Principles of inorganic chemistry

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PRINCIPLES OF
INORGANIC CHEMISTRY

2^^^

■^"■^ o

PRINCIPLES

OF

INORGANIC CHEMISTRY

BY

HARRY C. ^ONES

PROFESSOB OF PHYSICAL CHBMI8TBT IH THS
JOHNS HOPKINS UNIVKBSITT

THIBD EDITION

THE MACMILLAN COMPANY

LONDON: MACimXAN & CO., IffD.

1906

-^

COPTBIOHT, 1906,

By the MACMILLAN COMPANY.

Set up, electrotyped, and published January, 1903. Rq>rinted
May, 1904 ; October, 1906.

N'ortaioati l^rrff

I n I iiahiiiu ^ I II itiTviifk A Hmlth Co.

Ni.iHiMi.l, Mum., C.S.A.

•T17

PREFACE

Inoroanig Chemistry within the last few years has undergone
remarkable developments. This is due chiefly to generalizations
which have been reached through physical chemistry. We can see
most clearly what these developments are by comparing the inor-
ganic chemistry of twenty years ago with that of to-day. Until
recently the more important generalizations upon which the science
of inorganic chemistry rested were : The conservation of mass and
energy ; the laws of definite and multiple proportion^ and combining
weights ; the law of Avogadro, and the periodic system. Inorganic
chemistry was built upon these generalizations, and consisted largely
in a description of the compounds formed as the result of the inter-
action of matter in terms of these laws. Relations between the
composition and properties of compounds of different elements were
pointed out, which were more or less deep-seated and far-reaching.

Within the last fifteen years several newly discovered generali-
zations have been added to those longer known, and some of these
have been shown to be fundamental to the whole science of chem-
istry. The more important of these generalizations are : The theory
of electrolytic dissociation ; the law of mass action ; the phase rule,
and Faraday's law as the basis of chemical valence.

That these generalizations are of the very greatest importance for
inorganic chemistry is obvious to any one who is familiar with the
facts of physical chemistry arid of inorganic chemistry. Take the
theory of electrolytic dissociation, put forward by Van't Hoff and
Arrhenius. We know to-day that nearly all inorganic reactions are
reactions between ions; molecules and atoms as such having noth-
ing to do with the reactions. They simply serve to furnish the
ions, which are chemically the active agents. This obviously neces-
sitates a fundamental change in our conceptions of chemical phe-
nomena. It is not the uncharged atoms which react chemically,
but these become chemically active only when they carry an elec-
trical charge. The chemistry of atoms and molecules is thus
largely replaced by the chemistry of ions.

Similarly, the law of mass action of Guldberg and Waage has
produced a fundamental change in our method of regarding chemi-

v

66960

VI PREFACE

cal reactions. It has not only shown that mass is an important
factor in determining the magnitude of any given reaction, but in
many cases can actually determine the direction of the reaction. In
this law we have not simply a qualitative statement of the effect of
mass, but a quantitative relation mathematically formulated.

The phase rule of Willard Gibbs has also played its part in the
recent developments in inorganic chemistry. It has laid special
stress upon the conditions of equilibrium between the different
phases of the same and different substances, and has predicted the
existence of unknown substances, many of which have recently
been found. The phase rule is a beautiful, short-hand expression
of great masses of facts, and it gives us a comprehensive grasp of
these facts which without it would be impossible.

The application of Faraday's law as the basis of chemical valence
is not a new conception, but the importance of this application has
only recently become apparent. The importance of ions, which are
charged atoms or groups of atoms, and the study of electrochemical
phenomena in general have made prominent the fact that the law of
Faraday is a fundamental law of chemistry as well as of physical
chemistry. If we do not recognize this relation, the term " chemical
valence" is without exact significance and meaning; when based
upon the law of Faraday — a law to which thus far no exception is
known — valence has an exact physical basis, and this is an impor-
tant step for the development of chemistry.

The object of the present work is not to abandon the older gen-
eralizations, nor even the older methods of treating chemical phe-
nomena in so far as they cannot be improved. On the contrary, the
attempt has been made to retain those genei-alizations at their full
value, and special stress is laid upon the periodic system, which has
apparently fallen in certain directions somewhat into disrepute.
This generalization reached by Lothar Meyer and Mendel^eff, which
has been the philosophy of inorganic chemistry for so many years,
and in terms of which so much has been discovered, is still of very
great value. That it has serious defects no one can doubt; that
these are far exceeded by its merits is strongly impressed upon the
writer.

The aim of this book is to add to the older generalizations those
recently discovered, and to apply them to the phenomena of inor-
ganic chemistry in such a way that they may form an integral part
of the subject, and, at the same time, be intelligible to the student.
Why should we continue to teach the chemistry of atoms to students
on the ground of its being a little simpler, perhaps, than the chem-

PREFACE VU

istry of ions, or on any other ground, if we know that it is not in
accordance with the recently discovered facts ? Or why should we
continue to teach purely descriptive chemistry when the science of
chemistry has outgrown this stage, and many of the most important
relations have been accurately formulated in terms of the simpler
mathematics ?

These are questions which need only to be asked in order that
their answer may be made apparent. If a student can grasp the
conception of an atom and cannot add to this the idea of the atom
carrying an electrical charge, his hope of ever learning anything of
chemical phenomena in general is not bright.

The second point, the introduction of elementary mathematics
into chemistry, may seem to be more serious. The earlier text-
books on inorganic chemistry have been characterized by the ab-
sence of anything resembling a mathematical symbol, and chemistry
has come to be known as a non-mathematical science. It is, how-
ever, obvious to any one who has traced the development of science
that this condition of things cannot last. Physics has passed
through the stages through which chemistry is now passing. Fara-
day, the leading physicist of his day, was not a mathematician, but
how different at present, when to be a physicist one must first be a
fairly good mathematician. Indeed, it has been well said, that the
state of development of any branch of natural science can be meas-
ured by the extent to which mathematical methods have been
applied to it. That chemistry will become more and more mathe-
matical is just as certain as that it will develop.

There seems to be no good reason why we should refrain for a
moment from introducing simple, algebraic symbols into a compara-
tively elementary text-book in inorganic chemistry, where these
best serve to bring out the relations between phenomena. They
are introduced without question on almost every page of elementary
works on physics, and have come to be taken as a matter of course.
The same class of students, and frequently the same students, use
these works and corresponding texts on chemistry. Why should
chemists be hampered by being comx)elled to describe phenomena
at length when these could be formulated in a single line ? The
time has come when they need not be, and the earlier elementary
mathematics is introduced into text-books on chemistry, the better
for chemistry and for the chemist.

The writer has refrained from introducing unproved theories and
disputed questions as far as possible, since the student for whom
this work is meant is scarcely at a stage to properly appreciate and

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PREFACE TO THE SECOND EDITION

The short time that has elapsed since the appearance of the first
edition of this work but serves to show how great a demand exists
on the part of teachers and students of chemistry for the introduc-
tion of the recently discovered generalizations of physical chemistry
into their science. We can now add to the description of reactions
and substances formed as the final results of such reactions, certain
of the laws which condition them and to which they in general con-
form. The final result will be a science of chemistry, not as exact,
perhaps, as physics, but sufficiently exact to enable us to deal with
many phenomena in terms of the simpler mathematics. This is the
direction in which all branches of natural science are tending, and
by which their state of development may be fairly gauged.

A large number of minor corrections and changes have been
introduced into the second edition of this work. To most of these
my attention has been called by a large number of friends, to whom
I wish to take this opportunity to extend my hearty thanks.

It is to be hoped that any one who may discover an error in this
edition will kindly communicate the fact to the author.

HAKRY C. JONES.

PREFACE TO THE THIRD EDITION

The importance of physical-chemical generalizations for the
teaching of general chemistry is now generally recognized. This is
shown by the large demand for text-books of which these generaliza-
tions are made the basis, and by the increasing number of such
works that are being produced.

The teaching of even elementary chemistry has been fundament-
ally modified in the light of these recently discovered relations, and
teachers who desire to keep abreast with progress and to give their
students the results of the latest and most fruitful developments, are
becoming familiar with these relations and using them in the class-
room. This applies to the colleges fully as much as to the
universities.

The progress in this direction that has been made in the last ten
years will probably be greatly outstripped during the next decade.

In this edition the atomic weights recommended by the Interna-
tional Committee for 1906 are used.

Ht C* Jt

CONTENTS

CHAPTER I

PAOl

Intboduction 1

The Study of Nature. Relations between Chemistry and Physics.
Elements and Compounds. The Number of Elements'and Compoands.
The Chemical Elements. Chemical Combination.

CHAPTER n

Generalizations 7

The Science of Chemistry. Generalization. The Law of the Con-
servation of Mass. The Law of Constant Proportion. The Law of
Multiple Proportions. The Law of Combining Weights. The Atomic
Theory. The Correlation and Conservation of Energy. Importance
of the Conservation of Energy for the Science of Chemistry.

