained
by evaporation. It is in this way that pure alkali chromates are
made.
3. By addition of potassium dichromate to a solution of a salt of a
metal whose chromate is insoluble, the chromate and not the dichro-
mate is precipitated. This is in consequence of the fact, that there is
always a little hydrion and Cr04" (equation (2) above) in the solution
of the dichromate :
2Ba(N03)2 4- KjCr207 + H20 <=» 2BaCr04J + 2KN03 + 2HN08.
Being essentially an acid salt, the dichromate produces a salt and an
acid, as any acid salt would do. For example :
Ba(N08)2 + KHS04 <=± BaSO4| + KN08 + HN03.
Soluble chromates are, naturally, equally good precipitants of insoluble
ones.
4. The dichromates of potassium and sodium melt when heated, and,
at a white heat, decompose, giving the chromate, chromic oxide, and free
oxygen. To make the equation, we note that the dichromate, for
example K2Cr207, consists of K2Cr04 + Cr08, and the latter, if alone,
will decompose thus : 2O03 — » Cr203 + 30. Since the product must
i L contain a multiple of 02, the equation is :
4K2O207 -> 4K2Cr04 + 2O203 + 302.
5. With free acids the dichromates give powerful oxidizing mix-
tures, in consequence of their tendency to form chromic salts. Since
the latter correspond to the oxide Cr2O3 and the former to Cr03, the
passage from the former to the latter must furnish 30 for every
2Cr03 transformed. In dilute solutions, unless a body capable of
being oxidized is present, no actual decomposition, beyond the libera-
tion of chromic acid * occurs. When concentrated hydrochloric acid is
used, this acid itself suffers oxidation :
K2O2O7 + 8HC1 -+ 2KC1 + 2CrCl3 + 4H20 ( + 30)
(30) + 6HC1 -> 3H2O + 3C1, _
K2Cr2O7 + 14HC1 -> 2KC1 + 2CrCl3 + 7H20 + 3CL,
When sulphuric acid is employed, an oxidizable substance such as
* Not shown as a distinct stage in the subsequent equations.
THE CHROMIUM FAMILY. RADIUM 727
hydrogen sulphide (cf. p. 374), sulphurous acid, or alcohol must be
present :
K2O207 + 4H2S04 -> K2S04 + Crs(S04)8 + 4H2O ( + 3O) (1)
" (30) + 3H2SOS -* 3H2S04 (2)
or (30) + 3C2H5OH -> 3C2H40 1 + 3H2O (2')
[alcohol] [aldehyde]
In each case the usual summation of (1) and (2), with omission of the
3O gives the equation for the whole action. When (1) is dissected,
K20, 2O08 giving 3S03,Cr203+30 is found to be its essential content.
In practice, this sort of action is used for the purpose of making
chromic salts, and for its oxidizing effects, as in the preparation of
aldehyde and in the dichromate battery (q.v.~).
6. When a body which is not merely oxidizable, but is an active
reducing agent, is employed, the dichromate may be reduced without
the addition of any acid. For example, when warmed with ammonium
sulphide, a dichromate gives chromic hydroxide and free sulphur :
K2Cr207 + 3(NH4)2S + 7H20 _> 2Cr(OH)3 + 3S + 2KOH + 6NH4OH.
If the reducing body is less active, the change may nevertheless take
place under the influence of light. Thus, when paper is coated with
gelatine containing a soluble chromate or dichromate, and, after being
dried, is exposed to light, chromic oxide is formed by reduction, and
combines with the gelatine to give an organic compound. This product
will not swell up or dissolve in tepid water as does pure gelatine.
This action is used in many ways for purposes of artistic reproduction.
Thus, if the gelatine mixture is made up with lampblack, and, after the
coating has dried, is covered with a negative and exposed to light, the
parts which were protected from illumination may afterwards be
washed away, while the " carbon print " remains. The gelatine layer
can be transferred to wood or copper before washing. When materials
of different colors are substitxited for the lampblack, prints of any
desired tint may be made by the same process.
Insoluble Chromates. — A number of chromates, formed by pre-
cipitation with a solution of a soluble chromate or dichromate, are
familiar. Thus, lead chromate PbOr04 is vised as a yellow pigment.
By treatment with lime-water it gives a basic salt of brilliant red color
— " chrome-red " Pb2OCr04. Salts of calcium give a yellow, hydrated
calcium chromate CaCrO4, 2HO« analogous to gypsum, and, like it, per-
728 INORGANIC CHEMISTRY
ceptibly soluble in water (0.4 : 100 at 14°). Barium chromate BaCr04
is also yellow. Being a salt of a feeble acid, it interacts with active
acids, and passes into solution. Like calcium oxalate (cf. p. 599), it is
not soluble enough to be attacked by acetic acid. Strontium chromatei
however, is soluble in acetic acid. Silver chromate is red, and inter-
acts easily with acids. It will be observed that there is a close corre-
spondence between the relative solubilities of the chromates and the
sulphates.
Chromyl Chloride. — This compound corresponds to sulphuryl
chloride S02C1^, and is made by distilling a dichromate with a chloride
and concentrated sulphuric acid :
K2Cr207 + 4KC1 + 3H2S04 -» 2Cr02Cl2 + 3K2S04 + 3H20.
The hydrochloric acid liberated from the chloride may be supposed to
interact with chromic acid from the dichromate :
CrO2(OH)2 + 2HC1 -* CrO.Cl, + 2H20.
Chromyl chloride is a red liquid, boiling at 118°. It fumes strongly in
moist air, being hydrolyzed by water. This action is the reverse of
that shown in the last equation. No corresponding bromine and iodine
compounds are known ; and when a bromide or iodide is treated as de-
scribed above, the halogens are liberated by oxidation, and no volatile
compound of chromium appears. Hence, when an unknown halide is
mixed with potassium dichromate and sulphuric acid and distilled, and
the vapors are caught in ammonium hydroxide, the finding of a chromate
in the distillate demonstrates the existence of a chloride in the original
substance :
L, + 4NH4OH -> (NH4)2Cr04 + 2KH4C1 + 2H20.
This action is used as a test for the presence of traces of chlorides in
large amounts of bromides or iodides.
Chromic Anhydride. — The oxide Cr03 is made as described above
(par. 1, p. 725), and is often called chromic acid. It is soluble in
water, and combines with the latter to some extent, giving dichromic
acid H2.Cr2O7. In acidified solution it is much used as an oxidizing
agent for organic substances. It interacts with acids in the same way
as do the dichromates, giving chromic salts and furnishing oxygen to
THE CHROMIUM FAMILY. RADIUM 729
the oxidizable body. When heated by itself, it loses oxygen readily,
and yields the green chromic oxide : 4Cr03 — > 2O203 + 30a.
Perchrotnic Acid. — When hydrogen peroxide (cf. p. 306) is pres-
ent in solution with free chromic acid, a deep-blue, unstable compound
is formed. It can be extracted from the solution by means of ether,
and its formation is used as a delicate test for hydrogen peroxide
(q.v.}. It is supposed to be a mixed anhydride of hydrogen peroxide
with dichromic acid, perhaps H2Cr208 (cf. p. 397).
CHROMIC COMPOUNDS.
Chromic Chloride. — A hydrated chloride CrCi,, 6H20 is ob-
tained by treating the hydroxide Cr(OH)8 with hydrochloric acid and
evaporating. When heated, this is hydrolyzed, and chromic oxide re-
mains. The anhydrous chloride is formed by sublimation, as a mass of
brilliant, reddish- violet scales, when chlorine is led over a heated mix-
ture of chromic oxide and carbon (cf. Silicon tetrachloride) :
Cr208 + 3C + 3C12 -> 2CrClg + 3CO.
In this form the substance dissolves with extreme slowness, even in
boiling water, but in presence of a trace of chromous chloride or
stannous chloride it is easily soluble. The solution is green, as are
all solutions of chromic salts after they have been boiled, but on stand-
ing in the cold, blue crystals of CrCl3, 6H20 are deposited. These give
a blue solution containing Cr*" -f 3C1', but boiling reproduces the
green color. The green material is a basic salt of an exceptional
nature. It has lost one unit of chlorine by hydrolysis, and one of the
two others is not precipitated by nitrate of silver. The substance is
supposed, therefore, to contain a complex cation and to have the formula
CrClOH.Cl.
Chromic Hydroxide. — When ammonium hydroxide is added to
a solution of a chromic salt, a hydrated hydroxide of pale-blue color,
Cr(OH)3, 2H20, is thrown down. This loses water by stages, giving
intermediate hydroxides such as Cr(OH)3 and CrOOH, and finally
Cr208. It interacts with acids, giving chromic salts. It also dis-
solves in potassium and sodium hydroxides to form green solutions of
chromites of the form KCr02. When the solutions of the alkali
chromites are boiled, the free hydroxide, present in consequence of
730 INORGANIC CHEMISTRY
hydrolysis, is converted into a greenish, less completely hydrated, and
less soluble variety. This begins to come out as a precipitate, and
soon the whole action is reversed. Insoluble chromites, such as that
of iron Fe(Cr02)2, are found in nature. Many of them, like Zn(Cr02)2
and Mg(Cr02)2, may be formed by fusing the oxide of the metal with
chromic oxide ; the action being similar to that used in making zinc-
ates (p. 648) and aluminates (p. 691).
Chromic Oxide. — This oxide is obtained as a green, infusible
powder by heating the hydroxide ; or, more readily, by heating dry
ammonium dichromate ; or by igniting potassium dichromate with sul-
phur and washing the potassium sulphate out of the residue :
(NH4)2Cr207 -> Na + 4H20 + Cr208,
K2Cr207 + S -» K2S04 + Cra08.
Chromic oxide is not affected by acids, but may be converted into the
sulphate by fusion with potassium bisulphate. It is used for making
green paint, and for giving a green tint to glass. When the oxide, or
any of the chromic salts, is fused with a basic substance such as an
alkali carbonate, it passes into the form of a chromate, absorbing the
necessary oxygen from the air. If an alkali nitrate or chlorate is added,
the oxidation goes on more quickly. The alkaline solution of the
chromites may be oxidized, for example, by addition of chlorine or
bromine, and chromates are formed.
Chromic Sulphate, — This salt crystallizes in reddish-violet crys-
tals of a hydrate Cr2(S04)3,15H20, and may be made by treating the
hydroxide with sulphuric acid. It gives reddish-violet, octahedral
crystals of chrome-alum (cf. p. 687), K2S04,Cr2(S04)3, 24H20, when
mixed with potassium sulphate. This double salt is most easily ob-
tained by reducing potassium dichromate in dilute sulphuric acid by
means of sulphurous acid (p. 727), and allowing the solution to crystal-
lize. The solution of the crystals, either of the pure sulphate or of
the alum, is bluish-violet (Cr***), but when boiled becomes green. The
green compound is formed by hydrolysis and is gummy and uncrystal-
lizable. It even yields products which do not show the presence either
of the Cr*" or the SO/' ion. It seems to be formed thus :
2Cr2(S04)3 + H?0 <± Cr40(S04)4.S04 + H2S04.
THE CHROMIUM FAMILY. RADIUM 731
The green materials revert slowly to the violet ones by reversal of the
above action when the solution remains in the cold, and so crystals of
the sulphate or of the alum are obtainable from the green solutions.
Chromic Acetate. — This salt, Cr(C2H302)3, is made by treating
the hydroxide with acetic acid, and a green solution of it is used as a
mordant by calico-printers (cf. p. 689).
CHROMOUS COMPOUNDS.
By the interaction of chromium with hydrochloric acid, or by re-
ducing chromic chloride in a stream of hydrogen, chromous chloride
CrCl2 is formed. The anhydrous salt is colorless, and its solution is
blue (Cr"). Like stannous chloride, it is very easily oxidized by the
air, a solution of it containing excess of hydrochloric acid being used
in the laboratory to absorb oxygen :
4CrCl2 + 4HC1 + 02 -> 4CrCl3 + 2H20.
Chromous hydroxide is obtained as a yellow precipitate when alka-
lies are added to the chloride. With sulphuric acid it gives chromoua
sulphate CrS04, 7H20, which is isomorphous with the vitriols (p. 649).
Chromous salts give with ammonium sulphide a black precipitate
of chromous sulphide, and with sodium acetate a red precipitate of
chromous acetate. The latter is not very soluble, and is less quickly
oxidized by the air than any of the other chromous compounds.
Analytical Reactions of Chromium Compounds. — The
chromic salts give the bluish- violet trichromion Cr***, or the green com-
plex cations, and may be recognized in solution by their color. The
chromates and dichromates give the ions Cr04" and Cr207", which are
yellow and red respectively. From chromic salts, alkalies and ammo-
nium sulphide precipitate the bluish-green hydroxide, and carbonates
give a basic carbonate which is almost completely hydrolyzed to
hydroxide. By fusion with sodium carbonate and sodium nitrate, they
yield a yellow bead containing the chromate. The chromates and
dichromates are recognized by the insoluble chromates which they
precipitate, and by their oxidizing power when mixed with acids. All
compounds of chromium give a green borax bead containing chromic
bo rate, and this bead differs from that given by compounds of copper
(cf. p. 626), which is also green, in being unreducible.
732 INORGANIC CHEMISTRY
MOLYBDENUM, TUNGSTEN, URANIUM.
As was stated at the opening of the chapter, these elements give
acid anhydrides of the form X08, and acids and salts of the form
H2X04. They also give salts of the form It^K201 corresponding to
the dichromates. Uranium has base-forming properties as well.
Molybdenum. — This element is found chiefly in wulfenite
PbMo04 and molybdenite MoS2. The latter resembles black lead
(graphite), and its appearance suggested the name of the element
(Gk. /u,oAu)8£uva, lead). The molybdenite is converted by roasting into
molybdic . anhydride Mo08. When this is treated with ammonium
hydroxide, or with sodium hydroxide, ammonium molybdate (NH4)2
Mo04, or sodium molybdate Na2Mo04, 10H20 is obtained. The metal
itself is liberated by reducing the oxide or chloride with hydrogen.
When pure it resembles wrought iron, and, like iron (q.v.), takes up
carbon and shows the phenomena of tempering. The oxides [MoO ?],
MOjjOg, Mo02, and MoO8 are known, but the lower oxides are not basic.
The chlorides MogClg, MoCl^ MoCl4, and MoCl5 have been made. The
chief use of molybdenum compounds in the laboratory is in testing for
and estimating phosphoric acid. . When a little of a phosphate is added
to a solution of ammonium molybdate in nitric acid, and the mixture
is warmed, a copious yellow precipitate of a phosphomolybdate of
ammonium (NH4)3P04,llMo08, 6H20 is formed. The compound is
soluble in excess of phosphoric acid and in alkalies, but not in dilute
mineral acids.
Tungsten. — The minerals scheelite CaW04 and wolfram FeW04
are tungstates of calcium and iron respectively. By fusion of wolfram
with sodium carbonate and extraction with water, sodium tungstate
Na2W04, 2H20 is secured. It is used as a mordant and for rendering
muslin fire-proof. Acids precipitate tungstic acid H,W04,H20 from
solutions of this salt. The element gives the oxides W02 and W03,
the latter being formed by ignition of tungstic acid. The chlorides
WC12, WC14, WC15, and WC16 are known, the last being formed
directly, and the others by reduction. A hard variety of steel contains
5 per cent of tungsten.