CHAPTER m

OXTGEN 16

Occurrence in Nature. Preparation of Oxygen. Hydrogen Dioxide.
Substances bum readily in Oxygen. Explanation of the Above Re-
sults. Combustion. The Phlogiston Theory of Combustion. The
Kdle of Oxygen in Combustion. Increase in Weight in Combustion.
Oxygen used up in Combustion. Rapid and Slow Oxidation. Meas-
urement of the Heat of Combustion. Heat of Formation and of
Decomposition. Names of the Compounds formed with Oxygen.
Certain Physical Properties of the Element Oxygen. The Pressure of
Oxygen varies with the Conditions. The Iaw of Boyle for Gases.
The Law of Gay-Lussac for Gases. The Determination of the Abso-
lute Zero of Temperature. The Combined Expression of the Laws of
Boyle and Gay-Lussac. The Liquefaction of Oxygen. Properties of
Liquid Oxygen. Power of Oxygen to enter into Chemical Combination.

Ozone 29

Allotropic Modification of Oxygen. Preparation of Ozone. Prop-
erties of Ozone. Transformation of Ozone into Oxygen. The Differ-
ence between Oxygen and Ozone.

ix

CONTENTS

CHAPTER IV

PAOB

Hydrogen 83

Occurrence. Preparation of the Element Hydrogen. Combination
of Hydrogen with Oxygen. Mixture of Hydrogen and Oxygen affected
by the Presence of Certain Substances. Catalytic Reactions and
Catalyzers. Relations by Volume in which Hydrogen and Oxygen
combine. Heat Energy produced when Oxygen and Hydrogen com-
bine. The Oxyhydrogen Blowpipe. Dry Hydrogen will not combine
with Dry Oxygen. The Reducing Power of Hydrogen. Compounds
of Hydrogen with Other Metals. Hydrogen Present in All Acids. '
Nascent Hydrogen. Certain Physical Properties of the Element
Hydrogen. The Liquefaction of Hydrogen. Can the Absolute Zero
be realized Experimentally? Properties of Liquid Hydrogen. The
Hydrogen Spectrum. Electrolysis of Hydrogen.

CHAPTER V
Water and Hydrogen Dioxide 46

Occurrence of Water. Water as it occurs in Nature is Impure.
Mineral Waters. Purification of Water. Filtration. Water not an
Element, but a Compound. Composition of Water. Chemical Be-
havior of Water. Water a Stable Compound. Physical Properties
of Water. Boiling-point Heat of Vaporization. The Freezing of
Water. Heat of Fusion of Ice. Heat of Condensation of Steam and
of Solidification of Water. Superheating and Supercooling of Water.
The Vapor-tension of Water in its Different States of Aggregation.
The Temperature-pressure Diagram of Water. The Phase Rule.
Other Physical Properties of Water. Solvent Power of Water. Un-
saturated, Saturated, and Supersaturated Solutions. Limited and
Unlimited Solubility. Properties of Water affected by Dissolved Sub-
stances. The Dissociating Power of Water.

Hydrogen Dioxide. Preparation and Purification. Properties of
Hydrogen Dioxide. Hydrogen Dioxide a Good Oxidizing Agent.
Hydrogen Dioxide also a Reducing Agent Catalytic Decomposition
of Hydrogen Dioxide. Relations of Water and Hydrogen Dioxide.

CHAPTER VI
Determination op Relative Atomic Weights

Combining Numbers and Atomic Weights. Chemical Methods of
determining Combining Numbers. Molecular Weights determined
from the Densities of Gases. Avogadro's Hypothesis. Avogadro's
Hypothesis and Molerular Wcishts. Atomic Weights from Molecular
Weights. Atomic Woichts from Specific Heats. Isomorphism an
Aid in determining Atomic Weights. Most Accurate Method of
determining Atomic Weights. Table of Atomic Weights.

CONTEXTS Xl

CHAPTER VII

TAQM

Determination op the Molecular Weiohts of Oases and op Dis-

soLTBD Substances 82

Densities and Molecular Weights. Method of Dumas. The Method
of Gay-LuH8a4\ ilofinann^s M<Miitication of the Gay-Lussac Method.
The Gas-displacement Method of Victor Meyer. Method of Bunsen*
Results of Vapor-density Measurements. Abnormal Vapor-densities,
Apparent Exceptions to the Law of Avogadro. Explanation of the
Abnormal Vapor-densities. Dissociation of Vapors diminished by an
Excess of One of the I'roilucts of Dissociation.

The Iaiw of Mass Artion. The Work of Guldberg and Waage.

Molecular Weights of Dissolved Suhstqners, lK>termination of the
Molecular Weights of Dissolved Substanct^s by the Freezing-point
Method. Apparatus devised by Beck man n. Method employed by
Beckmann. Determination of the Molecular Weights of Dissolved
Substances by the Boiling-point Method. Boiling-point Method of
Beckmann. Carrj'ing out a Molecular Weight Determination with
the Beckmann Apparatus. Boiling-iK)int Apparatus of Jones.

CHAPTER VIII
Osmotic Pressure and the Theory of Electrolytic Dissociation . 100

Osmotic Pressure. Demonstration of Osmotic Pressure. Morsels
Method of preparing Semipermeable Membranes. Measurement of
Osmotic PresHure.

Relations between Osmotic Pressure and Oas-pressure. Boyle's
Law for Osmotic Pressure. Gay-Lussac's Law for Osmotic Pressure.
Avogadro's I-»aw applied to the Osmotic Pressure of Solutions. Causes
of Gas-pressure and Osmr)tic Pressure. Exceptions to the Applica-
bility of the Gas Laws to c>sninti(! Pressure.

Origin of the Theory of Klertrolytic Dissociation. The Problem as
it was left by Vaift Iloff. Work of Arrlienias. The Theory of Elec-
trolytic Dissociation. Measurement of PLlectroIytic Dissociation. The
Conductivity Method. Method of measuring the Conductivity of
Solutions. Calculation of the Dissociation from Conductivity Meas-
urements.

CHAITER IX

Chlorine 116

Chlorine an Element or a Compound. Occurrence and Preparation
of Chlorine. Chemical Properties of Chlorine. Action of Chlorine on
Hydrogen. Action of Chlorine on Water. Action of Chlorine on
Certain Organic Compounds. Chlorine Hydrate. Certain Physical
Pn.)i>crrtie8 of Chl«)rine. Liquefaction of Chlorine. Comparative In-
activity of Dry Chlorine. Hydrr»<hloric Acid. Volume Relations in
which Hydrogen and Chlorine combine. Prejiaration of Hydrochloric

xu

CONTENTS

Acid. Chemical Properties of Hydrochloric Acid. TJeflnRjoti of an
Acid. DeUjction of Hydrochloric Acid. Physical Properties of Hy-
drochloric Acid. A^jueous Soloiioa of Hydrochloric Add. Thermo-
chemistry of Hydrochloric Acid.

Compounds of Chlorine wUk Oj^^gtn and Hydrogen. Compounds
of Chlorine wiUi Oxygen, Comp4>uiid!j of CUIorint! with Oxygen and
Hydrogen. Hypochlorous Acid. Fropcrtjeu of Hyptjchloroua Acid.
Calcium Hypochlorite. Chlorine Monoxide. Chloric Acid. Proper-
ties of Chloric Acid. Chlorates. The Chlorine Ion and the Ion of
Chlorates. Perchloric Acid, Properties of rerchloric Acid, Chhirine
Seploxide. Chlorine Dioxide and Chloroita Acid. Power of Chlorine
to combine with Oxygen* Valency, Faraday's Law the Baals of
Chemical Valence.

CHAPTER X
The Periohic System

Hypothesis of Prout. The Triads ot Doberelner. The Octaves of
Newlanda. The Periodic System of Mendel^ff, Lothar Meyer, and
Brauner, Chemical Pr*»perties and Atomic Weights. Combining
Power. Relations within the Groups. Basic and Acid Properties.
Physical Properties and Atomic Weights. Atomic Volumes. Old
Atomic Weights corrected and New Elements predicted by Means
of the Periodic System. Imperfections in the Periodic System* Gen-
eral Scheme to be followed.

136

CHAPTER XI

BaOMIItS, lODtXS, PLtJOfitKK .

Brom(ney Occurrence and Preparation, Cliemical Properties of
Bromine, Detection of Bromine. Bromine Atoms and Broujinc Ions.
Physical Properties of Bromine. Hydrobrotiiic Acid, Proj>eriiea of
Hydrobromic Acid. Compounds of Bromine with Oxygen and Hydro-
gen, Bronuc Acid. Compound of Bromine with Chlorine.

hidinr^ Occnrrence and Preparation. Chemical Properties of Iodine.
Detection of Iodine. I>etection of lodii^e in the Presence of Bromine
and Chlorine. Physical Properties of Iodine, Hydriodic Acid<
Compounds of Iodine with Oxygen and Hydrogen. Compounds of
Iodine with Chlorine. Compound of Iodine with Bromine.

Flttnrinet Occurrence and Preparation. Chemical I'roperties of
Fluorine. Physical Properties of Fluorine. Hydrofluoric Acid.
Compound of Fluorine with Iodine. Coraparisoii of the Several Acids
formed by the Halogens,

1£2

CHAPTER Xn

S^LFHUS

171

Occurrence and Purification. Chemical Properties of Sulphur.
Physical Properties of Sulphur. Vapor-density of Sulphur. The
Temperature-pressure Dla^um of Sulphur.