Uranium. — This element is found chiefly in pitchblende, which
contains the oxide U808 along with smaller amounts of many other ele-
THE CHROMIUM FAMILY. RADIUM 733
rnents. By roasting the ore with carbonate and nitrate of sodium, and
extracting with water, an impure solution of sodium uranate NaU04 is
obtained. Acids precipitate the insoluble, yellow diuranate Na2U207,
6H20. This salt is used in making uranium glass, which shows a
yellowish-green fluorescence. The property is due to the fact that
the wave-length of part of the invisible, ultra-violet rays of the sun-
light are shortened, and a> greenish light is therefore in excess. The
oxides are U02 a basic oxide, XJ208, U808 the most stable oxide,
U08 uranic anhydride, and U04 a peroxide.
When the oxide U02 is treated with acids, it gives uranous salts
such as uranous sulphate U(S04)2,'4H20. Uranic anhydride and
uranic acid interact with acids, giving basic salts, such as U02S04,
3^H2O, and U02(N03)2, 6H20, which are named uranyl sulphate, uranyl
nitrate, and so forth. They are yellow in color, with green fluores-
cence! Ammonium sulphide throws down the brown, unstable uranyl
sulphide U02S from their solutions.
RADIUM.
The Discovery of the Element.* — It was Becquerel who first
noticed (1898) that all compounds of uranium gave out a radiation
capable of affecting a photographic plate covered with black, light-
proof paper. The Becquerel rays, however, required several days to
produce a distinct effect. These rays had a second, equally remarkable
property. Ordinary air is an extremely poor conductor of electricity,
and for this reason a well-insulated, electrically charged body, such as
an electroscope, will retain its charge for a long time. Yet a few
tenths of a gram of any uranium compound, brought within 3 or 4 cm.
of the charged body, rendered the air a conductor, and the charge was
quickly lost. The air, under these conditions, is said to be " ionized,"
but the positive and negative ions it contains are probably large
molecular complexes. This property makes possible the quantitative
measurement of radio-activity, or the rate of production of Becquerel
rays. We simply have to compare the times required for the discharge
of an electroscope by different specimens of radio-active matter.
The radio-activity of every pure uranium compound is proportional
to its uranium content. The ores are, however, relatively four times
as active. This fact led M. and Mme. Curie to the discovery that the
* I am indebted to my colleague Professor H. N. McCoy for the material of
which, the followiug is a slightly condensed version.
734 INORGANIC CHEMISTRY
pitchblende residues, from which practically all of the uranium had
been extracted, were nevertheless quite active. About a ton of the
very complex residues having been separated laboriously into the con-
stituents, it was found that a large part of the radio-activity remained
with the compound of barium. The barium chloride, after being puri-
fied until no other elements could be detected in it by ordinary chemi-
cal tests, was sixty times as active as uranium. The percentage of
chlorine contained in it was almost identical with that in the ordinary
salt, and it differed from this only in its photo-active and " ionizing "
powers. By repeated systematic recrystallization, however, a portion
was separated which gave distinctly the spectrum of a new element and
a diminished percentage of chlorine. Finally, a product free from
barium, and three million times as active as uranium, was secured. The
nature of the spectrum and the chemical relations of the element, now
named radium, placed it with the metals of the alkaline earths. The
ratio by weight of chlorine to radium in the compound was 35.45 : 112.5,
so that, on the assumption that the element is bivalent, its atomic
weight is 225. With this value it occupies a place formerly vacant
in the periodic table.
Properties of Radium Compounds. — Radium has not been
isolated, and the chloride and bromide are usually employed. The
enormous photo-activity and ionizing power of the compound has
been mentioned above. The rays, like X-rays, may cause severe
"burns." Many substances, like zinc-blende and the diamond, phos-
phoresce brilliantly when exposed to the rays. The most remarkable
property of the salts, however, is their constant evolution of heat in
relatively enormous quantities. One gram of the element, in com-
bination, evolves over 100 cal. per hour, and it is estimated that the
total amount of heat spontaneously produced by 1 g. would be about
1010 cal. To produce the same amount of heat by combustion, no less
than 300 kg. of hydrogen would have to be burned.
The Radiations of Radium Salts. — The Becquerel rays are
made up of three different radiations : The a-rays, which produce the
ionization, and are almost completely absorbed by a piece of paper ;
the /3-rays, which produce the photographic effects and readily penetrate
paper and even thin sheets of metal; the y-rays, which resemble
X-rays.
When ammonium carbonate is added in excess to a solution of a
THE CHROMIUM FAMILY. RADIUM 735
soluble salt of uranium, a precipitate is formed and redissolves. A
trifling undissolved residue, called uranium-X, however, possesses
itself of all the /3-ray activity of the original material. While the
activity of any ordinary uranium compound is constant, that of U-X
decreases rather rapidly, reaching half value in 22 days, and disappear-
ing entirely in a few months.
The uranium freed from U-X has at first only o-ray activity, but
after 22 days it has recovered half its former /3-ray activity, and
after a few months the whole. The removal of equal quantities of
U-X may be repeated, in this way, at intervals of a few months, any
number of times. There is thus a continuous production of* U-X from
uranium.
It has been found that the /3-rays are identical with cathode rays,
and, in terms of the molecular hypothesis, consist of minute particles,
electrons or corpuscles, shot out with a velocity approaching that of
light. The mass of each corpuscle is about T(^(T °^ that °^ an at°m
of hydrogen (cf. p. 222), and bears a negative charge of electricity
equal to that on a chlorine ion in solution.
The a-rays consist of particles about twice as heavy as a hydrogen
atom, and move with enormous velocity. Each of these particles bears
a positive charge equal to that on one atom of hydrion.
Rutherford and Soddy account for the facts on which the above
hypothetical statements are based by the disintegration hypothesis.
All atoms are considered as groups of minute particles in a state of
rapid orbital motion. The atoms ef radio-active elements are not
perfectly stable, and, occasionally, disintegration occurs, one or more of
the particles being shot out. These particles constitute the a- and
/2-rays. The heating effect of radium salts is thus due to the change
of the kinetic energy of the moving particles into heat when they
encounter some other body.
The y-rays probably originate by the impact of /3-rays, just as
X-rays are produced by cathode rays.
The Decay of an Element. — Uranium (at. wt. 238.5) has the
heaviest known atom. The disintegration of its atom gives a-rays and
leaves the lighter atoms of U-X. In a similar manner, radium gives
rise to a new radio-active, gaseous body, the radium emanation. This,
like U-X, is not permanent, and loses half its activity in four days. In
doing so it produces a series of new radio-active substances. These,
with their times of decay to half value, are as follows (Rutherford) :
736 INORGANIC CHEMISTRY
Ra^Em— >A-+B^C— >D->E-> F
1500 yrs. 4 days 3 min. 26 min. 19 min. 40 yrs. 6 days 142 days
Radium thus loses its activity. To account for its presence in the ore,
we must, therefore, suppose that it is being continuously produced from
some source. This source must be uranium, for every known ura-
nium ore contains radium (McCoy) and radium emanation (Boltwood)
in amounts proportional to the uranium content. Furthermore, it is
asserted by Soddy that the radium emanation is slowly re-formed in a
uranium compound previously freed from radio-active substances,
although Boltwood has failed to confirm this observation.
It has recently been found (Ramsay and Soddy) that helium (p. 430)
is one of the decomposition products of the radium emanation.
The phenomena of radio-active substances lead undeniably to the
startling conclusion that some, if not all, of the elements are capable
of spontaneous decomposition. The transmutation of the elements, in
the manner indicated, seems to be clearly established, since it appears
certain that the element uranium, produces another undoubted element,
radium, which, in turn, yields the element helium. Besides the U — Ra
series of radio-active substances, others are known. Thus, all of the
compounds of thorium are radio-active, and other less well-character-
ized substances, such as polonium, actinium, radio-tellurium, radio-lead,
etc., have been separated from pitchblende. It is probable that radio-
lead is radium-D, and that radio-tellurium is radium-F. There is also
considerable evidence that many other elements are very slightly radio-
active, perhaps TTTV(r as active as uranium. This may mean that all
elements undergo an extremely slow, spontaneous decay. At present
there is no known means of hastening the rate of radio-active change.
Exercises. — 1. Construct equations showing the interactions of
(a) chromic oxide and aluminium, (b) strontium nitrate and potassium
dichromate in solution, (c) potassium hydroxide and chromic hydroxide,
and the reversal on boiling, (d) chlorine and potassium chromite in
excess of alkali (what is the actual oxidizing agent ?).
2. What volume of oxygen at 0° and 760 mm. (a) is obtainable
from one formula-weight of potassium dichromate (par. 4, p. 726), (b)
is required to oxidize one formula- weight of chromous chloride ?
3. To what classes of actions should you assign the three methods
of making chromic oxide (p. 730) ?
CHAPTER XLIII
MANGANESE
The Chemical Relations of the Element. — Manganese stands,
at present, alone on the left side of the eighth column of the periodic
table. The right side is occupied by the halogens. It is never uni-
valent, as the halogens are, but its heptoxide Mn207 and the corre-
sponding acid, permanganic acid HMnO4, are in many ways closely
related to the heptoxide of chlorine and perchloric acid HC104. Of
the lower oxides of manganese, MnO is basic, and MiLjOg feebly basic.
Mn02 is feebly acidic, Mn03 more strongly so, and permanganic acid
(from Mn207) is a very active acid. Contrary to the habit of feebly
acidic and feebly basic oxides, such as those of zinc, aluminium, and
tin, the basic oxides of manganese are not at all acidic, and the acidic
oxides (with the possible exception of Mr^Og) are not also basic. There
are thus the five following, rather well-defined sets of compounds,
showing five different valences of the element. Of these the first,
fourth, and fifth are the most stable and the most important.
1. Manganous compounds, MnO, Mn(OH)2r MnS04, etc. These
compounds resemble those of the magnesium family (and those of
Fe**). The salts of weak acids, such as the carbonate and sulphide,
are easily made, and there is little hydrolysis of the halides. The
salts are pale-pink in color.
2. Manganic compounds, Mn203, Mn(OH)8, Mn2(S04)8, [MnCl8].
The salts resemble the chromic and aluminium salts in behavior, but
are even less stable than those of quadrivalent lead. They are com-
pletely hydrolyzed by little water. The salts are violet in color.
3. Manganites, Mn02, H2Mn08, CaMn03. The alkali manganites
are strongly hydrolyzed, like plumbates and stannates.
4. Manganates, MnO8, H2Mn04, K2Mn04. The salts resemble the
sulphates and chromates, but are much more easily hydrolyzed. The
free acid resembles chloric acid in that when it decomposes it yields
a higher acid (HMn04) and a lower oxide (Mn02). The salts are
green in color.
5. Permanganates, Mn207, HMn04(hydrated), KMn04. The salts
737
738 INORGANIC CHEMISTRY
resemble the perchlorates. and are not hydrolyzed by water. They are
reddish-purple in color.
It will be seen that the element manganese changes its character
totally with change in valence, and in each form of combination
resembles some set of elements of valence identical with that which it
has itself assumed.
Occurrence and Isolation. — The chief ore is the dioxide, pyro-
lusite Mn02, which always contains compounds of iron. Other man-
ganese minerals are : braunite Mn203 ; the hydrated form, manganite
MnO(OH) ; hausmannite Mn304; and manganese spar MnC03. The
last is isomorphous with calcite. The metal is most easily made by
reducing one of the oxides with aluminium by Goldschmidt's method.
Physical and Chemical Properties. — The metal manganese
has a grayish luster faintly tinged with red. It rusts in moist air, and
easily displaces hydrogen from dilute acids, giving manganous salts.
Its alloys with iron, such as ferro-manganese (20-80 per cent mangan-
ese), are used in the arts.
Oxides. — Manganous oxide MnO is a green powder, made by re-
ducing any of the other oxides with hydrogen. Hausmannite Mn304
is red. An oxide having this composition is formed when any of the
other oxides is heated in air, oxidation or reduction, as the case may
be, taking place (cf. p. 701). This oxide corresponds to minium
Pb304 (p. 701) rather than to Fe3O4, for with dilute acids it gives a
soluble manganous salt and a precipitate of the dioxide :
Mn2Mn04 + 4HN03 -» 2Mn(N03)2 + H4Mn04|.
The hydrated dioxide H4Mn04 subsequently loses water. Haus-
mannite also forms square prismatic crystals. In view of its behavior
with acids and its crystalline form, it is thought to be an orthoman-
ganite of manganese Mn2Mn04, rather than a derivative of manganic
oxide, Mn(MnO2)2, which would be a spinelle (p. 686). The magnetic
oxide of iron Fe (Fe02)2 belongs to the regular system, like the spinelles.
Manganic oxide Mn203 is brownish-black, and is formed by heating any
of the oxides in oxygen. In dilute acids it behaves as if it were a
manganite of manganese Mn.Mn03, for it gives a manganous salt and
manganese dioxide. Yet compounds of trivalent manganese are known,
and this may be one.
MANGANESE 739
Manganese dioxide Mn02 is black, and is most easily prepared in
pure condition by gentle ignition of manganous nitrate. The hydrated
forms of the oxide are produced by reactions like those just mentioned,
and by adding a hypochlorite or hypobromite to manganous hydroxide
suspended in water. Manganese dioxide is not a peroxide in the re-
stricted sense (cf. p. 308). It is used for manufacturing chlorine,
although electrolytic processes are now driving it out of this field. In
glass-making (q.v.'), it is employed to oxidize the green ferrous silicate,
derived from impurities in the sand, to the pale-yellow ferric com-
pound. The amethyst color of the manganous silicate which is formed
tends to neutralize this yellow. The dioxide forms the depolarizer in
the Leclanche cell (p. 673).
Manganese trioxide is a red, unstable powder. Manganese hept-
oxide is a brownish-green oil (see below).
When any of these oxides is heated with an acid, a manganous salt
is obtained. Salts of this class are, in fact, the only stable substances
in which manganese is combined with an acid radical. In this action
the oxides containing more oxygen than does MnO give off oxygen, or
oxidize the acid (cf. p. 172). When the oxides are heated with bases,
in the presence of air, manganates are always formed. With the
oxides containing a smaller proportion of oxygen than Mn08 oxygen
is taken from the air.
Manganous Compounds. — The manganous salts are formed by
the action of acids upon the carbonate or any of the oxides. Thus the
chloride MnCl^ 4H20 is obtained in pale-pink crystals from a solution
made by treating the dioxide with hydrochloric acid and driving off
the chlorine liberated by oxidation (p. 171). The hydroxide Mn(OH)2
is formed as a white precipitate when a soluble base is added to a solu-
tion of a manganous salt. This body passes into solution when ammo-
nium salts are added, and cannot be precipitated in their presence on
account of the formation of molecular ammonium hydroxide and the
suppression of hydroxidion (cf. magnesium hydroxide, p. 644). The
hydroxide quickly darkens when exposed to the air and passes over
into hydrated manganic oxide MnO (OH).
Manganous sulphate gives pink crystals of a hydrate. Below 6°
the solution deposits MnS04, 7H20, which is a vitriol (p. 649). Be-
tween 7° and 20° the product is MnSO4, 5H20, asymmetric and isomor-
phous with CuS04, 5H20. Above 25° monosymmetric prisms of
MnSO4, 4H2O are obtained. These hydrates have different aqueous
740 INORGANIC CHEMISTRY
tensions and may be formed from one another by lowering or raising
the pressure of water vapor around the substance (p. 122). The signi-
ficance of the temperatures proper to the crystallization of each (cf. pp.
573, 608, 623) is that a given solid hydrate can be formed only in a solu-
tion which is saturated with respect to that hydrate and has the same
aqueous tension as the hydrate. These conditions are necessary to that
state of equilibrium between the solution and the hydrate on which the
co-existence of solution and hydrate during crystallization depends (cf.
p. 160). Hence the hydrates with the larger proportions of water, and
the higher aqueous tensions, are formed in the colder solutions which
contain less of the solute when saturated and have therefore at a given
temperature themselves relatively high aqueous tensions.
The presence of a foreign dissolved body, since it will lower the
vapor tension of the solution, may similarly cause the formation of a
lower hydrate. Thus, at the ordinary temperature, calcium sulphate solu-
tion has a higher aqueous tension than gypsum, and therefore gypsum
is deposited from it, and anhydrite will turn into gypsum if placed in it.
But calcium sulphate solution containing much of the chlorides of
sodium and magnesium has a lower aqueous tension than gypsum, and
so anhydrite is deposited, and gypsum in contact with such a solution
would lose its water of hydration. This explains the deposition of
anhydrite in the salt layers (cf. p. 603).
Manganous carbonate MnC03 is a white powder formed by pre-
cipitation. The sulphide MnS is obtained as a green powder by lead-
ing hydrogen sulphide over any of the oxides. A flesh-colored,
hydrated form MnS, H2O is more familiar and is precipitated by am-
monium sulphide from manganous salts. It interacts with mineral
acids and even with acetic acid, so that it cannot be precipitated by
hydrogen sulphide (cf. p. 651).
The manganous salts of weak acids, such as the carbonate and
sulphide, darken when exposed to 'air and are oxidized, with formation
of hydrated manganic oxide. As Ave have seen, manganous hydroxide
is similarly oxidized and these salts are precisely the ones which
should furnish the hydroxide by hydrolysis. While there is a general
resemblance between the manganous salts and the stannous, chromous,
and ferrous salts, the manganous salts of active acids are not oxidized
by the air as are the corresponding salts of the other three metals.
Manganic Compounds. — The base of this set of compounds,
manganic hydroxide Mn (OH)3, is slowly deposited by the action of
MANGANESE 741
the air on an ammoniacal solution of a manganous salt in salts of
ammonium. The chloride MnCl3 is present in the liquid obtained by
the action of hydrochloric acid upon manganese dioxide (cf. p. 171), but
loses chlorine very readily and cannot be isolated. Double salts such as
MnCl3, 2KC1 and MnF3, 2KF; 2H20 are known. Manganic sulphate
Mn2(S04)3 is deposited as a violet-red powder when hydrated man-
ganese dioxide is heated with concentrated sulphuric acid at 160°.
It is deliquescent and is rapidly hydrolyzed in the cold even by a
little water, giving the brownish-black hydroxide :
Mn2(S04)8 + 6H20 -> 2Mn(OH)3 + 3H2S04.
The caesium-manganic alum Cs2S04, Mn2(S04)8, 24H20 seems to be the
most stable derivative.
Manganites, — Although manganese dioxide interacts when fused
with potassium hydroxide, simple salts derived from H2Mn03 ( = H20,
Mn02) or H4Mn04 (= 2H20,Mn02) are not formed. The products are
complex, as K2Mn5On. Some less complex manganites are formed in
the Weldon process for utilizing the manganous chloride obtained in
manufacturing chlorine. The liquor is mixed with slaked lime, and
air is blown through the mass of calcium and manganous hydroxides
which is thus obtained. Black manganites of calcium, such as
CaMn03(= CaO,Mn02) and CaMn205(CaO,2MnO2) are thus formed:
Ca(OH)2 + 2Mn(OH)2 + 02 -> CaMn205 + 3H20,
and when afterwards treated with hydrochloric acid they behave like
mixtures of manganese dioxide and calcium oxide. As we have seen
(p. 738), the oxides Mn304 and Mn203 may be manganites of manga-
nese.
Manganates. — When one of the oxides of manganese is fused
with potassium carbonate and potassium nitrate a green mass is
obtained. The green aqueous extract deposits potassium manganate
K2Mn04 in rhombic crystals, which are isomorphous with those of potas-
sium sulphate, and are almost black :
K2C03 + Mn02 + 0 -» K2Mn04 + C02.
The acid H2MnO4, itself unknown, must be weak, for the potassium
salt is easily hydrolyzed. The salt remains unchanged in solution
only in presence of free alkali, the hydroxidion of the alkali combin-
742 INORGANIC CHEMISTRY
ing with and suppressing the hydrion of the water whose combination
with the Mn04" ion constitutes the hydrolysis. When the concentra-
tion of the hydroxidion is reduced by dilution, or, better still, when a
weak acid such as carbonic acid or acetic acid is used to neutralize it,
the salt is hydrolyzed, according to the partial equation :
K2Mn04 -f- 2H20 -+ 2KOH ( + H2Mn04). (1)
The free acid immediately changes so that a part is oxidized to per-
manganic acid, giving a purple-red color to the solution, and a part is
reduced to manganese dioxide, giving a black precipitate. The trans-
formation is similar to that of chloric acid (p. 275). The equation
may be made by noting that manganic acid has the composition H20,
Mn03 and changes so as to yield H20,Mn207 and MnO2. Thus each
molecule of H2Mn04, in forming a molecule of Mu02, yields one unit
of oxygen, while 2(H20,MnG3) -f- 0 are required to give H20,Mn207 +
H20:
3(H20,Mn03) -+ H20, Mn207 + Mn02 + 2H20
or (3H2Mn04) -» 2HMn04 + Mn02 + 2H20. (2)
In consequence of the presence of potassium hydroxide (equation (1))
the product is potassium permanganate :
2KOH + 2HMn04 -> 2KMn04 + 2H20. (3)
Multiplying equation (1) by 3, omitting the manganic acid, and add-
ing the three partial equations, we have the equation for the action as
it really occurs :
2H20 -» 4KOH + 2KMn04 + Mn02.
In terms of the ions the equation is simpler :
3Mn04" + 2H* -+ 20H' + 2MnO/ + MnO2.
The alkaline solution of potassium manganate intei ,ts readily
with oxidizable substances. Thus oxalic acid is converted into car-
bonic acid, and alcohol into acetic acid. The details of the change
depend upon the amount of free alkali present and the nature of the
product of oxidation. Lower oxides of manganese such as Mn02 are
usually precipitated.
Permanganates. — Potassium permanganate KMnO4 is made by
hydrolysis of the manganate as shown above, and is obtained, in purple
MANGANESE 743
crystals with a greenish luster, by evaporation of the solution. The
crystals are rhombic prisms, isomorphous with potassium perchlorate.
To avoid the loss of manganese thrown down as dioxide, the action is
carried out commercially by passing ozone through the solution of the
manganate :
2K2Mn04 + 03 + H20 -» 2KMn04 + 02 + 2KOH.
Sodium permanganate is made in a similar manner. It is not obtain-
able in solkl form, but its solution is known as " Condy's disinfecting
fluid." This liquid owes its properties to the oxidizing power of the
salt. Permanganic acid is a very active acid, that is, it is highly ion-
ized in aqueous solution. A solid hydrate of the acid may be secured
in reddish-brown crystals by adding sulphuric acid to a solution of
barium permanganate and allowing the nitrate to evaporate :
Ba(MnO4)2 + H2S04 + zH2O ^± BaSO4 J. -f 2HMnO4, xH2Q.
This hydrate decomposes, on being warmed to 32°, and yields oxygen
and manganese dioxide. When a very little dry, powdered potassium
permanganate is moistened with concentrated sulphuric acid, brownish-
green, oily drops of permanganic anhydride (manganese heptoxide)
Mn2O7 are formed. This compound is volatile, giving a violet
vapor, and is apt to decompose explosively into oxygen and manganese
dioxide. Its oxidizing power is such that combustibles like paper,
ether, and illuminating-gas are set on fire by contact with it.
Potassium permanganate is much used for oxidations. The actions
are different according as the substance is employed (1) in alkaline,
(2) in acid, or (3) in neutral solution.
1. When an alkali, such as potassium hydroxide, is added, the
action by which the permanganate is formed is reversed, and the solu-
tion becomes green from the production of the manganate :
'J4KMn04 -f 4KOH -» 4K2Mn04 + 2H2O + O2,
or 4Mn04' + 40H' -> 4Mn04" + 2H20 + O2.
When a substance capable of being oxidized is present, the reduction
proceeds further and manganese dioxide is precipitated. Schemati-
cally : Mn2O7 — > 2Mn02 -+- 30, so that two molecules of the perman-
ganate, in alkaline solution, can furnish three chemical units of
oxygen to the oxidizable body.
744 INORGANIC CHEMISTRY
2. With an acid, the amount of oxygen available is greater, for the
manganous salt of the acid is formed :
Mn207 -> 2MnO + 50.
Thus, when sulphuric acid is added to potassium permanganate solu-
tion, and sulphur dioxide is led through the mixture, we have :
2KMn04 + 3H2S04 -^ K2S04 + 2MnS04 + 3H20(+ 50) (1)
(5Q)+5H2S08 -> 5H2S04 (2)
2KMnO4 + 3H2S04+5H2S08 -> K2S04 + 2MnS04 + 3H20 + 5H2SO4
In this case, since sulphuric acid is a product, the preliminary addi-
tion of the acid was superfluous. In other cases, the partial equation
(1), showing the available 50, remains the same, while the other par-
tial equation varies with the substance being oxidized. Thus, with
hydrogen sulphide as reducing agent, we have :
(0) + H2S -> H/) + S x 5 (2')
and with ferrous sulphate, we get ferric sulphate :
2FeS04 + H2S04( + 0) -> Fe2(S04)3 + H20 x 5 (2")
As before (2') and (2") must be multiplied throughout by five, before
summation is made. Since the permanganate is deep-purple in color,
while the manganous salt is almost colorless, this sort of action can be
used without an indicator for quantitative experiments. The stan-
dard solution of the permanganate is added from a burette until the
purple color ceases to disappear ; and the amount used enables us to
calculate the quantity of ferrous salt, oxalic acid, nitrous acid, or other
oxidizable substance, which was present. The last named substance
is oxidized to nitric acid.
3. AVhen dry potassium permanganate is heated, it decomposes
as follows :
2KMn04 -» K2Mn04 + Mn02 +- 02.
The neutral solution resembles that of potassium dichromate in oxi-
dizing substances which are reducing agents, but is more active.
Thus when the powdered salt is moistened with glycerine, the mass
presently bursts into flame. The fingers are stained brown, receiving
a deposit of manganese dioxide, in consequence of the reducing power
of the unstable organic substances in the skin. The destruction
of minute organisms by Condy's fluid results from a similar action.
MANGANESE 745
Analytical Reactions of Manganese Compounds. — The ions
commonly encountered are dimanganion Mn *, which is very pale-pink
in color, permangananion MnO/, which is purple, and mangananion
MnO/', which is green.
The manganous compounds give with ammonium sulphide the
flesh-colored, hydrated sulphide which is soluble in acids. Bases give
the white hydroxide, which darkens by oxidation, and is soluble in
salts of ammonium. The black, hydrated dioxide is precipitated by
hypochlorites.
All compounds of manganese confer upon the borax bead an ame-
thyst color which, in the reducing flame, disappears. A bead of
sodium carbonate and niter becomes green on account of the formation
of the manganate.
Exercises. — 1. Consider the valence of manganese in the oxides
Mn8O4 and Mn208, on the theory that they are manganites.
2. What do we mean by saying that (a) chromous chloride is stable
(p. 119), but easily oxidized by the air, (i) permanganic acid is an
active acid, (c) permanganic acid is an active oxidizing agent in pres-
ence of an acid ?
3. Formulate the oxidations of oxalic acid and of nitrous acid by
potassium permanganate in acid solution.
CHAPTER XLIV
IRON, COBALT, NICKEL .
THE elements iron (Fe, at. wt. 55.9), cobalt (Co, at. wt. 59), and
nickel (Ni, at. wt. 58.7) are not corresponding members of successive
periods, like the families hitherto considered. They are neighboring
members of the first long period, lying between its first and second
octaves (p. 308), and form a transition group between the adjoining
elements within those octaves. Thus, iron forms ferrates M2IFeVIO4
and ferric salts FemCl8, as well as ferrous salts Fe^CLj. These
resemble the chromates and manganates, the chromic and manganic
salts, and the chromous and manganous salts, respectively. Cobalt
forms cobaltic and cobaltous salts, like Co2m(S04)3 and Co1 C12. Nickel
enters only into nickelous salts, like MC12, and thus links iron and
cobalt with copper and zinc which are both bivalent elements. The
free metals of this family are magnetic, iron showing this property
strongly and cobalt very distinctly.
IRON.
Chemical Relations of the Element. — The oxides and hydrox-
ides FeO and Fe(OH)2, Fe203 and Fe(OH)3 are basic, the former more
strongly so than the latter. The ferrous salts, derived from Fe(OH)2,
resemble those of the magnesium group and those of Cr" and Mn"* and
are little hydrolyzed. The ferric salts, derived from Fe(OH)3, re-
semble those of Cr*" and Al"" and are hydrolyzed to a considerable
extent. Ferric hydroxide is even less acidic, however, than is chromic
hydroxide. Iron gives also a few ferrates K2Fe04, CaFe04, etc., de-
rived from an acid H2FeO4 which, like manganic acid H2Mn04 (p. 742),
is too unstable to be isolated. Complex anions containing this element,
such as the anion of K4.Fe(CN)6, are familiar, but complex cations
containing ammonia are unknown.
The ferrous salts differ from most of the manganous salts and re-
semble the chromous and stannous salts in being easily (although not
quite so easily) oxidized by the air. They pass into the ferric
condition.
746
IRON, COBALT, NICKEL
747
Occurrence. — Free iron is found in minute particles in some
basalts, and many meteorites are composed of it. Meteoric iron can
be distinguished from specimens of terrestrial origin by the fact that
it contains 3-8 per cent of nickel. The chief ores of iron are the
oxides, haematite Fe2O3 and magnetite Fe804, and the carbonate FeC03,
siderite. The first is reddish and columnar in structure ; but black,
shining, rhombohedral crystals, known as specularite, are also found.
Hydrated forms, like brown iron ore 2Fe2Og, 3H20, are also common.
Siderite is pale-brown in color and rhombohedral, isomorphous with
calcite. When mixed with clay it forms iron-stone and, with 20-25
per cent of coal in addition, black-band. Pyrite FeS, consists of
golden-yellow, shining cubes or pentagonal dodecahedra. It is used,
on account of its sulphur, in the manufacture of sulphuric acid, but,
from the oxidized residue, iron of sufficient purity is obtained with
difficulty. Compounds of iron are con-
tained in chlorophyl and in the blood
(haemoglobin), and doubtless play an
important part in connection with the
vital functions of these substances. By
interaction with organic compounds of
iron present in the tissues, ammonium
sulphide blackens the skin, ferrous sul-
phide being formed.
Metallurgy. — The ores of iron are
usually first roasted in order to decom-
pose carbonates and oxidize sulphides.
They are then reduced with coke. Ores
containing lime or magnesia are mixed
with an acid flux, such as sand or clay-
slate, in order that a fusible slag may
be formed. Conversely, ores containing
silica and clay are mixed with lime-
stone. With proper adjustment of the
ingredients the process can be carried
on continuously in a blast furnace (Fig. FIG. 105.
105). The solid materials thrown in
at the top are converted, as they slowly descend, completely into
gases which escape and liquids (iron and slag) which are tapped
off at the bottom. Heated air is blown in at the bottom through
748 INORGANIC CHEMISTRY
tuyeres, and the top is closed by a bell which descends for a moment
when an addition is made to the charge. The gases, which contain
much carbon monoxide, are led off and used to heat the blast or to
drive gas-engines.
The main action takes place between the carbon monoxide, present
in consequence of the excess of carbon, and the oxide of iron :
Fe304 + 4CO +± 3Fe + 4C02.
Since the action is a reversible one, a large excess of carbon monoxide
is required. Different smelters use different proportions, the ratio of
the amount actually used to that supplied varying from 1 : 2 to 1 : 15.