CONTENTS xiii

rAOB

Compoiincis of Solphar with Hydrogen. Hydrogen Sulphide.
Chemical Proiwrties of Hydrogen Sulphide. Reversible Chemical
Beactions. Acid Sulphides. Diasociation of Hydrogen Sulphide.
Physical Properties of Hydrogen Sulphide. Hydrogen Persulphides.

Compounds of Sulphur with Oxygen and Hydrogen, Sulphur Diox-
ide. Sulphurous Acid. Strength of Sulphurous Acid. Sulphur Tri-
oxide. Properties of Sulphur Trioxide. Sulphuric Acid. Chemical
Properties of Sulphuric Acid. Physical Properties of Sulphuric Acid.
Dissociation of Sulphuric Acid. Scientific and Technical Uses of
Sulphuric Acid. Other Compounds of Sulphur with Oxygen and
Hydrogen. Thiosulphuric Acid. Hyposulphurous Acid. Pyrosul-
phuric Acid or Disulphuric Acid. Persulphuric Acid. Polythionic
Acids.

Compounds of Sulphur with the Halogens and Oxygen. Com-
pounds of Sulphur with Chlorine. Compounds of Sulphur with Chlo-
rine and Oxygen.

CHAPTER Xra

8BLKirn7x AHD Tellurium 197

Selenium. Compounds of Selenium. Tellurium. Compounds of
TeUurium.

CHAPTER XrV

NmooBH 200

Occurrence and Preparation. Chemical Properties of Nitrogen.
Physical Properties of Nitrogen.

Compounds of Nitrogen with Hydrogen. Ammonia. Chemical
Properties of Ammonia. Composition of Ammonia. Physical Prop-
erties of Ammonia. Liquid Ammonia. Ammonium. Ammonium
Amalgam. Hydrazine. Properties of Hydrazine. Triazoic Acid.

CHAPTER XV

NSUTKALIZATION OP ACIDS AND BaSES 210

Ammonium Hydroxide. Bases are Hydroxyl Compounds. Acidity
of Bases and Basicity of Acids. Indicators. Theory of Indicators.
Salts. Heat of Neutralization. Explanation of the Constant Heat of
Neutralization of Strong Acids and Strong Bases. Neutralization of
Weak Acids and Bases. Explanation of the Results with Weak Acids
and Bases. Explanation of the Law of the Thermoneutrality of Solu-
tions of Salts.

Compounds of Nitrogen \oUh Oxygen and Hydrogen. Ammonium
Hydroxide. Measurement of the Dissociation of a Weak Base like

XIV CONTENTS

PAOB

Ammoniam Hydroxide. Law of Kohlrausch. Hydroxylamine.
Compounds of Nitrogen with Oxygen. Nitrous Oxide. Nitric Oxide.
Nitrogen Sesquioxide or Nitrogen Trioxide. Nitrogen Dioxide or
Nitrogen Peroxide. Nitrogen Pentoxide. Acid Compounds of Nitro-
gen with Oxygen and Hydrogen. Hyponitroua Acid. ?Jitrous Acid.
Nitric Acid. Chemical Properties of Nitric Acid. Physical Proper-
ties of Nitric Acid. Detection of Nitric Acid. Dissociation of Nitric
Acid and Nitrates. Fuming Nitric Acid. Aqua Regia.

Compounds of Nitrogen with the Halogens, Compounds of Nitrogen
with Chlorine and Bromine. Compound of Nitrogen with Iodine.
Compounds of Nitrogen with Oxygen, Hydrogen, and Sulphur. Ni-
trosyl Sulphuric Acid.

CHAPTER XVI

Tub Atmosphbrio Air and Certain Rare Elements occurring in it 235

The Atmospheric Air. Composition of the Atmosphere. Is the
Air a Mixture or a Compound ? Physical Properties of Atmospheric
Air. Liquid Air.

Argon, Helium, Krypton, Neon, Xenon. Argon. Number of
Atoms in the Molecule of Argon. Helium, Neon, Krypton, and Xenon.

CHAPTER XVII
Phosphorus 242

Occurrence and Preparation. Properties of Phosphorus. Yellow
Phosphorus. Red Phosphorus. Metallic Phosphorus. White Phos-
phorus. Compounds of Phosphorus with Hydrogen. Compounds of
Phosphorus with Oxygen and Hydrogen. The Acids of Phosphorus.
Orthophosphoric Acid. Dissociation of Phosphoric Acid. Detection
and Determination of Phosphoric Acid. Pyrophosphoric Acid. Meta-
phosphoric Acid. Hypophosphoric Acid. Phosphorous Acid. Meta-
phosphorous Acid. Ilypophosphorous Acid. Strengths of the Acids
of Phosphorus.

Compounds of Phosphorus with the Halogens. Phosphorus Tri-
chloride. Phosphorus Pentachloride. Phosphorus Oxychloride.

CHAPTER XVIII
Arsenic 255

Occurrence and Preparation. Properties of Arsenic. Compound
of Arsenic with Hydrogen (arsine), AsHs.

Compounds of Arsenic with Oxygen and Hydrogen. Compounds of
Arsenic with Oxygen. Arsenic Trioxide. Arsenic Pentoxide. Arse-
nious Acid. Arsenic Acid. Compounds of Arsenic with the Halo-
gens. Compounds of Arsenic with Sulphur. Sulpho-salts of Arsenic.

CONTENTS XV

CHAPTER XIX

AnriMorr 261

Oocarrence and Preparation. Properties of Antimony. Compound
of Antimony with Hydrogen(8tibine), SbHf

Compounds of AnUmoniff with Oxygen and Hydrogen. Oxides of
Antimony. Acids of Antimony. Compounds of Antimony with the
Halogens. Compounds of Antimony with Sulphur. Compounds of
Antimony with Sulphur and the Metals. Hard Lead.

CHAPTER XX

Bismuth 967

Occurrence and Properties. Compounds of Bismuth with Oxygen
and Hydrogen. Bismuth Chloride. Bismuth Sulphide.

CHAPTER XXI

VAiTADiuif, CoLUMBiuM, Nbodtmium, Praseodtmium, Tahtalum . . 270

Vanadium. Columbium. Praseodymium and Neodymlum. Tan-
talum.

CHAPTER XXn

Carbon 278

Allotropic Forms of Carbon. Amorphous Forms of Carbon. The
Different Forms of Carbon contain Different Amounts of Energy.
Physical Properties of Carbon.

Compounds of Carbon. Compounds of Carbon with Hydrogen.
Compounds of Carbon with Oxygen. Carbon Monoxide. Thermo-
chemistry of Carbon Monoxide. Water-gas. Carbon Dioxide. Prep-
aration of Carbon Dioxide. Chemical Properties of Carbon Dioxide.
Reduction of Carbon Dioxide by Plants. Physical Properties of
Carbon Dioxide. Discovery of Continuity of Passage from the Liquid
to the Gaseous State. The Kinetic Theory of Liquids. Carbon Sub-
oxide. Compounds of Carbon with Oxygen and Hydrogen. Com-
pounds of Carbon with the Halogens. Compound of Carbon with
Sulphur. Compound of Carbon with Nitrogen (cyanogen). Hydro-
cyanic Acid. Cyanic and Sulphocyanic Acids.

The Role of Carbon in producing Light. Illumination. Candle
and Oil-lamp. Coal-gas, Water-gas. Flames and their Luminosity.
Bunsen Burner, Blowpipe. Effect of cooling the Flame. The Acety-
lene Light. The Welsbach Light. The Electric Light. Measurement '
of the Relative Intensities of Light.

XVI CONTENTS

CHAPTER XXm

Silicon

The Element Silicon. Silicon Hydride or Hydrogen Silicide. Sili-
con Dioxide. The Acids of Silicon. Dialyzer. Crystalloids and
Colloids. Metasilicic Acid. Polysilicic Acids. Conversion of Silicates
into Carbonates. Compounds of Silicon with the Halogens. Com-
pound of Silicon with Carbon — Carborundum.

CHAPTER XXIV

Gbrmaitiuii, Titanium, Zirconium, Cerium, and Thorium • • . 906
Germanium. Titanium. Cerium. Thorium.

CHAPTER XXV

BoxoN 807

Occurrence, Preparation, and Properties. Boron Trioxlde. Boron
Nitride Compounds of Boron with Other Elements. Summary.

CHAPTER XXVI
Thb Mitals 810

CHAPTER XXVn

Thb Alkau Mrtals — Lithium, Sodium, Potassium, Rubidium, and

Cjcsium 812

Occurrence of the Element Sodium. Preparation of Sodium. Prop-
erties of Metallic Sodium.

Compimmh of Sodium with Oxygen and Hydrogen. Sodium Hy-
dride. SiHliuui IVroxide. Sodium Hydroxide. The Chemistry of
Sixlium the Chemistr>' of the Sodium Ion. Compounds of Sodium
with the Halogens. Sodium Chloride. I\irification of Sodium Chlo-
ride. S(xlium Bromide, Sodium Lxiide. Sodium Hypochlorite,
Chlorate and Bromate. Sodium Triazoate and Sodium Amide.
Sixlium Nitrate. Sodium Nitrite. Sodium Hydrosulphide and Sodium
SulphitieH. Sodium Sulphite. Sodium Sulphate. Acid Sodium Sul-
phate and Stxiium l^-rosnilphate. Sodium Thiosulphate. Sodium
Carlwnate. The Ix» Blanc Method. The Solvay or Ammonia Process.
Acid S<xiiuni Carbi^nates. Hydrolysis of the Carbonates. The Plios-
phates of S<xiiuni. Sodium Ammonium Phosphate. Sodium Borate
or Tetraborate. S*xiium Silicate. The Sodium Salt of Pyroanti-
monius Acid. S^xlium Acetate. Sodium Cyanide. Spectrum of
SoiUum.

CONTENTS XVll

CHAPTER XXVin

PAOB

Potassium 343

Occurrence and Preparation. Properties of PotaBsiom. Potassium
Hydride. Potassium Peroxide. Potassium Hydroxide. Compounds
of Potassium with the Halogens. Potassium Chloride. Potassium
Bromide. Potassium Iodide. Potassium Fluoride. Potassium Chlo-
rate. Potassium Perchlorate. Potassium Hydrazoate and Potassium
Amide. Potassium Nitrate. Gunpowder. Potassium Nitrite. Com-
pounds of Potassium with Sulphur. Compounds of Potassium with
Sulphur and Oxygen. Potassium Sulphate. Potassium Carbide.
Potassium Carbonate. Acid or Primary Potassium Carbonate. Phos-
phates of Potassium. Silicates of Potassium. Potassium Silico-
fluoride. Potassium Pyroantimoniate. Potassium Cyanide. Potassium
Sulphocyanate. Oxalates of Potassium. Detection of Potassium.

CHAPTER XXIX

LrraiuM, Rubidium, Cesium, and Ammonium 352

Lithium, Discovery, Preparation, and Properties. Compounds of
Lithium. Rubidium, Occurrence, Preparation, and Properties. Com-
pounds of Rubidium. Cassium, Occurrence, Compounds. Ammo-
nium, Ammonium Hydroxide. Ammonium Chloride. Ammonium
Hydrazoate or Triazoate. Ammonium Nitrite. Ammonium Nitrate.
Ammonium Hydrosulphide, Sulphide, and Polysulphides. Ammonium
Sulphate. Ammonium Carbonate. Phosphates of Ammonium. Char-
acteristics of the Alkali Metals in General.

CHAPTER XXX

Calcium, Strontium, and Barium

Calcium, Occurrence, Preparation, and I*roperties. Calcium Oxide
or Lime. Calcium Hydroxide or Slaked Lime. Compounds of Cal-
cium with the Halogens. Calcium Hypochlorite — Bleaching Powder.
Sulphides of Calcium — Calcium Hydrosulphide. Calcium Sulphate.
Calcium Carbide. Calcium Carbonate. Primary or Acid Calcium
Carbonate. Phosphates of Calcium. Calcium Silicate. Glass.
Varieties of Glass. Calcium Oxalate. Detection of Calcium.

Strontium. Occurrence, Preparation, and Properties of Strontium.
Salts of Strontium. Detection of Strontium.

Barium. Oxides of Barium. Barium Hydroxide. Barium Chlo-
ride. Barium Sulphate. Barium Carbonate. Barium Phosphates.
Other Insoluble Compounds of Barium. Detection of Barium. Detec-
tion of the Alkaline Earths — Calcium, Strontium, and Barium.

xvm CONTENTS

CHAPTER XXXI

PA«B

Thb Maombsium Group — Glucinum, Maonbsidm, Zinc, Cadmium,

Mbrcubt 888

Olucinum. Magnesium. Magnesium Oxide and Magnesium Hy-
droxide. Magnesium Cliloride. Magnesium Sulphate. Magnesium
Carbonate. Phosphates of Magnesium. Silicates of Magnesium.
Other Compounds of Magnesium. Separation of Magnesium from the
Elements of the Calcium Group.

Zinc. Zinc Oxide and Hydroxide. Zinc Chloride. Zinc Sulphide.
Zinc Sulphate. Zinc Carbonate.

Uses of Zinc in Primary Batteries, Demonstration of the Solution-
tension of Metals. The Relative Solution-tensions of Some of the
More Common Metals. Solution-tension of Metals and Primary Cells.
Concentration Element. The Daniell Cell.

Cadmium. Salts of Cadmium.

Mercury. Properties of Mercury. Amalgams. Molecular Weights
of Metals in Mercury. Mercurous and Mercuric Oxides. Mercurous
and Mercuric Chlorides. Mercuric Bromide and Iodide. Mercuric
Sulphide. Mercurous and Mercuric Sulphates. Mercuric Cyanide.
Action of Ammonia on Salts of Mercury. Variable Valence.

CHAPTER XXXn

Thb Eabth Mbtals — Aluminium and the Rabb Elements, Scan-
dium, Gallium, Yttrium, Indium, Lanthanum, Ytterbium, Thal-
lium, AND Samarium 407

Aluminium^ Occurrence and Preparation. Properties of Aluminium.
Alloys of Aluminium. Aluminium Amalgam. Aluminium Oxide.
Aluminates. Aluminium Chloride. Aluminium Sulphide. Alu-
minium Sulphate, 'llie Alums. Aluminium Carbide and Carbonate.
Silicates of Aluminium. Applications of Aluminium Silicates. Porce-
lain. Detection of Aluminium.

Scamlium. Gallium. Yttrium. Indium. lAznthanum. Ytter-
bium. JTiallium. Samarium.

CHAPTER XXXm

Iron 410

Iron, Occurrence and Preparation. Properties of Iron. Impure or
Commercial Iron. The Thomas-Gilchrist Converter. Oxides of Iron.
Ferrous and Ferric Compounds. Ferrous and Ferric Hydroxides.
Ferrous and Ferric Chlorides. Sulphides of Iron. Ferrous Sulphate.
Ferric Sulphate. Potassium Ferrocyanide. Copper Ferrocyanide.
Potassium Ferricyanide. Change in Color with Change in Electrical
Charge. Other Salts of Iron. Ferrates.

CONTENTS XIX

CHAPTER XXXIV

TAQM

Cobalt ahd Nickel 431

Cobalt. Cobaltons and Cobaltic Compounds. Oxides and Hydrox-
ides of Cobalt Cobaltoos Salts. Double Cyanides of Cobalt. Double
Nitrite of Cobalt. Action of Ammonia on Solutions of Cobalt Salts.

Nickel. Compounds of Nickel.

CHAPTER XXXV
Manoanesb 4S6

Oxides of Manganese. Hydroxides of Manganese. Manganous
Salts. Manganic Compounds. Tetravalent Manganese. Valence
and Properties of Manganese. Manganous Acid. Manganic Acid.
Permanganic Acid. Potassium Permanganate. Color of Perman-
ganates.

CHAPTER XXXVI
Chromium 445

Oxides of Chromium. Hydroxides of Chromium. Valence and
Properties of Chromium. Cbromoas Salts. Chromic Salts. Chromic
Chloride. Chromites. Chromic Acid. Chromates. Dichromates.

The Ions Cr04 and CrsO?. Chlorides of Chromic Acid. Perchromic
Acid. Detection of Chromium.

CHAPTER XXXVn
Molybdenum, Tunosten, Uranium 458

Oxides of Molybdenum, Molybdic Acid. Compounds of Chlorine
with Molybdenum.

Tungsten. Chlorides of Tungsten. Tungstic Acid.

Uranium. Oxides of Uranium. Chlorides of Uranium. Uranium
Radiation. Other Radiactive Substances.

CHAPTER XXXVIII

Copper 460

Occurrence and Preparation of Copper. Properties of Copper.
Alloys of Copper. Oxides of Copper. Cupric Hydroxide. Chlorides
of Copper. Cupric Chloride. Sulphides of Copper. Copper Sulphate.
Copper Carbonate. Other Copper Salts.. Precipitation of Copper
by Zinc. Another Method of Ion Formation.

XX CONTEXTS

CHAPTER XXXIX

PAGS

Silver and Gold 467

Preparation of Silver, Properties of Silver. Colloidal Silver.
Alloys uf Silver. Silvering. Oxides and Hydroxide of Silver. The
Silver Ion. Silver Ciiloride. Silver Bromide. Photography. Silver
Iodide. Silver Nitrate. Silver Sulphide. Silver Sulphate. Silver
t 'arbonate. Other Compounds of Silver.

Gold. Metallurgy of Gold. Properties of Gold. Oxides and
Hydroxides of Gold. Salts of Gold.

CHAITER XL

Lbai>, Tik 478

Occurrence, Preparation, and Properties of Lead. Precipitation of
Lead by Metaln. Oxiiies of Lead. Hydroxides of Lead. Chlorides
of Lead. Iodide i>f lA*ad. Ix*ad Nitniti^ I^ad Sulphide. I^ad
Sulphate. Lead Persulphate. L<*ad Carbonate. Lead Chromate.
Lead Acetate. The Storage Battery or Accumulator.