The actual ratio by volume of carbon dioxide to carbon monoxide
required for equilibrium with the two solids is, 61 : 39 at 650°, and
7 : 93 at 800°.
In the upper part of the furnace, the heat (400°) loosens the texture
of the ore. Further down, the temperature is higher (500-900°),
and the carbon monoxide reduces the oxide of iron to particles of soft
iron. A temperature high enough to melt pure iron is barely reached
anywhere in the furnace, but, a little lower down, by union with carbon,
the more fusible cast iron (1200°) is formed and falls in drops to the
bottom. It is in this region also that the slag is produced. If the flux
had begun sooner to interact with the unreduced ore, iron would have
been lost by the formation of the silicate. The iron collects below
the slag, and the latter flows continuously from a small hole. The
former is tapped off at intervals of six hours or so from a lower
opening.
Cast Iron and Wrought Iron. — Pure iron is not manufactured,
and indeed would be too soft for most purposes. Piano-wire, however,
is about 99.7 per cent pure. The product obtained from the blast fur-
nace contains about 90 per cent of iron along with 3-4.5 per cent of car-
bon, often nearly as much silicon, varying proportions of manganesef
and some phosphorus and sulphur. The last four ingredients are lib-
erated from combination with oxygen by the carbon in the hottest part
of the furnace and combine or alloy themselves with the iron. Cast
iron does not soften before melting, as does the purer wrought iron, but
melts sharply at 1150-1250° according to the amount of foreign material
it contains. When suddenly cooled it gives -white cast iron which is
very brittle and looks homogeneous to the eye, all the carbon being
present in the form of carbide of iron Fe8C in solid solution in the
IRON, COBALT, NICKEL
749
metal. By slower cooling, time is permitted for the separation of part
of the carbon as graphite and for other changes (see below), and gray
cast iron results. Spiegel iron is cast iron made from ores containing
5-20 per cent of manganese and the usual proportion of carbon. Ferro-
manganese contains 20-80 per cent of the same element.
"Wrought iron is made by heating the broken " pigs " of cast iron
upon a layer of material containing oxide of iron and hammer-slag
(basic silicate of iron) spread on the bed of a reverberatory furnace.
The carbon, silicon, and phosphorus combine with the oxygen of the
oxide, and the last two pass into the slag. The sulphur is found in the
slag as ferrous sulphide. On account of the effervescence due to
the escape of carbon monoxide, the
process is called "pig-boiling." The
iron is stirred with iron rods (" pud-
dled") and stiffens as it becomes
purer, until finally it can be with-
drawn in balls (" blooms ") and freed
from slag under the steam-hammer.
It now softens sufficiently for weld-
ing below 1000° and melts at 1600°
or higher, according to its purity.
If it still contains more than a trace
of combined phosphorus it is brittle when cold (" cold short "). A
little surviving sulphide of iron makes it brittle when hot ("red-
short ") and unsuitable for forging. Wrought iron should contain
only 0.1-0.2 per cent of carbon. The above operations are now largely
performed by machinery.
Steel. — This is a variety of iron almost free from phosphorus,
sulphur, and silicon, and containing usually not more than 1.5 per cent
of carbon. "Mild steel" contains 0.2-0.5 per cent of carbon. Steel
combines the properties of wrought and of cast iron, being hard and
elastic, and at the same time available for forging and welding when
the proportion of carbon is not too high.
Steel is made largely by the Bessemer process. The molten cast
iron is poured into a converter (Fig. 106) and a blast of air (A) is
blown through it. The oxidation of the manganese, carbon, silicon,
sulphur, and a little of the iron gives out sufficient heat to raise the
temperature of the mass above the melting-point of wrought iron.
The required proportion of carbon is then introduced by adding pure
FIG. 106.
750 INORGANIC CHEMISTRY
cast iron, Spiegel iron, or charcoal, and the contents, first the slag, and
then the molten steel, are finally poured into molds by turning the
converter. When the cast iron contains much phosphorus, the oxide
of this element is reduced again by the iron as fast as it is formed by
the blast. In such cases a basic lining containing lime and magnesia
takes the place of the sand and clay lining of the ordinary Bessemer
converter, and a slag containing a basic phosphate of calcium is pro-
duced. This modification constitutes what is known as the Thomas-
Gilchrist process. The slag ("Thomas-slag") when pulverized forms
a valuable fertilizer (cf. p. 605).
In the Siemens-Martin, or open hearth process, the cast iron is
melted in a saucer-shaped depression lined with sand, and scraps of
iron plate (for dilution) and haematite, or some other oxide ore, are
then added in proper proportions. The materials are heated with gas
fuel for 8-10 hours until a sample shows, under the hammer, that the
process is complete. The product is then drawn off through a hole
and cast in molds.
Properties of Steel' — When steel is heated to redness and
cooled slowly, it is comparatively soft. Sudden chilling, however, ren-
ders it as hard as glass. By subsequent, cautious heating the hardness
may be reduced to any required extent, and this treatment is called
" tempering." The sufficiency of the heating is judged roughly by the
interference colors caused by the thin film of oxide which forms on the
surface. Thus a pale-yellow color (430-450°) serves for tempering
razors, a decided yellow (470°) for pen-knives, a brown (490-510°) for
shears, a purple (520°) for table-knives, a blue (530-570°) for watch-
springs and sword-blades, and a black-blue (610°) for saws. Except
in the case of watch-springs, these films are afterwards removed by the
grinding.
To understand this behavior it must be noted that there are three
states of solid iron resembling the rhombic and monoclinic states of
sulphur (cf. p. 368). The form stable below 765° is known as a-ferrite
(wrought iron). It is magnetic and can hold little carbide of iron in
solid solution. Above 765° this changes into /2-ferrite which, likewise,
holds little of the carbide in solution, but is not magnetic. At 890°
this changes into y-ferrite, a non-magnetic form in which the carbide is
soluble. When allowed to cool, iron assumes these forms in the reverse
order. If, now, a fluid solution of carbon in iron, suitable for steel, is
suddenly chilled, a great part of the cold mass is a supercooled solid
IRON, COBALT, NICKEL 751
solution of carbide of iron in y-f errite. This material is called marten-
site and is very hard and brittle. It is less stable at ordinary tem-
peratures than is a-ferrite, but, as is the case with yellow phosphorus
(p. 359) and amorphous sulphur (p. 370), the low temperature having
once been reached, transformation into the more stable form is there-
after exceedingly slow. The material is hard steel.
When the molten steel (solution of carbon in iron) is allowed to
cool so sloivly that equilibrium can be reached at every step, a compli-
cated series of changes ensues. First the mass solidifies (at or before
1130°) to a mixture of martensite (y-ferrite with carbide in solid solu-
tion up to 2 per cent) and graphite.* As the temperature now falls
very slowly, more graphite separates until, at 1000°, 1.8 per cent re-
mains in solution. From this point the dissolved carbide of iron
(cementite Fe8C containing 6.6 per cent of carbon) is separated. At
670° pure a-ferrite also begins to appear. The final result is a mechani-
cal mixture of a-ferrite (wrought iron), carbide of iron, and, if the
original amount of carbon was sufficiently large, graphite. These com-
ponents may be recognized by making a microscopic study of a polished
surface, and their formation may be followed by chilling the specimen
at any desired stage. The soft iron which predominates in the product
of slow cooling makes the whole soft. Heating to a high temperature
and sudden chilling gives the homogeneous solid solution of the carbide
in y-ferrite once more and restores the qualities characteristic of steel.
Moderated reheating (tempering) of the chilled mass results in more
or less partial accomplishment of the changes proper to slow cooling,
and consequently in a more or less close approach to the condition
which results from this.
The difference between the effect of rapid and slow cooling of cast
iron (p. 748) can now be made clear. Rapid cooling leads to the
omission of the intervening steps enumerated above and, if something
like 5 or 6 per cent of carbon is present, the material turns almost
completely into the carbide (cementite). This is white cast iron.
With slower cooling, much graphite separates, and the product, gray
cast iron, contains much less of the carbide and much more free iron.
The various changes which occur in cooling steel are retarded by
the presence of foreign substances, just as, with sulphur (p. 370),
foreign substances delay the change from S^ to S\ and permit the
supercooling of the former and its appearance in the form of amorphous
* When molten cast iron, containing 3-4.6 per cent of carbon, is cooled in this
fashion, the amount of graphite may be considerable.
752 INORGANIC CHEMISTRY
sulphur. Manganese, nickel, and other metals, in particular, greatly
reduce the facility with which y-ferrite passes into ft- and o-f errite at
890° and 765°. Thus iron with 12 per cent of manganese, when
chilled from a high temperature, contains only supercooled y-ferrite
and is non-magnetic. It has to be kept for hours (instead of a few
minutes) at a temperature below 765°, say 500-600°, before it goes
over into a-f errite. Manganese is thus a valuable constituent of steel be-
cause, by favoring the survival of the y-ferrite in which alone the carbon
is soluble, it permits the manufacture of a homogeneous steel contain-
ing an unusually large proportion of dissolved carbon, and allows
slower cooling without loss of temper.
Pure Iron. — The pure metal may be made by reducing the
purified oxalate in a stream of hydrogen or by Goldschmidt's method
(p. 540).
Chemical Properties. — When exposed to moist air, iron receives
a loosely adherent coating of rust (2Fe203,Fe(OH)3). There is still
uncertainty as to how the product is formed. It may result from
displacement of the hydrogen of carbonic acid, the oxygen assisting
(cf. p. 624), and subsequent hydrolysis of the carbonate and oxidation
of the ferrous hydroxide. The fact that alkalies prevent rusting favors
this view, for they should diminish the amount of hydrion. According
to another theory, water and iron are simultaneously oxidized :
Fe + 02 + H2O -» FeO + H202
and the hydrogen peroxide immediately combines with the ferrous
oxide :
2FeO + H202 -> Fe202(OH)2. (2)
The presence of the peroxide cannot be demonstrated in this case,
perhaps because it is used up very rapidly. In the rusting of zinc,
however, it is always found (cf. also p. 636).
Iron burns in oxygen and interacts with superheated steam, giving
Fe304. A superficial layer of this oxide adheres firmly and protects
the iron from the action of the air (Barff's process for prevention of
rusting).
Iron displaces hydrogen easily from dilute acids. Steel and cast
iron, which contain iron, its carbide, and graphite, give with cold dilute
acids almost pure hydrogen, and the carbide and graphite remain un-
attacked. More concentrated acids, however, particularly when warm,
IRON, COBALT, NICKEL 753
give off, along with hydrogen, hydrocarbons formed by interaction
with the carbide (p. 543). The odor of the gas is due to com-
pounds of sulphur and phosphorus. With dilute nitric acid, iron gives
ferrous nitrate and ammonium nitrate (cf. tin, p. 694) and with the
concentrated nitric acid ferric nitrate and oxides of nitrogen. It has
little action upon alkalies.
After being dipped in very concentrated nitric acid, iron becomes
passive (cf. chromium, p. 723), and no longer displaces hydrogen and
other elements lying below it in the electromotive series. A sharp
blow, however, produces a change which spreads over the surface from
the point struck, and the metal becomes active once more.
Ferrous Compounds. — Ferrous chloride is obtained as a pale-
green hydrate FeCl2, 4H20 by interaction of hydrochloric acid with
the metal or the carbonate. The anhydrous salt sublimes in colorless
crystals when hydrogen chloride is led over the heated metal. At a
high temperature the vapor of ferrous chloride has a density correspond-
ing to the simple formula FeC^, but at lower temperatures there is
much association (p. 242) and the formula approaches Fe2Cl4. In
solution the salt is oxidized by the air to a basic ferric chloride :
4Fe" + 02 + 2H20 -> 4Fe*" + 40H/
In presence of excess of the acid, normal ferric chloride is formed.
With nitric acid, ferric chloride and nitric oxide are produced (p. 442).
Ferrous hydroxide Fe(OH)2 is thrown down as a white precipitate,
but rapidly becomes dirty-green and finally brown, by oxidation. It
dissolves in solutions of salts of ammonium, being, like magnesium
hydroxide (p. 644), sufficiently soluble in water to require an appreci-
able concentration of OH' for its precipitation. The ammonium salts
convert this into molecular ammonium hydroxide. Ferrous oxide FeO
is black, and is formed by heating ferrous oxalate in absence of air.
It may be made also by cautious reduction of ferric oxide by hydrogen
(at about 300°), but is easily reduced further to the metal.
Ferrous carbonate is found in nature, and may be made in slightly
hydrolyzed form by precipitation. The precipitate is white but rapidly
darkens and finally becomes brown, the ferrous hydroxide produced by
hydrolysis being oxidized to the ferric condition. The salt interacts
with water containing carbonic acid after the manner of calcium car-
bonate (p. 482), and hence is found in solution in natural (chalybeate)
waters.
754 INORGANIC CHEMISTRY
Ferrous sulphide may be formed as a black, metallic-looking mass
by heating together the free elements. It is produced by precipita-
tion with ammonium sulphide, but not with hydrogen sulphide. It
interacts readily with dilute acids. The precipitated form is slowly
oxidized to ferrous sulphate by the air.
Ferrous sulphate is obtained by allowing pyrites to oxidize in the
air and leaching the residue :
2FeS2 + 702 + 2H20 -» 2FeS04 + 2H2S04.
The liquor is treated with scrap iron and the neutral solution evapo-
rated until a hydrate FeS04, 7H20, green vitriol, or " copperas," is
deposited. This substance forms green crystals belonging to the
monosymmetric system, but gives also mixed crystals in which it is
isomorphous with the rhombic vitriols (cf. Magnesium and Zinc sul-
phates, p. 649). The crystals are efflorescent, and become also brown
from oxidation to a basic ferric sulphate :
4FeS04 + 02 + 2H.P -» 4Fe(OH)S04.
With excess of sulphuric acid and air, or an oxidizing agent, such as
nitric acid, ferric sulphate is formed. The anhydrous salt forms a
molecular compound with nitric oxide, and the solution becomes brown
when this gas is led through it, an unstable complex ion, perhaps
FeNO*", being produced (cf. p. 443). The double salts of the form
(NH4)2S04,FeS04,6H2O (Mohr's salt) are not efflorescent, and in solid
form are less readily oxidized than is ferrous sulphate. The ferrous sul-
phate is used in dyeing and in making ink. The extract of nut-galls con-
tains tannic acid, HC14H909, which, with ferrous sulphate, gives ferrous
tannate. This salt is oxidized by the air to the ferric condition, and the
ferric compound is a fine, black precipitate which can be suspended in
a solution of gum-arabic. The resulting material is ink. The black
streaks seen below nail-heads in oak and other woods are due to the
formation of ferrous carbonate and its interaction with the tannic acid
in the wood.
Ferric Compounds. — By leading chlorine into a solution of
ferrous chloride, and evaporating until the proper proportion of water
alone remains, a yellow, deliquescent hydrate of ferric chloride, FeCl3)
6H20 is obtained. When this is heated still further, hydrolysis takes
place and the oxide remains. When chlorine is passed over heated
iron, the anhydrous salt sublimes in dark scales, which are red by
IRON, COBALT, NICKEL 755
transmitted light. At a high temperature the formula of the vapor is
FeCl3, but at lower temperatures, in consequence of association, the
density increases and the formula approaches Fe2Cl6. In solution,
the salt, like other ferric salts, can be reduced to the ferrous condition
by boiling with iron :
21V" + Ye -* 3Fe".
The same reduction is effected by hydrogen sulphide and by stannous
chloride (cf. Mercuric chloride, p. 655) :
2Fe"*+ S"-»2Fe** + S|.
2Fe"* + Sn" -» 2Fe"+ Sn"".
The last action shows that ferrous salts are less active reducing agents
than are the stannous salts. The ferric ion is almost colorless, the
yellow-brown color of solutions of ferric salts being due to the pres-
ence of ferric hydroxide produced by hydrolysis. The color deepens
when the solution is heated, and fades again very slowly, by reversal
of the action, when the cold solution is allowed to stand. On the
other hand, the hydrolysis may be reversed and the color may be almost
destroyed, particularly in the case of the nitrate, when excess of the
acid is added to the solution :
Fe(N08)s + 3H2O «^ Fe(OH)8 + 3HN03.
Ferric iodide is reduced by the hydriodic acid produced by its own
hydrolysis, and hence ferrous iodide does not unite with iodine to form
this compound. The case is similar to that of cupric iodide (p. 621).
Ferric hydroxide, Fe(OH)3, appears as a brown precipitate when a
base is added to a ferric salt. It does not interact with excess of the
alkali. In this form the substance is a hydrogele (p. 523) and dries to
the oxide without giving definite intermediate hydrated oxides. The
hydrates, Fe4Og(OH)6 (brown iron ore) and Fe2O(OH)4 (bog iron ore),
however, are found in nature. The hydroxide passes easily into colloidal
solution in a solution of ferric chloride, and by subsequent dialysis
through a piece of parchment (cf. p. 523) the salt can be separated, and
a pure aqueous solution of the hydroxide obtained. This solution is
red in color, shows no depression in the freezing-point, and is not an
electrolyte. It deposits the hydroxide as a brown precipitate when
ionogens are added to the solution.
Ferric oxide, Fe203, is sold as "rouge" and " Venetian red." It is
made from the ferrous sulphate obtained in cleaning iron ware which
756 INORGANIC CHEMISTRY
is to be tinned or galvanized, and in other ways in the arts. The salt
is allowed to oxidize, and the ferric hydroxide, thrown down by the
addition of lime, is calcined. A purer form is produced by dry distil-
lation of the basic ferric sulphate, an operation which used to be under-
taken on a large scale for making Nordhausen sulphuric acid (p. 388).
This oxide is not distinctly acidic, but by fusion with more basic ox-
ides, compounds like franklinite Zn(Fe02)2 may be formed. It is
reduced by hydrogen, at about 300° to ferrous oxide (which catches fire
spontaneously in the air) and at 700-800° to metallic iron.
Magnetic oxide of iron Fe304 or lodestone is found in nature, and
is formed by the action of air (hammer-scale), steam, or carbon dioxide
on iron. It forms octahedral crystals like the spinelles (p. 686), and is
assumed to be Fe(Fe02)2.
Ferric sulphide may be made by fusing together the free elements
but is not obtained by precipitation. Soluble sulphides first reduce the
ferric salt to the ferrous conditions, liberating sulphur,
. Fe2(S04)3 + (NH4)2S _» 2FeS04 + (NH4)2S04 + S,
and then give ferrous sulphide (p. 376).
Ferric sulphate is formed by oxidation of ferrous sulphate, and is
obtained as a white mass by evaporation. It gives alums, such as
(NH4)2S04, Fe2(S04)3, 24H20, which are almost colorless when pure, but
usually have a pale reddish-violet tinge.
Pyrite. — The mineral pyrite FeS2 (Fools' gold) is the sulphide of
iron which is most stable in the air. It is found in nature in the form
of glittering, golden-yellow cubes, octahedrons, and pentagonal dodeca-
hedrons. It is not attacked by dilute acids, but concentrated hydro-
chloric acid slowly converts it into ferrous chloride and sulphur. It
is reduced by hydrogen to ferrous sulphide.
Cyanides. — When potassium cyanide is added to solutions of fer-
rous or ferric salts, yellowish precipitates are produced, but the simple
cyanides cannot be obtained in pure form. These precipitates interact
with excess of the cyanide giving soluble complex cyanides of the forms
4KCN,Fe(CN)a and 3KCN, Fe(CN)3 respectively. These are called
ferro- and ferricyanide of potassium, respectively.
Ferrocyanide of potassium, " yellow prussiate of potash," is made
by heating nitrogenous animal refuse, such as blood, with iron filings
IRON, COBALT, NICKEL 757
and potassium carbonate. The resulting mass contains potassium
cyanide and ferrous sulphide, and when it is treated with warm water
these interact and produce the ferrocyanide :
2KCN + FeS -» Fe(CN)2 + K2S,
4KCN + Fe(CN)2 -> K4.Fe(CN)6.
The salt is made also from the cyanogen contained in crude illumi-
nating-gas. It forms large, yellow, monosymmetric tables with three
molecules of water of hydration. The solution contains almost ex-
clusively the ions K* and Fe(CISr)6"", and gives none of the reactions
of the ferrous ion Fe". The corresponding acid H4Fe(CN)6 may be
obtained as white crystalline scales by addition of an acid and of
ether (in which the substance is less soluble than in water) to the salt.
The acid is a fairly active one but is unstable and decomposes in a
complex manner. Other ferrocyanides may be made by precipita-
tion. That of copper Cu2Fe(CISr)6 is brown, and ferric ferrocyanide
Fe4[Fe(CN)6]3 has a brilliant blue color ("Prussian blue"). The
ferrous compound Fe2Fe(CN)6, or perhaps K2FeFe(CN)6, is white but
quickly becomes blue by oxidation. The ferrocyanides are not poison-
ous.
Ferricyanide of potassium. This salt is easily made from the
ferrocyanide by oxidation :
2K4Fe(CN)6 + C12 -* 2KC1 + 2K8Fe(CN)6,
or 2Fe(CN)6"" + 01, -» 2Fe(CN)6'" + 201'.
It forms red monosymmetric prisms. The free acid H3Fe(CN)6 is un-
stable. Other salts may be prepared by precipitation. Ferrous ferri-
cyanide Fe3[Fe(CISr)6]2 is deep-blue in color (" TurnbulFs blue "). With
ferric salts only a brown solution is obtained.
Ferric thiocyanate Fe(C!N"S)3 is formed by interaction of soluble
thiocyanates with ferric salts (cf. p. 250). It is deep-red in color and
gives a blood-red solution in water. Since both the ions are colorless,
the solution must contain chiefly the molecular salt. Its formation
furnishes a very delicate test for traces of ferric salts.
Iron Carbonyls. — When carbon monoxide is led over finely
divided iron at 40-80°, or under eight atmospheres pressure at the
ordinary temperature, volatile compounds of the composition Fe(CO)4,
the tetracarbonyl, and Fe(CO)5, the pentacarbonyl, are formed. When
758 INORGANIC CHEMISTRY
the gaseous mixture is heated more strongly, the compounds decom-
pose again, and iron is deposited. Illuminating-gas burners frequently
receive a deposit of iron from this cause.
Ferrates. — A red solution of potassium ferrate, K2Fe04, is ob-
tained by passing chlorine through caustic potash in which ferric
hydroxide is suspended. The salt crystallizes in red, rhombic prisms,
isomorphous with the sulphate and chromate of potassium. It alters
quickly in solution, in consequence of hydrolysis and subsequent de-
composition of the ferric acid, depositing ferric hydroxide and giving
off oxygen. Barium, strontium, and calcium salts are formed as red
precipitates by double decomposition.
Analytical Reactions of Compounds of Iron. — There are 'two
ionic forms of iron, diferrion Fe", which is very pale-green, and tri-
ferrion Fe*%%, which is almost colorless. The yellow color of ferric
salts is due to hydrolysis. Ammonium sulphide forms, with both,
black ferrous sulphide which is soluble in dilute acids. The hydrox-
ides are white and brown respectively, and ferrous carbonate is white.
With ferric salts, soluble carbonates yield the hydroxide. With ferro-
cyanide of potassium, ferrous salts give* a white, and ferric salts a blue,
precipitate. With ferricyanide of potassium the former give a deep-
blue precipitate, and the latter a brown solution. Ferric thiocyanate
is deep-red. From ferric solutions barium carbonate throws down fer-
ric hydroxide. When sodium acetate is added in excess to a ferric
salt, a red, little ionized, but much hydrolyzed, acetate is formed.
When the solution is boiled the hydrolysis is increased, and an insol-
uble, basic ferric acetate is thrown down. With borax, iron com-
pounds give a bead which is green in the reducing flame, and colorless
or, with much iron, yellow or even brown when oxidized.
COBALT.
The Chemical Relations of the Element. — Cobalt forms
cobaltous and cobaltic oxides and hydroxides CoO and Co(OH)2, Co208
and Co(OH)s, respectively, which are all basic, the former more so
than the latter. The cobaltous salts are little hydrolyzed, but the
cobaltic salts are completely decomposed by water. The latter also
liberate readily one-third of the negative radical, after the manner of
manganic salts, becoming cobaltous. Complex cations and anions con-
taining cobalt are very numerous and very stable.
IRON, COBALT, NICKEL 759
Occurrence. — Cobalt is found along with nickel in smaltite CoAs2
and cobaltite CoAsS. The pure metal may be made by Goldschmidt's
process, or by reducing the oxalate, or an oxide, with hydrogen.
Physical and Chemical Properties, — The metal is silver- white,
with a faint suggestion of pink. It is less tough than iron, and has
no commercial applications. It displaces hydrogen slowly from dilute
acids, but interacts readily with nitric acid.
Cobaltous Compounds. — The chloride may be made by treating
the oxide with hydrochloric acid. It forms red prisms of CoCl2,
6H20, and when partially or completely dehydrated becomes deep-
blue. Writing made with a diluted solution upon paper is almost
invisible, but becomes blue when warmed and afterwards takes up
moisture from the air, and is once more invisible (" sympathetic
ink"). Most cobaltous compounds are red when hydrated or in solu-
tion (Co**) and blue when dehydrated. The blue color assumed by a
strong solution of cobaltous chloride, when it is warmed, or when
hydrochloric acid is added to it, is explained by some chemists as
being due to repression of the ionization of the salt, and by others
as being due to the formation of the complex anion of the salt
Co**.CoCl//. By addition of sodium hydroxide to a cobaltous salt, a
blue basic salt is precipitated. When the mixture is boiled, the red
hydroxide, Co(OH)2 is formed. This becomes brown through oxidation
by the air. It interacts with ammonium hydroxide, giving a soluble
ammonio-cobaltous hydroxide (<*/. p. 661), which is quickly oxidized
by the air to an ammonio-cobaltic compound (see below). It dissolves
also in salts of ammonium as magnesium hydroxide does (p. 644).
When dehydrated it leaves the black cobaltous oxide. Cobaltous sul-
phate, CoS04, 7H2O, isomorphous with magnesium sulphate, and the
nitrate, Co(N08)2, 6H20, are familiar salts. The black cobaltous sul-
phide, CoS, is precipitated by ammonium sulphide from solutions of
all salts, and even by hydrogen sulphide from the acetate, or a solution
containing much sodium acetate (cf. p. 650). Once it has been formed,
it does not interact even with dilute hydrochloric acid, having appar-
ently changed into a less active form. A sort of cobalt glass, made
by fusing sand, cobalt oxide, and potassium nitrate, forms, when
powdered, a blue pigment ("smalt") used in china-painting and by
artists.
760 INORGANIC CHEMISTRY
Cobaltic Compounds. — By addition of a hypochlorite to a solu-
tion of a cobaltous salt, cobaltic hydroxide Co(OH)3, a black powder,
is precipitated. Cautious ignition of the nitrate gives cobaltic oxide,
Co208. Stronger ignition gives the commercial oxide, which is a cobalto-
cobaltic oxide Co3O4. Cobaltic oxide dissolves in cold hydrochloric
acid, but the solution gives off chlorine when warmed. By placing
cobaltous sulphate round the anode of an electrolytic cell, crystals of
cobaltic sulphate, Co2(S04)s, have been made and cobaltic alums have
also been prepared.
Complex Compounds. — Potassium cyanide precipitates from
cobaltous salts a brownish-white cyanide which interacts with excess
of the reagent, giving a solution of potassium cobaltocyanide K4Co
(CN)6 (cf. p. 756). This compound is easily oxidized by chlorine, or
even when the solution is boiled in the air, and the colorless potassium
cobalticyanide is formed :
4K4Co(CN)6 + 2H20 + 02 -» 4K8.Co(CN)6 + 4KOH.
The solution gives none of the reactions of Co**", and with acids the
very stable cobalticyanic acid, H3Co(CN)3, is liberated.
When acetic acid and potassium nitrite are added to a cobaltous salt
the latter is oxidized by the nitrous acid (liberated by the acetic acid)
and a white complex salt K3.Co(N02)6, nR20 (= Co(N02)3,3KN02),
potassium cobaltinitrite, is thrown down.
Cobaltic salts give with ammonia complex compounds which are
many and various. The cations often contain negative groups, and
are such as Co(NH3y, Co(NH3)5Cl" and Co(NH3)5N02". Usually
the solutions give none of the reactions of cobaltic ions, and often fail
likewise to give those of the anion of the original salt.
NICKEL.
The Chemical Relations of the Element. — Nickel forms nick-
elous and nickelic oxides and hydroxides NiO and Ni(OH)2, Ni203, and
Ni(OJI)3, but only the former are basic. The nickelous salts resemble
the cobaltous and ferrous salts, but are not oxidizable into correspond-
ing nickelic compounds. Since there are no nickelic salts, there are
here no analogues of the cobalticyanides or the cobaltinitrites. The
complex nickelous salts, like the complex cobaltous salts, and unlike
the complex cobaltic salts, are unstable, and so give some of the
reactions of Ni".
IRON, COBALT, NICKEL 761
Occurrence. — Nickel occurs free in meteorites and in niccolite
NiAs and nickel glance NiAsS. It is now manufactured chiefly from
garnierite, a silicate of nickel and magnesium found in New Caledonia.
Properties. — The metal is white, with a faint tinge of yellow,
is very hard, and takes a high polish. It is used in making alloys,
such as German silver (copper, zinc, nickel, 2:1:1) and the "nickel"
used in coinage (copper, nickel, 3 : 1). Nickel plating on iron is
accomplished exactly like silver plating (p. 632). The bath contains an
ammoniacal solution of ammonium-nickel sulphate (NH4)2SO4, NiS04,
6H20, and a plate of nickel forms the anode.
The metal rusts very slowly in moist air. It displaces hydrogen
with difficulty from dilute acids but interacts with nitric acid.
Compounds of Nickel. — The chloride NiCLj', 6H20, is made by
treating one of the oxides with hydrochloric acid, and is green in
color (when anhydrous, brown). The sulphate, NiS04, 6H20, which
crystallizes in square prismatic forms at 30-40°, is the most familiar
salt. The heptahydrate NiS04, 7H20, obtained from cold solutions, is
isomorphous with magnesium sulphate. Nickelous hydroxide, Ni(OH)2,
is formed as an apple-green precipitate, and when heated leaves the
green nickelous oxide, NiO. It dissolves in ammonium hydroxide,
giving a complex nickel-ammonia cation. It is soluble also in salts of
ammonium (cf. p. 661). By cautious ignition of the nitrate, nickelic
oxide, Ni^, is formed as a black powder. The oxides and salts, when
heated strongly in oxygen give the oxide Ni304. The last two oxides
liberate chlorine when treated with hydrochloric acid, and give nickel-
ous chloride. Nickelic hydroxide, Ni(OH)s, is a black precipitate
formed when a hypochlorite is added to any salt of nickel. Nickelous
sulphide is thrown down by ammonium sulphide, and behaves like
cobaltous sulphide (p. 759). It forms a brown colloidal solution when
excess of the precipitant is used, and is then deposited very slowly.