Tin. Preparatit)n au<l Properties of Tin. Allotropic Forms of Tin.
Alloys of Tin. The 'I'in Ions. Stannous and Stannic Oxides. Stan-
nous and Stannic Hydroxides. Stannous Chloride. Stannic Chloride.
Sulphides of Tin.

CHAPTER XLI

RuTHEifiuii, RHODirM, Palladiitm, 0smii:m, Iridium, Platinum . . 489
Ruthenium. Rodiutn. Palladium. Osmium. Iridium. Plati-
num. Projjerties of Platinum. Uses of Platinum. Colloidal Solution
of Platinum. Oxides and Hydroxides of Platinum. Chlorides of
FlaUnnm. Sulphides of Platinum. Double Cyanides of Platinum.

PRINCIPLES OF
INORGANIC CHEMISTRY

PRINCIPLES OF INORGANIC CHEMISTRY

CHAPTER I

onraoDucnoN

Stiiy of Vafene. — The study of nature is not limitiMl to tlm
world in which we lire, but to the universe of wlnoli our world
forms only a reiy small part The study of the variouN iu«|MM*tN of
naturey or natural science, forms the greatest and nioHt conipnOmti
sive chapter of human knowledge. Indeeil, so great In iho lli«lil nf
natural science and so different the methods whioh nm fMiiplnyil
in studying nature, that no one mind can compreht^nd morn tliiiii ii
small part of what has already been learned.

Natural science can, however, be subdivided into a nntiiU*r of
branches, which are all related, but which )KmH<fNN liilinnMit dtlTfi
ences sufficient to distinguish the one from tho othnr^ and In Anmn
cases these differences are quite marked.

Astronomy has to do chiefly with the study of ilm inoiiniin itiid
relations of bodies as a wholain the univorHi', whil«« idiynlin, rtifin
istry, geology, and biology are concerned primarily wif li plMiMuni'ihi
which manifest themselves on the earth. Of tlH*Ni<, pliyniin iind
chemistry are far more closely allied than ilM»y iim to K*'ol'»^fV I Imi
science of the formation of the surface of tlin nirtli, or lo liioloj/y
the science of living matter.

Selations between Chemistry and Physios. Wliiln it \n imponti)
ble at this stage to give a compreh<7nHivi*' conci'ptioti of ttin ifhif iohd
and differences between chemistry and )>liynif:H, riTtaiii fiindainKotuI
distinctions can be pointed out.

Connect a piece of copper with an idcciric l*fitt«iry, iind tin v\tu'\nf
current will flow through it. The coppor wliil« rarry hik tlm rm i^nt
has properties which are different from <opp<*r ilironj^li whirl) no
electricity is passing. Disconnect ilio copp<*r from tli« «d«M;triij l»jit
tery, and it possesses its original pniperti<*».

Heat a piece of copper gently, and Kom<t nf iin pro)H^rti»'« an*
changed. It will give out heat to surrounding^ t)h'yu'XH\ it will
B 1

2 PRINCIPLES OF INORGANIC CHEMISTRY

occupy a larger volume when hot than at ordinary teraperatnrea*
Allow the copper which has been warmed to cool, and it will possegg
again its original properties. Hammer the copper or bend it, and it
will remain copper* Changes of this kind are known as physical.

If, on the other hand, we heat a piece of copper to rt^ness in
the presence of the air, a far more fundamental eliange takes place.
The copper disappears and a black snbst-ance is formed which has
properties quite different from the original copper. The black sub-
atance does not look like a metal. It cannot be drawn out into
wire. It weighs considerably more than the original cop[ier, and in
general has properties sufficiently different from the original copper
to show that we are dealing with an entirely different substance.
If the black powder is now cooled to the original temperature of the
copper, it retains its own characteristic properties.

It is obvious, therefore, that by heating to a sufficiently high
temperature in the presence of the air, the copper has been trans-
formed into something else, and that the new substance is not
rctransformed into copper when the original temperature is again
restored.

The change effected in the latter case is, then, far more funda^
mental than in the former. While certain proi>erties of the copper
were changed by passing an electric current through it^ or by gently
warming it, these properties were restored again when the current
was interrupted, or when the temperature was allowed to fall. The
original copper remained copper. In the latter case, however, the
composition of the substance was changed^ and this is characteristic of
chemical activity. The Buhstances which react chemically lose many
of their characteristic properties, and give rise to new substances
with very different properties.

The distinction between physical and chemical change is not always
as sharp as in the example given above. Certain phenomena mani-
fest themselves which belong, strictly speaking, neither to chemistry
nor to physics, but occupy a position midway between the two, A
comparatively new branch of science which deals with these phe-
nomena has come into prominence in the last fifteen years. This is
known as physical chemistry.

Although the distinction between physical and chemical changes
ifl not always a sharp one^ yet, in raost cases, there is no serious diffi-
culty in deciding to which class a given set of phenomena belongs.
In general, any change which does not affect the composition of sub-
stances is physical, while change in compOBition is characteristic of
chemical trans formation a.

INTRODUCTION 8

Blements and Componndf. — If we look about ns, we recognize that
nature is made up apparently of a great many substances. The soil
and the rocks differ greatly in composition in different localities, and
are always more or less complex. Water exists everywhere, and the
air is a mixture of many substances. When we turn to living matter
we find the complexity greatly increased. The simplest living being
is composed of very complex substances, and the more highly devel-
oped organisms contain a countless number of substances.

This is the way the problem of the com|>osition of the external
world, as recognized by our senses, presents itself at first. It, how-
ever, becomes greatly simplified when we study the composition of
things in a systematic and comprehensive manner. All known sub-
stances fall into two great classes, — those that cannot be decom-
posed into simpler substances and those that can. Take the piece
of copper already referred to. By no process known to man can it
be decomposed into anything simpler than copper. It can be caused
to unite with other things and form substances more complex than
copper, but deal with it as you will, and it cannot be decomposed into
anything else. On the other hand, take the well-known substance
water, add a little acid to it to make it conduct, and pass an electric
current through it. The water will be decomposed into two simpler
substances, both of them gases, and they will be set free the one at
the one pole, the other at the other.

There is, therefore, a fundamental difference between copper and
water — the one cannot be decomposed into simpler substances, the
other can be decomposed into two substances, both of which differ
fundamentally in their properties from the substance water.

Substances like copper which cannot be decomposed into any-
thing simpler are known as dementSy while those substances which
can be decomposed into simpler things are known as compounds.

The Number of Elements and Compounds. — W^hile the chemical
compounds already known number more than a hundred thousand,
the number of chemical elements which have thus far been discov-
ered is only about seventy-five. When we consider that a compound
is made up of two or more of these elementai-y substances, the whole
problem of the composition of substances is vastly simplified. We
can, then, refer every compound known, both in inanimate nature
and in the realm of living matter, to a comparatively few elementary
substances.

Take the rocks which are most familiar on the surface of the
earth ; they are made up chiefly of not more than a dozen elementary
substances. In addition to this dozen elements they may contain a

PRINCIPLES OF INORGANIC CHEMISTRY

number of other elementary substances, but these are present in rela^
tively araall quantities-
Water, which covers such a large portion of the surface of the
earth, is, as we have seen, made np of two elenieuts.

The atmosphere which is so essential to life is made up chiefly of
two elements^ containing, however, a number of other elements and
compounds in relatively small quautities.

If we turn to living matter, we find a very large number of chem-
ical compounds, and the greatest complexity represented. Cellulose,
starch, albumen, are among the most complex substances known to the
chemist, yet an analysis of these substances brings out the surprising
fact that they eontam scarcely more than a half-dozen elements.

The number of chemical elements know a to us at present is, as
already stated, about seventy-five. This number has been largely
increased during the last few years, and there are good scientific
reasons, as we sball learn, for believing that elements exist which
have not yet been discovered. It should, however, be stated that if
such elements exist at all, it is highly probable that they occur only
in small quantities, or in comparatively obscure places, otherwise
they woidd have been discovered by one investigator or another
using chemical, physical, or physical chemical methods.