With potassium cyanide and a salt of nickel the greenish nickelous
cyanide, Ni(CN)2, is first precipitated. This dissolves in excess of the
reagent, and a complex salt K2Ni(CN)4, H20( = 2KCN,Ni(CN)2) may
be obtained from the solution. This salt is of different composi-
tion from the corresponding compounds of cobalt and iron, and is less
stable. Thus, with bleaching powder, it gives Ni(OH)8 as a black
precipitate. When the solution is boiled in the air no oxidation to a
complex nickelicyanide occurs, and indeed no such salts is known.
762 INORGANIC CHEMISTRY
This fact enables the chemist to separate cobalt and nickel, for when
the mixed cyanides are boiled and then treated with bleaching powder,
the cobalticyanide is unaffected. With potassium nitrite and acetic
acid no insoluble compound corresponding to that given by cobalt salts
is formed by salts of nickel. The only known compound which could
be formed, 4KN02,Ni(N02)2, is soluble. This action also is used for
the purpose of separation.
When finely divided nickel, made by reducing the oxide or oxalate
with hydrogen at a moderate temperature, is exposed to a stream of
cold carbon monoxide, nickel carbonyl Ni(CO)4 is formed. This is a
vapor and is condensable to a colorless liquid (b.-p. 43° and m.-p. —25°).
The vapor is poisonous. When heated to 150-180° it is dissociated and
nickel is deposited. Cobalt forms no corresponding compound.
Analytical Reactions of Compounds of Cobalt and Nickel. —
The cobalt ion Co" is pink, and the nickelous ion Ni** green. The
reactions used in analysis have been described in the preceding para-
graphs. With borax, cobalt compounds give a blue bead, and nickel
compounds a bead which is brown in the oxidizing flame and cloudy,
from the presence of gray, metallic nickel, when reduced.
Exercises. — 1. What would be the interactions of calcium car-
bonate when fused with sand and with clay, respectively ?
2. Make equations representing (a) the oxidation of ferrous
chloride by air, (£) the hydrolysis of ferrous carbonate and the oxida-
tion of ferrous hydroxide, (c) the oxidation of ferrous sulphate with
excess of sulphuric acid by hypochlorous acid, (d) the formation of
ferrous and ferric tannates (p. 754), (e) the reduction of ferric chloride
by iron, by hydrogen sulphide, and by stannous chloride, respectively,
(/) the dry distillation of basic ferric sulphate, (<?) the formation of
ferric ferrocyanide and of ferrous ferri cyanide, (h) the hydrolysis and
decomposition of potassium ferrate.
3. Explain the solubility of cobaltous and nickelous hydroxide in
salts of ammonium.
4. Construct equations to show the formation (a.) of the insoluble
potassium cobaltinitrite (nitric oxide is given off), (b) of nickelic
hydroxide from nickelous chloride and sodium hypochlorite.
5. Tabulate in detail the chemical relations of the elements cobalt
and nickel, with especial reference to showing the resemblances and
differences.
CHAPTER XLV
THE PLATINUM METALS
THE remaining elements of Mendelejeff's eighth group divide them-
selves into two sets of three each. Just as iron, cobalt, and nickel
have similar atomic weights and much the same specific gravity,
(7.8-8.8), so ruthenium (Ru, at. wt. 101.7), rhodium (Rh, at. wt. 103),
and palladium (Pd, at. wt. 106.5) have specific gravities from 12.26 to
11.5. Similarly osmium (Os, at. wt. 191), iridium (Ir, at. wt. 193),
and platinum (Pt, at. wt. 194.8) form a triad with specific gravities
from 22.5 to 21.5. Chemically, ruthenium shows the closest resem-
blance to osmium, and both are allied to iron. Similarly, rhodium and
iridium, and palladium and platinum are natural pairs.
The six elements are found alloyed in nuggets and particles which
are separated from alluvial sand by washing. Platinum forms 60-84
per cent of the whole. The chief deposits are in the Ural Mountains,
smaller amounts being found in California, Australia, Borneo^ and
elsewhere. The components are separated by a complex series of
chemical operations.
Ruthenium and Osmium. — These metals are gray like iron,
while the other four are whiter and more like cobalt and nickel. They
also resemble iron in being the most infusible members of their
respective sets. Both melt considerably above 2000°. They likewise
resemble iron in uniting easily with free oxygen, while the other four
elements do not. Ruthenium gives Ru02 and even Ru04, although
the latter oxide is more easily obtained indirectly. Osmium gives
Os04, " osmic acid," a white crystalline body melting at 40° and boiling
at about 100°. The odor and irritating effects of the vapor recall
chlorine (Gk., 607*77, odor). The substance is not an acid, or even an acid
anhydride. The aqueous solution is used in histology, amd stains
tissues in consequence of its reduction by organic bodies to metallic
osmium. It is affected particularly by fat. Osmic acid also hardens
the material without distorting it. It will be observed that ruthenium
and osmium have a maximum valence of eight.
These elements resemble iron in giving ruthenates and osmates,
763
764 INORGANIC CHEMISTRY
like K2Ru04 and K20s04, but no corresponding oxide or acid. Potas-
sium ruthenate resembles potassium manganate and gives, when
diluted, an oxide of ruthenium and potassium perruthenate KKu04.
There are also salts corresponding to the oxides RugOj and Os208.
Rhodium and Iridlum. — These metals are not attacked by
aqua regia, while the other four are dissolved, more or less slowly.
They are harder than platinum, and indium is alloyed with this metal
for the purpose of increasing its resistance to the action of acids.
They resemble cobalt in having no acid-forming properties. Salts
corresponding to the oxides RlXjOg and Ir203 are formed, and iridium
gives compounds derived from Ir02 as well. The most familiar com-
pounds of iridium are the complex chlorides XglrClg (= SXClJrClj)
and XjIrClg ( = 2XCl,IrCl4). The solutions of the latter are red, and
the acid, chloroiridic acid H2IrClg, is often found in commercial chloro-
platinic acid H2PtCl6, and confers upon it a deeper color.
Palladium and Platinum. — Palladium is the only metal of
this family which is attacked by nitric acid. Palladium and platinum
form -ous and -ic compounds of the forms PdX2 and PdX4 respectively.
The oxides PdO and PtO and corresponding hydroxides are basic.
When quadrivalent, the metals appear chiefly in complex compounds,
like H^PtClfl, Hg-PdClg, in which the metal is in the anion. Platinum
gives also platinates derived from the oxide PtO2, and quadrivalent
platinum furnishes no well defined salts in which it constitutes by
itself the positive ion.
Palladium. — This metal, named from the planetoid Pallas, is
noted chiefly for its great tendency to absorb hydrogen. At 1000° it
takes up about 650 times its own volume. The amount absorbed
varies continuously with the concentration (pressure) of the hydrogen,
although not according to a uniform rule, and the case is therefore re-
garded as being one of solid solution. When a strip of palladium is
made the cathode of an electrolytic cell, over 900 volumes of hydrogen
may be occluded. This absorbed hydrogen, in consequence of the cat-
alytic influence of the metal, reacts more rapidly than does the gas, and
consequently a strip of hydrogenized palladium will quickly precipitate
copper and other metals less electro-positive than hydrogen and will
reduce ferric and other reducible salts :
CuS04 + H2 -» H2S04 + Cu, or Cu"+ H2 -> 2H' + Cu.
2FeCl3 + H2 -> 2FeCLj + 2HC1, or 2Fe"*+ H2-»2Fe" + 2H.
THE PLATINUM METALS 765
The palladious salts are such as PdCLj, PdS04, Pd(N08)2. Palladia
chloride, formed by treating the metal with aqua regia, is contained in
the solution in the form of chloropalladic acid H2PdClg (= 2HC1,
PdCl4) and gives difficultly soluble salts like K2PdCV When the
solution of the acid is boiled, however, chlorine is given off aud PdCl2
or H2PdCl4 remains in solution.
Platinum. — This metal (dim. of Sp. plata, silver) is grayish-white
in color, and is very ductile. At a red heat it can be welded. It does
not melt in the Bunsen flame but fuses easily in the oxy hydrogen jet.
On account of its very small chemical activity it is used in electrical
apparatus and for making wire, foil, and crucibles and other vessels
for use in laboratories. It unites, however, with carbon, phosphorus,
and silicon, becoming brittle, and forms fusible alloys with metals like
antimony and lead. Hence care has to be taken not to heat in vessels
made of it compounds from which these elements may be liberated. It
also interacts with fused alkalies, giving platinates, but the alkali car-
bonates may be melted in vessels of platinum with impunity. The
oxygen acids are without action upon it, but the free chlorine in aqua
reyia converts it into chloroplatinic acid K^PtC^.
The metal condenses oxygen upon its surface and it dissolves
hydrogen. The finely divided forms of the metal, such as platinum
sponge made by igniting ammonium chloroplatinate (NH4)2PtClg, and
platinum black made by adding zinc to chloroplatinic acid, show this
behavior very conspicuously. They cause instant explosion of a mix-
ture of oxygen and hydrogen, in consequenc of the heat developed by
the rapid union of that part of the gases which is condensed in the
metal. A heated spiral of fine platinum wire will continue to glow if
immersed in the mixture of alcohol vapor and oxygen formed by lead-
ing oxygen through liquid alcohol. The heat is developed by the inter-
action which takes place between the substances with great speed at the
surface of the platinum. Platinum sponge is the active constituent of
the contact-mass used in. making sulphur trioxide (p. 380).
Platinum is the only otherwise suitable substance which has the
same coefficient of expansion as glass, and it is consequently fused into
incandescent bulbs and furnishes the electrical connection with the
filament in the interior. Large amounts are also consumed in photo-
graphy. The price of the metal is subject to great variations, since a
rainy season in the Caucasus will render larger amounts accessible to
the miners, but, on the whole, the many applications which have been
found for it have tripled its price in the last twenty years.
766 INORGANIC CHEMISTRY
When special resistance to chemical or mechanical influences is
required, as in standard meters for international reference, or points
of fountain pens, the alloy with iridium is employed (cf. p. 240).
Compounds of Platinum. — Platinous chloride is made by pass-
ing chlorine over finely divided platinum at 240-250° or by heating
chloroplatinic acid to the same temperature. It is greenish and in-
soluble in water, but forms with hydrochloric acid the soluble chloro-
platinous acid H2PtCl4. The potassium salt of this acid K2PtCl4 is
used in making platinum prints (cf. p. 635). Bases precipitate black
platinous hydroxide Pt(OH)2 which interacts with acids but not with
bases. Gentle heating gives the oxide PtO and stronger heating the
metal. With potassium cyanide and barium cyanide soluble platino-
cyanides, K2Pt(CN)4, 3H2O and BaPt(CN)4, 4H20, are formed. These
substances, when solid, show strong fluorescence, converting X-rays as
well as ultra-violet rays into visible radiations. The barium salt is
used to coat screens on which the shadows cast by X-rays are received.
Platinic chloride PtCl4 may be made by treating the metal with
aqua regia and heating the chloroplatinic acid HjPtCL, so formed in
a stream of chlorine at .360°. When dissolved in water, however, it
gives a compound H2PtCl4O in which platinum is in the anion. Its
solution deposits red, non-deliquescent crystals of H2PtCl40, 4H20.
Chloroplatinic acid forms reddish-brown deliquescent crystals
HgPtClg, 6H20. With potassium and ammonium salts, it yields the
sparingly soluble, yellow chloroplatinates KgPtClg and (NH4)2PtCl6
(cf. p. 561), in solutions of which the platinum migrates towards the
anode and silver salts precipitate Ag2PtClg and not silver chloride.
Bases interact with chloroplatinic acid, giving a yellow or brown
precipitate of platinic hydroxide Pt(OH)4. This substance interacts
both with acids and with bases. In the former case the salts, which
probably are formed, have not been isolated. In the latter case plati-
nates, like Na2H10Pt3O12, H20, have been obtained. Both sets of plati-
num compounds interact with hydrogen sulphide, giving the sulphides,
PtS and PtS2 respectively. These are black powders which dissolve
in yellow ammonium sulphide solution, much as do the sulphides of
gold, arsenic, and other metals, giving ammonium sulphoplatinates.
There are numerous complex coinpounds with ammonia represent-
ing both the platinous and platinic series. They resemble the
ammonio-cobaltic compounds in their behavior.
INDEX
*#* Acids are all Hated under " acid " and salts under the positive radical.
ACCUMULATORS, 702
Acetylene, 496, 515
Acid, acetic, 498
ionization of, 578
antimonic, 716
arsenic, 711
arsenious, 711
boracic, 527
boric, 527
bromic, 277
carbolic, 441
carbonic, 481
chloric, 274
chlorosulphuric, 399
chlorous, 275
chromic, 728
disulphuric, 388
dithionic, 398
formic, 497
hydrazoic, 422, 452
hydriodic, 239, see Hydrogen iodide
hydrobromic, 233, see Hydrogen
bromide
hydrochloric, 93, see Hydrogen chlo-
ride
chemical properties, 185
of constant boiling-point, 182
hydrocyanic, 507
hydrofluoboric, 527
hydrofluoric, 242
hydrofluosilicic, 521
hypobromous, 276
hypochlorous, 176, 267
hyponitrous, 450
hypophosphorous, 469
hyposulphurous, 393
iodic, 277
metaphosphoric, 465, 468
metastannic, 696
muriatic, see Acid, hydrochloric
Acid, nitric, 93, 438
fuming, 440
oxidizing actions of, 446
preparation, 439
properties, 439
nitrosylsulphuric, 383
nitrous, 449
orthophosphoric, 465, 466, 732
oxalic, 499
palmitic, 505
pentathionic, 398
perchloric, 276
perchromic, 729
periodic, 278
permanganic, 743
persulphuric, 397
phosphorous, 469
picric, 441
pyrophosphoric, 467
tannic, 754
selenic, 402
silicic, 523
a-stannic, 696
sulphuric, 93, 382
chamber process, 383
constitution, 391
contact process, 380
fuming, 388
Nordhausen, 388
properties, 387, 389
uses, 393
sulphurous, 393
telluric, 403
tetraboric, 528
tetrathionic, 398
thiosulphuric, 396
trithionic, 398
Acidimetry, 351
Acids, 70, 92, 119, 281, 345
activity of, see Chemical activity
767
768
INDEX
Acids, complex, salts of, 360, 363, 536
constitution of, 470
dibasic, ionization of, 346, 374
ionic double decompositions of, 349
polythionic, 398
Activity, see Chemical activity
Adsorption, 476
Affinity, constant, 254
criticism of word, 29
inapt use of word, 110
see Chemical activity
Agar-agar, 313
Air, a mixture, 430
composition, 431
liquid, 434
solubility, 155
Albumins, 415
Alcohol, ethyl, 503
methyl, 503
Alizarine, 689
Alkalies, 348, 548
Allotropic modifications, 459
Alloys, 532
Alumina, 686
Aluminates, 685
Aluminium, 682
acetate, 688
carbide, 494
chloride, 684
hydroxide, 684
oxide, 686
sulphate, 686
sulphide, 687
Aluminothermy, 540
Alums, 547, 632, 687, 730, 741, 756
Amalgams, 532
Amethyst, 522
Ammonia, 417
Ammonia-soda process, 574
Ammonion, 568
Ammonium amalgam, 567
bromide, 564
carbamate, 566
carbonate, 565
chloride, 564
compounds, 420, 564
cyanate, 488, 566
hydroxide, 565
iodide, 564
nitrate, 451, 565
Ammonium sesquicarbonate, 566
sulphate, 566
sulphides, 566
sulphostannate, 697
thiocyanate, 566
Amorphous substances, 123, 139, 369,
475, 519
Analysis, application of ionic hypothe-
sis in, 343
of compounds of carbon, 623
recognition of cations ("metals")
in, 660
Analytical reactions, aluminium, 691
ammonium, 568
antimony, 720
arsenic, 720
barium, 612
bismuth, 720
cadmium, 652
calcium, 608, 612
chromium, 731
cobalt, 762
copper, 625
gold, 639
iron, 758
lead, 705
manganese, 745
magnesium, 645
mercury, 659
nickel, 762
potassium, 561
silver, 635
sodium, 577
strontium, 609, 612
tin, 697
zinc, 650
Anhydride, chromic, 728
hypochlorqus, 267
iodic, 278/279
nitric, 440
nitrous, 450
perchloric, 276
permanganic, 743
phosphoric, 464
sulphuric, 380
sulphurous, 378
Anhydrides, 71
mixed, 397, 463
various acids from one 278
Anions, 322
INDEX
769
Anode, 322
Antichlor, 397
Antimonites, 716
Antimony, 713
nitrate, 716
oxy chloride, 715
pentachloride, 715
pentoxide, 715
sulphate, 716
sulphides, 716
trichloride, 715
trioxide, 715
Antimonyl compounds, 716
Apatite, 455
Aqueous tension, 82, 88, 116, 135
Aqua regia, 448
Aragonite, 591
Argentic compounds, see Silver
Argol, 501
Argon, 428, 434
Arsenic, 707
pentoxide, 711
sulphides, 712
trichloride, 710
trioxide, 710
Arsine, 709
decomposition of, 252
Asbestos, 642
Asphalt, 492
Assaying, 639
Association, 242, 294
Atmosphere, 426
Atom, 196
definition of, 218
Atomic hypothesis, 217
Atomic volume, 407
Atomic weight, definition of, 201
Atomic weights, 50
determination of, 199
table of, 203, 204, inside rear board
Attributes of bodies, 35
Avogadro's hypothesis, 131, 190, 288
Azurite, 616
BACTERIA, nitrifying, 438
Balanced actions, 248
Barium, 609
carbonate, 610
chlorate, 611
chloride, 611
Barium, chloride, hydrated, 121
chromate, 728
dioxide, 611, see Peroxide
hydroxide, 612
nitrate, 612
oxide, 611
peroxide, 63, 303, 304, 611
dissociation of, 257
sulphate, 610
sulphide, 610
Baryta-water, 612
Bases, 71, 119, 281, 348
~ ionic double decompositions of, 350
solubilities of, 544
Basicity, 358
Battery-cells, 666, see Cells
Battery, theory of, 667
Bauxite, 685
Bead tests, 468, 528, 731, 745
Beer, 502
Benzene, 497
Benzine, 492
Beryllium, see Glucinum, 641
Bessemer process, 749
Bicarbonates, see Hydrogen carbonates
Bichromates, see Bichromates
Bismuth, 717
compounds, 718
Bisulphates, see Hydrogen sulphates
Bisulphites, see Hydrogen sulphites
Bleaching, 269, 394
powder, 266, 602
Blue-stone (blue vitriol), 120, 625
Body, meaning of term, 34
Boiling-point of mixed liquids, 182, 503
Boiling-points, elevation of, 162
of solutions, 162
Bone-ash, 456, 605
Bone black, 476
Borates, 528
Borax, 528
Boron, 526
carbide, 529
nitride, 529
trichloride, 527
trifluoride, 527
trioxide, 529
Boyle's law, 84, 288
Brandy, 502
Brass, 619
770
INDEX
Brimstone, 367
Erin's oxygen process, 63, 611
Bromine, 227
chemical properties, 230
physical properties, 229
preparation, 227
Bronze, 619
Brucite, 644
CADMIUM, 650
compounds, 651
Caesium, 563
Calamine, 646
Calcination, 456
Calcion, 608
Calcite, 591
Calcium, 589
carbide, 478, 602
carbonate, 591
chloride, 590
chromate, 727
fluoride, 591
hydride, 589
hydrogen sulphite, 395
hydroxide, 585, 596
nitrate, 602
oxalate, 597
oxide, 595
phosphates, 605
phosphide, 461
silicate, 606
sulphate, 603
sulphide, 604
Calculations, 58, 65, 91, 215
Calomel, 654
Calorie, 76
Cane-sugar, 500
Carbohydrates, 500
Carbon, 473
amorphous, 475
properties, 477
dioxide, 479
properties, 480
disulphide, 488
monoxide, 485
properties, 486
Carbonates, 482
Carbonates, acid, see Hydrogen car-
bonates
Carbonyl chloride, 487
Carborundum, 520
Carnallite, 551, 642, 643
Cassiterite, 693
Catalytic action, 75, 96, 111, 508
Cathode, 322
Cations, 322
Cause, in Science, 30
Caustic potash, 552
Cell, Bunsen, 673
Daniell, 672
dichromate, 673
dry, 673
gravity, 672
Leclanche, 673
storage, 702
Cells, battery, 666
concentration, 673
Cellulose, 500
Cement, 597
Cerium, 510, 682, 705
Chalk, 591
Chalybeate water, 483
Chance's process, 574
Characteristics of chemical phenom-
ena, 5, 16, 17, 26, 27, 41, 48
Charcoal, 476
properties, 476
Charles' law, 85, 87
Chemical activity, cause of, 28
measurement of, 28, 79, 111, 678,
679
of acids, 347, 356, 679
Chemical change, spontaneously pro-
gressing, 27
varieties of, 14, 99, 121, 176, 181, 187,
360 (ionic), 397, 520
Chemical relations of elements, 177,
226, 719
Chemistry, methods of work in, 36
Chinese white, 648
Cinnabar, 653
Chlorates, 272
Chlorides, electrolysis of, 169
interaction with acids, 179
solubility, 185
Chlorine, 168
bleaching action, 177
chemical properties, 174
chemical relations, 177
dioxide, 275
INDEX
771
Chloride, physical properties, 173
preparation, 168
test for, 175
uses, 177
Chloroform, 494
Choke-damp, 493
Chrome-alum, 730
Chromic acetate, 731
anhydride, 728
chloride, 729
oxide, 730
hydroxide, 729
sulphate, 730
Chromites, 729
Chromium, 722
trioxide, 728
Chromous compounds, 731
Chromyl chloride, 728
Clay, 525, 689
Coal, 477
Cobalt, 758
complex compounds, 760
zincate, 648
Cobaltic compounds, 760
Cobaltous compounds, 759
Coke, 475
Colloidal solution, 523
Colloids, 524
Columbium, 721
Combination, 14
Combining proportions, measurement
of, 43
Combining weights, 46
law of, 48
Combustion, 71
Common ion, effect of adding, 580
Complex cations, 534, 614, 622, 661, 662
cadmium, 652
.chromium, 729, 730
cobalt, 759, 760
copper, 620, 622, 623, 625
mercury, 659
nickel, 761
platinum, 766
silver, 629, 631
zinc, 648
Component, definition of, 34
Compound, 15
Compounds, binary, 103
molecular, 123, 443
Concentration, molar, 250
molecular, law of, 252
illustration of law of, 394
of gases, 80
Conditions, 35, 52
Conductivity, electrical, 325
Condy's disinfecting fluid, 743
Congo red, 355, 689
Conservation of mass, 17
Constant, depression, 291
ionization, 297
heat summation, law of, 78
Constituent, definition of, 34
Constitution of acids, 470
of substances, 104, 224, 279, 308, 391,
421, 441, 451, 466
Copper, 615
acetylene, 496, 516
properties, 618
Corpuscles, 222, 735
Couples, electrical, 96, 673
Corrosive sublimate, 655
Crayon, 591
Cream of tartar, 502
Critical phenomena, 133
Cryolite, 683
Crystal, forms, 137
structure, 140
Crystallization, water of, 123, 181, see
Hydrates
Cupric acetate, 624
bromide, 621
ionization of, 335
carbonate (basic), 623
chloride, 619
cyanide, 624
ferrocyanide, 286, 299, 757
hydroxide, 623
iodide, 621
nitrate, 623
oxide, 43, 622
sulphate, 624
hydrated, 120, 122
sulphide, 625
Cuprous chloride, 620
cyanide, 624
iodide, 621
oxide, 622
sulphide, 625
Cyanogen, 507
772
INDEX
DALTON'S law of partial pressures, 88
Deacon's process, 170
Decomposition, 15
potentials, 324, 675
Decrepitation, 556
Definite proportions, law of, 41
Deliquescence, 162
Densities of gases, 193
measurement of, 89, 90
Densities of solutions, 164
"Derived from," meaning of term, 278
Dextrose, 500
Dialysis, 523
Diamond, 474
Dichromates, chemical properties of,
725
Diffusion, in gases, 107
in liquids, 136
Dilution formula, Ostwald's, 298
Dimorphous substances, 369
Discharging potentials, 324, 675
Displacement, 95, 99
ionic, 342, 361
Dissociation, 121
hydrolytic, 181, see Hydrolysis
in solution, proofs of, 281, 289, 292,
293
of compounds (by heat), 208
pressure, 257
Distillation, 38
fractional, 399, 492
in vacuo, 276, 471
Dolomite, 642
Double decomposition, 15
Dulong and Petit's law, 211
Dyeing, 688
EARTHENWARE, 689
Efflorescence, 121, 123, 162
Electrolysis, 310, 675
of chlorides, 169
Electrolytes, 296, 310
Electromotive chemistry, 664
force, 665
series, 361, 670, 676
Electrons, 222, 735
Electroplating, 632
Element, definition of, 31, 32
not same as simple substance, 32,
57
Elements, base-forming, see Elements,
metallic
chemical relations of, 177, 226, 719
metallic, 119, 337, 404, 533
classification of, 537
Elements, molecular weights of, 206
negative, 322, 405
non-metallic, 119, 337, 405, 534
positive, 322, 404
quantities of terrestrial material, 33
Endothermal actions, 27
Energy, and chemical change, 19, 28
chemical, 25
conservation of, 23
definitions of, 23, 24
factors of, 677
free or available, 26
units employed in measuring, 25
Enzymes, 501
Epsom salts, 644
Equations, 54
making of, 55, 69, 208, 274, 447
molecular, 208
Equilibria, displacement of, 2f,7
Equilibrium, 117, 135
chemical, 246
ionic, 297
considered quantitatively, 578
Equivalents, 49, 102
Esters, 504
Ether, 147, 506
luminiferous, 24
Ethyl acetate, 504
sulphate, 504
Ethylene, 495
bromide, 496
Excess, meaning of term, 43
Exothermal actions, 27
Explanation, its nature, 10
Explosives, 452
FACTORS of energy, 77, 677
Fat, 505
Fehling's solution, 623
Fermentation, 501
Ferrates, 758
Ferric, compounds, 754, 755, 756
thiocyanate, 250, 757
Ferrous, compounds, 753, 754
sulphide, 12, 42, 371
INDEX
773
Fertilizers, 415, 467, 605
Filtrate, 11
Fire-damp, 493
Flame, 509
Bunsen, 511, 512
Flame, chemical changes in, 513
Flash-light powder, 643
Fluorine, preparation, 240
properties, 241
Fluorite, 591
Flux, 529
Formulae, 53
graphic, see Constitution of substances
making of, 55, see Equations
molecular, of compounds, 204
of simple substances, 205
structural, see Constitution of sub-
stances
Franklinite, 646, 756
Freezing mixtures, 164
Freezing-point, depression constant,
291, 294
Freezing-points, depression of, 163, 290
of solutions, 163, 290
Furnace, blast, 748
electric, 457, 479
reverberatory, 572
GALENA, 698, 705
Gallium, 681
Garnet, 525, 683
Gas, illuminating, composition of, 513
Gases, densities of, 193
drying of, 100
liquefaction of, 432
purification of, 100
solubility of, 153
Gasoline, 492
Gay-Lussac's law, see Charles' law
of combining volumes, 125
Germanium, 692
Glass, 606
Glauber's salt, 120, 158, 160, 576
Glucose, 500
Glycerine, 503
Gold, 635
compounds of, 638
Gram-molecular volume, 199
Grape-sugar, 500
Graphite, 475
Green fire, 611
Greenockite, 650
Gun-cotton, 441, 452, 504
Gunpowder, 557
Gypsum, 603
HALOGEN family, 226, 243, 279
Heat, of fusion, 115
of solution, 164
of vaporization, 118
Heavy-spar, 610
Helium, 436, 563, 736
Homologous series, 491
Hydrargyllite, 685
Hydrates (hydroxides), 120
(substances containing water of
crystallization), 120
of salts, composition of, 545
vapor tension of, 121
Hydrazine, 422
Hydrion, properties of, 346
Hydrocarbons, 490
properties of, 493
Hydrogele, 523
Hydrogen, 92
chemical properties, 108
physical properties, 106
preparation, 93, 95, 97, 99, 362
bromide, preparation, 230
properties, 232
chloride, chemical properties, 183
composition, 183
physical properties, 182
preparation, 178, 585
fluoride, physical properties, 241
preparation, 241
iodide, dissociation of, 254
interaction with sulphuric acid, 237
preparation, 237
properties, 239
pentasulphide, 376
peroxide, 126, 303
selenide, 401
sulphate, properties, 387
sulphide, 371
chemical properties, 373, 374
telluride, 403
Hydrolysis, definition of, 181
of halogen compounds, 534
of salts, 344
774
INDEX
Hydrosole, 523
Hydroxidion, properties of, 349
Hydroxylamine, 423
Hypobromites, 276
Hypochlorites, 265
Hypotheses, formulative, 141
Hypotheses, stochastic, 142
Hypothesis, 128
atomic, 217
Avogadro's, 131, 190
corpuscular, 222
ionic, 296, 317
kinetic-molecular, 128
molecular applied to solutions, 150
ICE, 115
Iceland spar, 591
Identification, means of, 35, 39
Illuminating-gas, composition of, 513
Indicators, 353, 355
Indigo, 688
Indium, 681
Ink, 754
Insoluble compounds, theory of preci-
pitation of, 581, 585, 601, 644
theory of solution of, 585, 598, 601,
621, 644, 648
Internal rearrangement, 15, 488
Inversion, 500
Iodine, 233
Iodine, chlorides, 244
pentoxide, 278, 279
preparation, 234
properties, 235
uses, 236
lodimetry, 236
lodoform, 494
Ion, common, effect of adding, 580
definition of, 321
product, 583
Ionic concentration, 585
hypothesis, 276, 317
objections to, 319
lonization, degrees of, 329
lonogens, 296
Indium, 764
Iron, 746
alum, 756
carbonyls, 757
cast, 748, 751
Iron, chemical properties, 752
cyanides, 756
wrought, 749
see Ferrous and Ferric
Isomers, 262, 488, 696
Isomorphism, 545
KAINITE, 360, 559
Kalion, 561
Kaolin, 525, 689
Kerosene, 492
Kieserite, 645
Krypton, 437
LAKES, 689
Lanthanum, 681, 682
Lampblack, 476
Law, definition of, 7
Dulong and Petit's, 211
Faraday's, 315
Gay-Lussac's, 125
Henry's, 154
Le Chatelier's, 260
misconceptions in regard to the na-
ture of, 7
of combining weights, 48
of mobile equilibrium, 260
of partition, 155
periodic, 410
van 't Hoff's, 260
Laws, two kinds of, 191
Lead, 698
acetate, 704
carbonate, 704
chlorides, 700
chromate, 727
hydroxide, 701
iodide, 700
nitrate, 703
oxides, 700, 701, 702
sugar of, 704
sulphate, 704
sulphide, 705
Le Blanc soda process, 572
Light, see Photochemistry
Lignin, 395
Lime, chloride of, see Bleaching powder
milk of, 596
superphosphate of, 467, 605
Limestone, 591, 592, 595
INDEX
775
Liquefaction of gases, 432
Litharge, 700
Lithium, 577
Litmus, 70, 355
Lodestone, 756
Lunar caustic, 631
MAGNALIUM, 683
Magnesia alba, 644
Magnesium, 642
ammonium arsenate, 645
carbonate, 644
chloride, 643
hydroxide, 644
oxide, 643
phosphates, 645
sulphate, 644
sulphide, 645
Malachite, 616
Manganates, 737, 741
Manganese, 737
dioxide, 63, 171, 398
oxides of, 738
Manganic compounds, 737, 740
Manganites, 737, 741
Manganous compounds, 737, 739, 740
sulphate, 739
Marsh gas, 493, 709, 714
Mass action, 250
Massicot, 700
Matches, 460
Matrix, 367
Matter, 23
Mercur-ammonium compounds, 659
Mercuric, cyanide, 658
fulminate, 658
oxide, 12, 61
thiocyanate, 658
Mercury, 652
chlorides, 654
iodides, 656
nitrates, 657
oxides, 656
sulphides, 657
Metallic elements, 119, 337, 404, 534
Metals, electrical conductivity of, 532
extraction from ores, 540
hydroxides of, 541
occurrence of, 539
oxides of, 541
Metals, physical properties of, 530
see Metallic elements
"Metals," recognition of, in analysis
660
Metastable condition, 159
Methane, 493
Methyl, formate, 504
orange, 355
Mica, 525
Microcosmic salt, 467, 567
Migration, ionic, 312
speed of, 314
Minium, 701
Mobile equilibrium, law of, 260
Mohr's salt, 754
Molar, volume, 199
weight, 197
Mole, 197
Molecular, compounds, 443
magnitudes, 140
weight, chemical unit of, 196
weights, 191
in solution, 289, 292, 305
of elements, 206
rule for measuring, 197, 199, 210
Molecule, the chemical, 198, 296
Molybdenum, compounds of, 732
Mordants, 689
Mortar, 597
Multiple proportions, law of, 41
NAPHTHA, 492
Nascent state, 272, 423, 446
Natrion, 577
Neon, 437
Neodymium, 682
Nernst lamp, 706
Nessler's reagent, 659
Neutralization, 266, 351
theory of, 353
thermochemistry of, 357
volume change in, 358
Nickel, 700
carbonyl, 762
compounds, 761, 762
Niobium, 721
Nitrates, 442
Nitric oxide, 442
Nitrites, 449
Nitrogen, 62, 415
776
INDEX
Nitrogen, chloride, 424
iodide, 425, 452
oxides, 438
oxygen acids of, 438
tetroxide, 444
Nitroglycerine, 441, 452, 504
Nitrosyl chloride, 448
Nitrous oxide, 450
Nomenclature, acids and salts, 263
ionic hypothesis, 296, 321
Non-ionic actions, 364
Non-metallic elements, 119, 405, 534
OCTANTS, law of, 406
Oleum, 388
Orpiment, 712
Osmium, 763
Oxidation, 70, 72, 110, 454
Oxides, 70
Oxygen, 61
atomic weight why sixteen, 196
chemical properties, 67
physical properties, 66
preparation, 63
Ozone, 300
Ozocerite, 492
PALLADIUM, 764
Paper, manufacture of, 395, 501, 686
Paris green, 624, 712
Parke's process, 627
Pattinson's process, 627
Pearl ash, 558
Perfect gas, 85
Periodates, 278
Periodic system, 407
applications of, 412
Peroxides, 308
Petroleum, 491
ether, 492
Pewter, 694
Phase rule, 592
Phases, 156
Phenol, 441
Phenolphthalei'n, 355
Phlogiston, 9
Phosgene, 487
Phosphine, 461, 469
Phosphonium compounds, 462
Phosphorus, 455
Phosphorus, acids of, 465
halides, 462
oxides, 464
pentachloride, 462
dissociation of, 255
pentoxide, 464
preparation, 455
properties, 457
red, 457
sulphides, 471
trichloride, 462
trioxide, 464
uses, 460
yellow, 457
Photochemistry, 458, 484, 633
of plants, 483
Photography, 633
Physics needed in the study of chem-
istry, 39, 65
Pink-salt, 695
Plaster of paris, 603
Platinum, 765
compounds of, 766
Plumbates, 702
Plumbic, see Lead.