There is, on the other band, the probability that substances which
we now regard as elementary » may prove to be compounds of still
simpler substances, A substance is for us an element, which we
have not been able thus far to break down into anything simpler. It
is quite conceivable, how eve r^ that as old methods are improved and
new ones devised, we may be able to effect decompositions not thus
far accomplished. One would naturally turn in this connection to
electrical methods, by means of which very high temperatures can be
easily realized. Since, however, this is pure speeulation, entirely un-
substantiated thii^ frir by fact, it is not profitable to pursue it farther.
The Chemical Elemeats. — Having learned what is meant by the
term "chemical element/' we naturally ask which are the elements
and what substances are compounds? In the following tal>le the
substances which have been shown with a reasonable degree of proba-
bility to be elementary, are given, together with the symbol which
is used for the element in question ; —

I

Aluminium , Al

Antimony Sb

Argon * A

Arsenic ..,.,.. As

Barium ,...•.* Ba

Bismuth Bi

Boron .,.«.... B

Bromine ....... Br

INTRODUCTIOK

9

Cadmium Cd

Caesium Cs

Calcium Ca

Carbon C

Cerium Ce

Chlorine CI

Chromium Cr

Cobalt Co

Columbium Cb

Copper Cu

Erbium (?) E

Fluorine F

Gadolinium G

Gallium Ga

Germanium Ge

Glucinum Gl

Gold Au

Helium He

Hydrogen H

Indium In

Iodine I

Iridium Ir

Iron Fe

Krypton Kr

Lanthanum La

Lead Pb

Lithium Li

Magnesium Mg

Manganese Mn

Mercury Hg

Molybdenum Mo

Neodymium Nd

Neon Ne

Xickel Ki

Nitrogen X

Osmium Os

Oxygen O

Palladium Pd

Phosphorus P

Platinum! Pt

Potassium K

Proseodymium Pr

Rhodium Rh

Rubidium Rb

Ruthenium Ru

Samarium Sm

Scandium Sc

Selenium Se

Silicon Si

Silver Ag

Sodium Xa

Strontium Sr

Sulphur S

Tantalum Ta

Tellurium Te

Thallium Tl

Thorium Th

Thulium Tu

Tin ... Sn

Titanium Ti

Tungsten W

Uranium U

Vanadium V

Xenon X

Ytterbium Yb

Yttrium Y

Zinc Zn

Zirconium Zr

The symbol of the element is usually the first letter, or this com-
bined with some other distinctive letter when several elements begin
with the same letter. In some cases, however, the symbol is not
taken from the English name of the element, but from the Latin.
Thus, the symbol for copper is Cu, from the Latin cuprum; the
symbol for iron Fe, from the Latin ferrum; etc.

Some elements occur in very large quantities in the earth, while

6 PRINCIPLES OF INORGANIC CHEMISTRY

others are comparatively rare. The following estimate of the com-
l>ositioti of that part of the earth which is accessible to us seems on
the whole the most reliable: —

Oxygen, jjercentage in the earth 50.0

Silicon, "^ a u u 25.0

Aluminium, "' " " *' 7.2

Iron, " a u ic 5Q

Cakium, " u u u 35

Magnesium, " a u u 2.5

Soilium, " ti u u 2.3

Potassium, " i( a u 2.2

Hydrogen, " u u u jq

Titanium, " u u u o.3

Carbon, " a u a q2

The earth is thus made up chiefly of nine elements, the reijiain-
der Avurring in comparatively small quantities.

Qfetoucal Combination. — Certain elements can combine with
-"t^rrskiu othor oloments and form compounds. Two elements may
omoLue JLud form a compound, or three, four, or more elements may
•ymoiUtf. We may, therefore, have compounds containing two, or a
■^a^ii ;u>[cer uumU^r of fslementary substances. Most chemical com-
ouuus A>ucuiu two or three elements, but some contain four, five, or
'•ju r -veu ^ larger number of elements.

■ Vhue .-ieuiencarv suKstances may combine with one another and
;ntt •jiui.»\jaud*» it is not true that any element can combine with
.iiy .'iiiet- -lemeuc. We shall learn that elements with widely dif-

iwu ;>rf.*Lit-ffite6'^aerally combine most readily, while the elements
•-.uae i-rvwecue* aw similar may not combine at all, or if they form
luiuuiULa. ars« .iw often readily decomposed.

'usa t«.«a&euu ■ouibiue and form a compound, this is expressed

■ ir.mm '.ne -vmooi* oi the elements with a plus sign between them
-^ iviMiUui -iii^ »f the eijuality sign, and the symbols of the

-.bUL acer uco the compound on the right-hand side
^l lauun. T^ius* '«bea the elements hydrogen and oxygen
L ai xniL *:*wr. "in* '-* expresseil chemically as follows ; —

I I 11,1,1 1:^ ^a>m J» 1 L-kemical equation.

j^jw.^ ir^sflittcrv • ousists^ in part, of a study of the ele-

--^ ^^ ^.^n^n.1^ jruiMSL Aew elcmcnts can form with one

.wt . -rmaB^skiasaatt^ iafcve been reached to which rhomical

A ^M^ewM:^^^tt»*f juqrb. To these we shall now turn.

CHAPTER II

The Sdenoe of (Slmustrj. — The stodj of chemical phetkraaffn^
like the study of phenomena in general, coDsisted at first in simple
observation and description. Two substances vere aIIov«d to r«act
chemically and the reaction was obserrecL The nator*; and properties
of the substances entering into the reaction were studied* abd then
the nature and properties of the products formed. Thij wa.% the
qualitative stage of chemistry.

Mere qualitative observations are followed by quantitative m^rav
uremeuts in the development of any branch of natural science. The
mind is not content with merely observing phenomena at lon;^ range,
as it were. It desires to study them in detail and quantitatively,
and this marks the second stage in the development of a branch of
science. The quantities of the substances which enter into a reaction
were carefully weighed, and the quantities of the pr^xiucts forroerL
The amounts of different substances which combine with one am/ther
were determined, and certain other changes which take place simul-
taneously with chemical transformations were studied.

Just as the qualitative stage in the development of any branch
of science leads to the discovery of a large number of facts, so the
quantitative period brings to light an enormous mass of details
which are placed upon record. Still we do not have a science. A
heterogeneous mass of isolated facts, however large and however well
established, is not a science. Indeed, the highest aim of a scienre
is not simply to observe and record facts.

Facts bear about the same relation to a science as the bricks to a
magnificent piece of architecture. They are absolutely essential to
it, but they are only a means to an end.

OeneralizatioiL — The highest aim of scientific investigation is
the discovery of wide-reaching relations between large numbers of
facts. Such relations when sufficiently comprehensive are known as
generalizations. Beyond these we cannot go. Whether they are
absolute truths to which all phenomena conform we cannot say,
because we cannot observe all the cases to which they apply. Take

7

8 PRINCIPLES OF INORGANIC CHEMISTRY

a simple example by way of illustration. It was early observed that
a body thrown upward from the surface of the earth will return
again to the surface. Repeated observations confirmed those first
made, but it remained for Newton to arrive at the generalization
known as the law of gravitation.

In a similar manner certain generalizations have been reached in
chemistry, which have been of fundamental importance in the de-
velopment of the science. Some of these will be considered in this
place, while others will be introduced in connections into which they
seem to enter naturally.

The Law of the Conservation of Kau. — We have seen that when
substances react chemically they disappear as such, and products are
formed having properties very different from the original substances.
When copper was heated in the air a black powder was formed
having properties which are very different from the original copper.
The question arises, are all the properties of the copper lost during
thd ahciiiical transformation, or have only some of them disappeared?
It 1h (^asy to convince ourselves that most of the properties of the
iMippor have been lost during the reaction, but it is a very much
tiioi'H (litttcult problem to determine whether all the properties have
itiMi4|t|marml. Take the property mass. Does the mass of the sub-
liiHtUHtM tintt^ring into a chemical reaction undergo any change during
tho ttmotloii ? This is a question very easy to raise but very difficult
Iri* miMWiir with any high degree of accuracy.

HituHi wnl^ht is a measure of mass, the problem reduces itself to
(luUsi'uiiuihii whothor there is any change in weight under similar
imuUiUuim wImmi oh(niii(Uil reaction takes place ? The weighings must
i»4t 14444(0 wwkU^v Miinilar conditions, before and after the reaction, since
tUis wi>i||h|i itf any llivun substance is a function of the conditions,
t»«|HH^i4il> Uu4 UiMtunoH from the centre of the earth — a body weigh-
ii4^ U4tau (44 4 iliH^p vallt^y than on a high mountain.

VVt* iu^u 4414* wor thin «i\u»Mtion then only to within the limit of ac-

• ■unkr.v 4»t tlu» i44oMt liitlhiMl chemical balance, and some of the most
.i> ( tii»vU> wiuk iu ihtt whole field of chemistry has been done in

• Mttui t Uu4i ^iih ihU )irotilom.

U Ih.^^ kHH»u ^MilH^iliiihiMl for a comparatively long time that if
Hull, im .u4> 4ih44»4|ji» ill maMN in chemical reaction it is very small.
I l»u, U.MUv^iu, Uh*\\>« vutiivly unanswered the question as to whether
i !»«♦.. .uikiUi. i4v>t> \h> 44 Wm/A* ohango in mass when substances react

\ I*. , .|n. .ii^»u Kkt 4AHV»4llY \h»ou Htudiecl with a degree of accuracy
^ lu. U i*.i . ^.i\\\\ k V44 .^n»uo^im»twl and never surpassed in the whole

GEXERAUZATIOXS 9

history of chemical investigatioiL The German i>hvsical chemist,
Landolt, of the University of Berlin, had constructed probably the
most accurate chemical balance which has ever l)4vii niaile. With
this he weighed the substances before the reatHion and tlun Wfij,'he*l
the products of the reaction. Although verj- slight ililT«*ri*nces in
weight were detected, yet these were in no ease grfat»-r than the
possible experimental error. His work and sul>seqii»'nt in vest i ora-
tions along the same line confirm the earlier concludiun that there
is no appreciable change in weight, and, consequently, no appreciable
change in mass in chemical reaction. This is known a.s the I'tic of
the consercntion oftnass.

The importance of this generalization for the scienre of chemis-
try cannot be easily overestimated. If mass did chrin^'*' in rhfiuical
reaction, it would be meaningless to work quantitatively where
chemical transformations take place. The whole sul)j»M-t of quanti-
tative analysis would be very different from what it is today, and an
exact science of chemistry would be next to impossible.

The law of the conservation of mass is sometini»»s ref»Trerl to as
the law of the conservation of matter. The former ♦•xpn-ssi.m is
greatly to be preferred to the latter, since it states just what has
been established by experiment The latter is i»ure tlu'ory, having
no known connection with fact.

The Law of Constant Proportion. — The second iini>ortant i;ener-

alization which was reached through the quantitative study of rjieini-

cal phenomena, was that the constituents of a choinical rompouncl

are always present in a constant proportion. If two substances reaet

{ chemically and form a third, they enter into combination in a con-

I stant proportion. The law may be formulated thus : —

I Any given chemical compound always contains the satn^ constituents,

and there is a constant proportion between the rnasses of the constituents
present.

The law of constant proportions was called in question in the
early years of the nineteenth century by the Frencli chemist,
Berthollet, in his great book, Essai de statique chiniique. Berthollet
was deeply impressed by the effect of the quantity of substance used
on the nature of the chemical reaction, and saw in outline what we
shall learn to be one of the most important laws of chemical activity.
He thought that not only the nature and magnitude of the reaction
were aifected by the masses of the substances used, but also the
composition of the products formed. Two substances could unite in
a great many proportions, and the composition of the product de-

10 PRINCIPLES OF INORGANIC CHEMISTRY

pended chiefly on the relation between the amounts of the substances
used.

The error of BerthoUet was corrected by Proust, who showed
that many of the substances which BerthoUet had supposed to
be compounds were mixtures of different substances. This, how-
ever, is not a severe reproach to BerthoUet, since the methods
for effecting separations and analyzing substances were very crude
indeed, at his time, and it arouses our deep admiration when we
consider what was accomplished under the conditions which then
jptevailed.

Subsequent work with the more refined methods has shown that
ithe law of constant proportion is a fundamental law of chemistry.
We should mention especially the classical work of the Belgian,
Stas. He tested this law with a thoroughness and accuracy which
have rarely been equalled in any branch of science. The result is
what has already been indicated. The law has stood the most
refined and crucial experimental test.

The Law of Koltiple Proportions. — While it is true that sub-
stances combine in constant proportions, it is also true that two
substances may combine in more than one proportion. The two
compounds methane and ethylene were analyzed, and it was
found that the ratio of carbon to hydrogen in the former was as
3 to 1 ; in the latter as 6 to 1. The latter evidently contains twice
as much carbon with respect to hydrogen as the former. A number
of analogous cases where two elements combine in more than one
proportion were examined, and the result was the discovery by
Dalton of the law of multiple proportions. This law may be formu-
lated thus : —

Iftico elements combine in more than one proportion, the masses of
the one which combine with a given mass of the other bear a simple,
rational relation to one another.

Since this law was proposed, great masses of facts which bear
upon it have been discovered. The result is that the law has been
found to hold thus far without an exception.

The importance of the law of multiple proportions for the science
of chemistry is very great. If this law did not govern chemical
reactions, the number of compounds which any two elements might
form with one another would be very great. As it is, any two
elements can form only a limited, and usually a comparatively
small, number of compounds with each other.

One further point should be mentioned in connection with this

GEXERALIZATIOXS 11

law. It shows that chemical reactions proceed br stej'S or li^ps, as
it were. One part of A combines with one of B. or with two of B,
or with three of B; not one part of A with one and a frar-ti^n of B,
or two and a fraction of B. The importance of this fa*-t as bearing
upon the possibility of an exact science of chemistry is very great
indeed. While it is impossible to see its sii^ificance at this stage
of our subject, it may be stated that it is this law more tlian any
other which, for a long time, made it difficult to apply mathematir-s
to chemical phenomena. These breaks, or latk of continuity, made
it extremely difficult to use the calculus in dealing with chemical
phenomena as it could be used in dealing with the phenomena of
physics, and are the most potent reason why chemistry has developed
so much more slowly than physics, and is still, strictly sjieaking. not
an exact science.

The Law of Combining Weights. — There is a thirrl law to which
the masses of substances which combine with one another conform.
If we determine the weights of different substances which combine
with a given weight of a definite substance, these weights, or simple
multiples of them, represent the quantities of the different sub-
stances which will combine with one another. Thus, '15.4.5 ]>arts
of chlorine combine with 1 part of hydrogen, and 79.90 parts of
bromine combine with 1 part of hydrogen. When chlorine and
bromine combine, SoAo parts of chlorine combine with 79.% parts of
bromine. Again, 40.1 parts of calcium combine with 10 parts of
oxygen, and 65.4 parts of zinc combine with 16 parts of oxygen. If
calcium and zinc combined^ 40.1 parts of calcium would combine
with 65.4 parts of zinc.

The quantities of substances which combine with one another have
been termed their combining numbers or combining weights, and the
law is known as the law of combining weights. The law may be
stated thus : —

Substances combine either in the ratio of their combining numbers^
or in simple, rational multiples of these numbers.

Of all the elements, hydrogen combines with other elements in
smaller quantity by weight than any other element. Its combining
number, being the least of all the elements, is taken as unity. We
shall become familiar with the combining weights of the elements in
another connection. Suffice it to say here that this law, like the laws
of constant and multiple proportions, has been subjected to the most
careful experimental test, and has been shown to be true to within
the limit of error of some of the most refined experimental work.

12 PRINCIPLES OF INORGANIC CHEMISTRY

The Atomic Theory. — The discovery of empirical relations such
as the three laws of chemical combination just considered, is of great
importance, and is absolutely essential to scientific progress; but
these are of interest chiefly as they lead to correct theories and
wide-reaching generalizations. Dalton raised the question. What do
the laws of definite and multiple proportions really mean? Why
do such relations obtain ? His answer is what has come to be known
as the scientific atomic theory, in contradistinction to the older
imaginative speculations about atoms and molecules. The view that
matter is composed of indivisible particles or atoms, which have
definite weights, and that chemical action takes place between these
particles, was to Dalton the only rational explanation of the laws of
multiple proportion and combining weights. If water is composed
of such ultimate, indivisible parts or atoms, then a constant number
of atoms of one substance combines with one atom of another sub-
stance to form a definite molecule of the compound, and we have the
law of constant proportions. One atom of one substance may com-
bine with one atom of another substance, or a number of atoms of
one Mubstance may combine with one of another to form a molecule,
Init the number must be a simple, rational, whole number; whence
ihn law of multiple proportions.

Hince the atoms have definite weights, and the laws of constant
and nniltijilo proportions are true, the law of combining numbers
follows art a nccensary consequence of the atomic theory.

Tim (ninHiion as to the size or mass of an atom is one which is
III lit <n»nh In m»ino doubt. We know that they are inconceivably small.
*riiU U «1h»wii by the fact that certain substances will continue to
Hlvo olT odoi'H for n long time, which fill a large space, and still not
h»«»M Hpprnoluhly in weight. The odoriferous particles must be
|»H*aiMil ill Mviiry part of the space, and although the substance will
W»mUihu» to III! UiIm Hpuce with such particles for months or longer,
i-tu* ^Mimmil of iiiiU|er which has volatilized is scarcely weighable.
'J'hU »«hovv** lliii ulinoMt \inlimited divisibility of which matter is
WM»f**»^» . »*»»*i i».v ilMlliiltion the atom is indivisible. The same fact
l«4 l«u«u^Ul. oiit li,Y dl«milving certain coloring matters such as the
miihuh \\\\^^ iu wiUor. Vnry mnall amounts of such substances can
Uii|*!Ml' 'Wi •ipiaooii^hUi iM»lor to comparatively enormous volumes of
\V'iU» Ho. ioluiiuui uuUlt^r must be capable of almost unlimited
\iU\.iUilii^ \u Muiiu VM \\\\h may be effected.

i'vi |».»|. ., uii lUw s\ Um\k\ \\\is U^Kt idea of the size of atoms has been

lu.k'i^'Jn I u«i ii.v L<»nl K^^lviu in England. In his own words:

lu» V4,mi. .\ uuuiliiip oi 14 glolni of glass as large as a pea, to be ma^-

nified np to the ssse of tbe «fcnl : -e-Aii ^:ctf:i;:i:»si3 it?a:r n.^r* S?*!
in the same pcc^MctkcL TLr TijgiT^i jcrui^x^ '•":iiji r* i'.»kr?r^-
grained than a heap of s=x^ iz^ju 'nn iriCihUj Icea» ft:«r<ie^crjk:=fri
than a heap of cri«kes balls."*

considering th^zs far txiTijncIj ibe -rxr.ez^ iriz^f .<rz.i;ii>:cj^ vlicl ^kke
place in chemical rcactiocs^ az>d i^v* !•: irT^^i vf::: ^rtiin 4^-CT■^kliI»-
tions which hare l«ii mcLiei. arac vLiti 2i* a: :l-r f.ei*ii:-,c of
the science of chemisXTj. Were we lo *.t»:c- ier* aa-i t«e-^-:i v<ir >r jdv
of the several elements, we w-xild I«arr ^r;M*i<-hed a elass v f j lie^
nomena whose importance ca:in<:< easily t* CTrr*t5T:3iar«*L

Whenerer we hare cbeniieal ncAC-iioii rak^i^g rlace we Lave Leax
liberated or absorbed, and Tis^^allj libenied. TLis l^rls^rs -^s lo a
study of the energy jchanges wbdeh are iiiseparably cc*nnei-:ed with
all chemical action.