Polarization, 324, 675
Polymerization, 242
Polysulphides, 376
Porcelain, 690
Potassamide, 420
Potassium, 549
aluminium sulphate, 687
arsenite, 712
bromate, 556
bromide, 552
carbonate, 557
chlorate, 64, 275, 555, 583
chloride, 550
chromate, 723
cyanate, 507, 559
cyanide, 558
dichromate, 725
ferrate, 758
ferricyanide, 757
ferrocyanide, 756
fluorides, 552
hydride, 550
hydrogen sulphate, 560
sulphide, 560
tart rate, 561
INDEX
777
Potassium, hydroxide, 552, 558
iodate, 556
iodide, 551
metantimoniate, 716
nitrate, 556
oxides, 555
percarbonate, 558
perchlorate, 556
permanganate, 742
plumbate, 702
pyrosulphate, 56Q
sulphate, 559
sulphides, 560
thiocyanate, 508, 559
zincate, 648
Potential, chemical, 678
differences (anions), 676
(cations), 670
discharging, 675
Precipitate, 13
Precipitation, in ionic actions, 339,
350
rule for, 587
rule of (Berthollet's), 259
Pressure, dissociation, 257
osmotic, 151, 283, 299
partial, 88
solution, 152
vapor, 115, 135
Principles, summary of, 188, 262
Properties, specific, 35
Proustite, 626, 713
Prussian blue, 757
Pumice-stone, 525
Pure, chemically, definition of, 34
Putrefaction, 501
Pyrargyrite, 626, 717
Pyrite, 42, 756
Pyrosulphates, 388
QUARTATION, 532
Quartz, 522
Quicklime, 592, 595
Quicksilver, 653
RADICALS, 93
organic, 494, 499
valence of, 104
Radio-activity, 733
Radium, 733
Realgar, 712
Red, fire, 658
heat, tempferature corresponding to,
73 ]
lead, 701
Reduction, 72; 110
Reversible act?ons, 176, 179, 237, 246
Rhodium, 764
Rhombohedron, 14
Roasting, 378
Rock crystal, 522
Rouge, 755
Rubidium, 563
Ruthenium, 763
SALAMMONIAC, 564
Saltpeter, 438, 569
Salts, 264, 281, 337
acid, 358
basic, 359
composition of hydrates of, 545
definition of term, 365
double, 360
general methods of making, 542
ionic double decompositions of, 337
mixed, 359
nomenclature of, 263
of complex acids, 360, 363
of hydrogen, 345
of hydroxyl, 348
precipitation of, 339, 585
solubilities of, 157, 544
Samarium, 681
Sand, 522
Saponification, 505
Scandium, 681
Scheele's green, 624, 712
Schlippe's salt, 717
Sciences, classification of, 3
Scientific method, 4, 6, 10, 16, 17, 23,
28-30, 128, 141, 191
Selenium, 401
Semi-permeable partitions, 284, 299
Shoenite, 559
Silicates, 524
Silicon, 518
carbide, 520
dioxide, 522
hydride, 519
tetrachloride, 520
778
INDEX
Silicon, tetrafluoride, 521
Silver, 626
arsenate, 632
carbonate, 631
chloride, 14
chromate, 632, 728
complex compounds of, 629
cyanide, 630, 632
halides, 629, 633
nitrate, 13, 631
orthophosphate, 632
oxides of, 630
-plating, 632
properties, 628
sulphate, 632
sulphide, 632
thiosulphate, 630
Simultaneous actions, 231, 272, 275
Slag, 540
Soap, 505
Soda, crystals, 573
washing, 573
Sodium, 569
aluminate, 99, 685
amalgam, 569
arsenite, 712
carbonate, manufacture, 571
properties, 575
chloride, 13, 570
chromate, 724
dichromate, 725
hydride, 569
hydrogen carbonate, 575
periodate, 279
hydroxide, 570
hyposulphite, 393
"hyposulphite," see Thiosulphate
nitrate, 14, 571
nitrite, 571
oxide, 571
pentasulphide, 376
peroxide, 303, 308, 571
phosphates, 577
plumbite, 701
pyroantimoniate, 577, 716
silicate, 577
stannate, 696
sulphate, 576
hydrated, 120, 158, 160
tetraborate, 528, 577
Sodium, thiosulphate, 576
zincate, 99
Soldering, 421
Solubility, curves of, 157, 158
data, 147, 154, 157, 531
independent, 153
influence of temperature on, 156
limits of, 146
measurement of, 147
of gases, 153
product, 583
separation of substances by, 273
Solution, 145
definition of, 165
heat of, 164
in two solvents, 155
of insoluble substances, 585, 598,
601, 621, 644, 648
scope of word, 146
Solutions, boiling-point of, 162
densities of, 164
freezing-points of, 163
general properties of, 145
is-osmotic, 285
molar, 149
normal, 148
saturated, 148, 153, 161, 581
solid, 146
standard, 236
supersaturated, 159
vapor tension of, 161
Solution tension, hypothesis of, 670
Solvay soda process, 574
Specific heats of elements, 211
Spectroscope, 561
Spectrum, 562
Speed of reaction, 111, 251, 394
affected by, catalysis, 75
concentration, 74, 250, 394
solution, 76, 283
temperature, 73
Spinelles, 686
Spirit of hartshorn, 418
Stable, meaning of term, 119
Stannous, see Tin.
Starch, 500
Steam, 115
Stearin, 506
Steel, properties of, 750
theory of tempering, 750
INDEX
779
Stibine, 714
Stibnite, 714
Storage battery, 702
Strontium, chloride, 608
hydrated, 121
chromate, 728
compounds, 608
peroxide, 305
Structure, molecular, 213, 224, see
Constitution.
Sublimation, 235, 463
Substance, meaning of term, 30, 32, 35
Substitution, 176
Suint, 557
Sulphantimoniates, 717
Sulphantimonites, 717
Sulpharsenates, 712
Sulpharsenites, 712
Sulphates, 390
acid, 390, see Hydrogen sulphates
Sulphides, 375
relative solubilities of, 651
Sulphites, 395
Sulphocyanates, see Thiocyanates
Sulphostannate, 697
Sulphur, 11, 367
chemical properties, 370
chemical relations, 377
crystalline form, 11
dioxide, 378
family, chemical relations of, 403
monochloride, 398
oxides of, 378
oxygen acids of, 381
physical properties, 368
tetrachloride, 399
trioxide, 380
Sulphuretted hydrogen, see Hydrogen
sulphide
Sulphuryl chloride, 399
Summary of principles, 188, 262
Superphosphate of lime, 467, 605
Sylvite, 551
Symbols, 53
Synthesis, 473
System, meaning of term, 158
TANTALUM, 721
Tartar-emetic, 716
Tellurium, 402
Temperature, critical, 1 34
Tests, 99
Thallium, 681
Theory, see Hypothesis
Thermochemistry, 76, 233, 271, 307, 357
Thionyl chloride, 399
Thio-, see Sulpho-
Thorium, 510, 705
Tin, 693
bromide, 695
chloride, 694, 695
hydroxide, 696
oxides, 696
-stone, 693
sulphides, 697
Titanium, 705
Titration, 352
Touch-paper, 557
Transformation by steps, 453
Transition points, 115, 118, 369, 594
Triethylamine, 156
Tungsten, 732
Turnbull's blue
Turquoise, 683
Type-metal, 699
ULTRAMARINE, 690
Units, 25, 58, 533
electrical, 665
Uranium, compounds, 733
Urea, 487
VALENCE, 101
definitions of, 102, 103, 106
exceptional, 106
multiple, 105
of radicals, 104
Vanadium, 721
Vapor, densities, 90
pressure, 115, 135
tension of solutions, 161
Velocity, see Speed of reaction
Venetian red, 755
Verdigris, 624
Vermilion, 658
Vinegar, 498
Vitriol, blue, 120, 625
green, 754
oil of, see Acid, sulphuric
white, 649
780
INDEX
Vitriols, 546, 649
Volume, atomic, 407
molar, 199
Volumetric analysis, 353
WATER, 113
chemical properties, 118
composition, 124
electrolysis, 676
gas, 485
hard, 113, 594
ionization of, 331
mineral, 113
of crystallization, see Hydrates
physical properties, 114
purification, 113
Weight, molar, 197
Weights, atomic, see under Atomic
weights
molecular, see under Molecular
weights
Weldon process, 741
Welsbach mantles, 705
Whiskey, 477, 502
White lead, 704
Wine, 501
XENON, 437
YTTERBIUM, 681
Yttrium, 681
ZINC, 646
acetate, 649
-blende, 646, 649
carbonate, 649
chloride, 647
hydroxide, 648
oxide, 646, 648
sulphate, 649
sulphide, 649
-white, 648
Zincates, 648
Zirconium, 705
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INTERNATIONAL ATOMIC WEIGHTS (1906).
ELEMENTS
SYMBOLS
AT. WT.
O = 16
ELEMENTS
SYMBOLS
AT. WT.
O=16
Aluminium . . .
Antimony ....
Al
Sb
27.1
120 2
Neodyinium . . .
Neon
Nd
Ne
143.6
20
Argon
A
39 9
Nickel
Ni
58 7
Arsenic
As
75.0
Nitrogen ....
N
14 04
Barium
Ba
137.4
Osmium ....
Os
191
Bismuth ....
Boron
Bi
B
208.5
11.0
Oxygen ....
Palladium
O
Pd
16.00
106 5
Bromine ....
Cadmium ....
Caesium ....
Calcium ....
Carbon
Br
Cd
Cs
Ca
c
79.96
112.4
132.9
40.1
12.00
Phosphorus .
Plantinum . .
Potassium
Praseodymium . .
Radium ....
P
Pt
K
Pr
Ra
31.0
194.8
39.15
140.5
225
Cerium
Ce
140.25
Rhodium ....
Rh
103 0
Chlorine ....
Chromium ....
Cobalt
Cl
Cr
' Co
35.45
52.1
59 0
Rubidium . . .
Ruthenium . . .
Samarium
Rb
Ru
Sa
-85.5
101.7
150 3
Columbium . . .
Copper .
Cb
Cu
94.
63 6
Scandium . . .
Selenium
So
Se
44.1
79 2
Erbium ....
Fluorine ....
Er
F
166.
19,
Silicon
.Silver
Si
Ag
28.4
107 93
Gadolinium . . .
Gallium ....
Germanium
Glucinum ....
Gold
Gd
Ga
Ge
Gl
Au
156.
70.
72.5
9.1
197 2
Sodium ....
Strontium . . .
Sulphur ....
Tantalum ....
Tellurium
Na
Sr
S
Ta
Te
23.05
87.6
32.06
183.
127 6
Helium
He
4
Terbium ....
Tb
160
Hydrogen ....
Indium . . . .
Iodine
H
In
I
1.008
115.
126.97
Thallium ....
Thorium ....
Thulium ....
Tl
Th
Tm
204.1
L'32.5
171
Iridium
Ir
193 0
Tin
Sn
119 0
Iron
Fe
55.9
Titanium
Ti
48 1
Krypton ....
Lanthanum . .
Lead
Kr
La
Pb
81.8
138.9
206 9
Tungsten ....
Uranium ....
Vanadium
W
U
V
184.
238 5
51 2
Lithium ....
Li
7.03
Xenon
Xe
128
Magnesium
Manganese. . .
Mercury ....
Mg
Mn
Hg
24.36
55.0
200.0
Ytterbium . . .
Yttrium ....
Zinc
Yb
Yt
Zn
173 0 .
89.0
65 4
Molybdenum . . .
Mo
96.0
Zirconium
Zr
'JO. 6
Only significant figures are given in this table. Where no figure follows the
decimal point, the value of the first decimal is uncertain.
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