Energy manifests itself in a nmtt-er of forms. We have lirfit
energy, heat energy, electrical energy, mei-banical energy, and iLese
are mutually convertible into one aniXher. That mechanical ent*wnb-
can be converted into heat is shown wherever friction exists;. Rub
together two pieces of metal and both become hot. . That heat
energy can be converted into light energy is illustrated by a piece of
metal which has been heated to incandescence. That heat energy
can be converted indirectly into electrical energy is shown by the
dynamo, and so on. This principle of the mutual convertibility of
the various forms of energy is known as the principle of (he cm-
relation of energy, and is an important generalization in physical
science.

That one form of energy can be converted qualitativel»f into
another is important, but far less important than the fact that one
form of energy can be converted quantitatively into another. When,
for example, mechanical energy disappears, as when a haninior falls
upon a metal plate, the heat energy produced is exactly equivalent to
the mechanical energy which has disappeared. If the heat ent»rj;y
produced under these conditions was transformed into work, it would
raise the hammer again exactly to its original position. This prin-
ciple, fundamental to the science of physics, is known as llio pn'ncijtle
of the conservation of energy. It says in words that no onerjxy can
be created or lost, and is analogous to the law of the conservation of
mass, which we have already studied.

Importance of the Conservation of Energy for the Science of
Chemistry. — The bearing of the conservation of energy upon chem-
istry may not appear at first sight. In addition to the forms of

14 PRINCIPLES OF INORGANIC CHEMISTRY

energy enumerated above we should add intrinsic energy, whicli is
frequently referred to as chemical energy. This form of energy
exists in practically all substances in larger or smaller amounts, and
is the form which is converted into heat when a piece of coal is
burned. The existence of this form of energy is essential to all
chenii(;al action, and is, therefore, absolutely essential to the science
of chemistry. It is this form of energy which is converted into
heat whenever a chemical reaction takes place with the evolution of
heat. Indeed, the transformation of intrinsic energy into heat lies
right at the foundation of most chemical reactions and is the chief
cause why such reactions take place. It is sometimes stated that
chemical reactions are accompanied by heat evolution. This state-
ment is misleading, since it lays stress upon the less imix)rtant phe-
nomenon. Indeed, it confuses cause and effect. We should probably
be much nearer the truth if we said that the thermal change was
accompanied by material transformations, which gave rise to new
products with properties for the most part different from those of
the original substances.

Although we cannot discuss this point more fully in the present
connection^ we can see that the energy changes which take place
during chemical reaction are of prime importance.

Although only a part of the intrinsic or chemical energy in the
substances which react is converted into heat or some other form of
energy during the reaction, yet this part which disappears is converted
quantitatively into other forms. The law of the conservation of
energy is, therefore, fundamental to the scientific study of chemistry.

CHAPTER III

OXYOBN (At. Wt. = 16.0)

Ocenrrenoe in Nature. — Oxygen is the most abundant of all' the
chemical elements. It forms about 88.8 per cent of all the water on
the earth, and about 23 per cent of the atmospheric air. It is an
important constituent of most of the rocks, and occurs in nearly all
living matter whether vegetable or animal. It is estimated in general
that about one-half of the earth's crust is composed of the element
oxygen.

Preparation of Oxygen. — Since oxygen occurs in such large quan-
tities in nature, we would think that we should turn to some natural
source for a supply of this element. It is, however, not very easy
to obtain pure oxygen from any natural source. It can be obtained
from the air, but not very readily. It is much more difficult to
separate it from its compounds in the rocks. It can be obtained
from water by decomposing the water by means of an electric cur-
rent, but there are far more economical and convenient means of
preparing oxygen than by the electrolysis of water.

One of the most convenient means of obtaining oxygen in the
laboratory is by heating potassium chlorate. This compound, which
is represented by the formula KClOj, contains about 39 per cent of
oxygen, and gives up all of its oxygen when moderately heated.
The decomposition of the chlorate proceeds in two distinct stages,
which we shall study later in more detail. The final result is as
indicated ; all the oxygen is set free and potassium chloride remains
behind. This is expressed by the following equation : —

2KC10s=2KCl+30j.

Another method of preparing oxygen is by Jieating mercuric oxide.
It is decomposed at once into metallic mercury and oxygen in the
sense of the following equation : —

2Hg0=2Hg4-0^

Hydrogen dioxide, a compound having the composition expressed
by the formula HjOj, when brought in contact with many substances

15

16 PRINCIPLES OF INORGANIC CHEMISTRY

such as the metals, or compounds which are themselves rich in
oxygen, gives up half of its oxygen, becoming water : —

2HA = 2H,0-hOj,

Oxygen can be readily obtained from the compound barium
dU/xide, When ordinary barium oxide, BaO, is heated and a current
of air passed over it, it takes up oxygen from the air, becoming
barium dioxide. When the dioxide is subjected to diminished press-
ure, it gives off oxygen and passes back again into barium oxide.

2Ba02 = 2BaO-hOj^

This is the most convenient means of obtaining oxygen from the
air in pure condition. The oxide of barium takes up oxygen from
the air, forming the dioxide of barium, which can in turn be decom-
posed into oxygen and oxide of barium. The latter can be converted
again into the dioxide and the process continued at will.

Snbstanoes burn readily in Oxygen. — One of the most character-
istic of the chemical properties of oxygen is the readiness with which
substances burn in it. Substances which burn comparatively slowly,
or will not burn at all in the air, often burn with the greatest readiness
in oxygen gas, emitting very bright light and evolving large quantities
of heat.

Fill a number of glass vessels with oxygen gas in the following
manner: First, fill the vessels with water ajid invert them in a
trough containing water. Place some potassium chlorate in a glass
reUjrt, connect a piece of rubber tubing with the neck of the retort,
and then heat the retort gently with a Bunsen burner. After all the
air 4ias been expelled, bring the end of the rubber tube beneath the
mouth of the glass vessel and continue to heat the retort. The
oxygen which is set free by the decomi)08ing potassium chlorate will
rise in the glass vessel and displace the water with which it is filled.

The arrangement of the apparatus for preparing oxygen is shown
in Fig. 1 . The glass retort E containing the potassium chlorate is
heated by the Bunsen burner B, The glass cylinder C is filled with
water and dips beneath the water in the glass trough T. The rubber
tube A is i)laced beneath the mouth of the glass cylinder after all
the air has been expelled from the retort, and the cylinder filled with
oxygen gas. Fill a number of such cylinders with oxygen gas and
the following experiments can be readily carried out

Ignite a pine splinter until it burns to a coal. Extinguish the
flame and plunge the splinter with the coal on the end into a vessel
containing oxygen. The splinter will burst again into flame.

OXYGEN

17

Place a piece of salphur in a deflagrating spoon of convenient
shape and size; ignite the sulphur and plunge it into a vessel filled
with oxygen. The sulphur, which in the air burns with a blue flame
of small luminescence, bursts into violent combustion in the oxygen,
evolving large amounts of heat and light.

A piece of carbon is placed in a similar sixx>n heated to redness,
and plunged into a vessel tilled with oxygen. The carbon burns
vigorously, with evolution of large amounts of heat and light.

Fio. 1.

While a piece of iron will not bum with any appreciable velocity
in the air, it burns very readily indeed in pure oxygen. This can be
shown as follows: Take a steel watch-spring and wrap one end
with cotton thread. Plunge this end into molten sulphur, when a
comparatively large amount of the sulphur will adhere to the thread.
Ignite the sulphur and then plunge the iron into the vessel of oxygen.
The sulphur will first burn vigorously and heat the iron to a very
high temperature. The iron will then burn in the oxygen with an
intense white light, and a large number of highly heated particles
will fly off from the iron, producing quite a pyrotechnic effect In
this experiment it is well to have the vessel containing the oxygen
placed upon a stone slab or immersed in a vessel containing water,
since otherwise the molten iron may fall upon the support to the
vessel and break it, thus interrupting the experiment. This experi-
ment illustrates particularly well the difference between combustion
in the air and in oxygen.

Another experiment which is frequently used to illustrate this
same point is the burning of phosphorus in air and in oxygen. While
c

18 PRINCIPLES OF INORGAKIC CIIEMISTRV

phosphorus bums quietly in the air, in pure oxygen the combustion
takes place with great violence. Introduce a small piece of phos-
phorus into a deflagrating s[>oon, ignite it, and immerse it in a vessel
tilled with oxygeu. The vessel should be large to avoid being broken
by the heat which is liberated in such large quantities. It is also
advisable to take the precautioa to wrap the vessel with a towel, to
avoid pieces of glasss from flying in case the vessel should break*

Explanation of the Above Result*, — The above results show
beyond question that certain substances which burn slowly in the
air, or do not burn at all, burn readily in pure oxygen. This natur-
ally raises the question why this is the case. The air, as we shall
learn, is essentially oxygen diluted with about four times its volume
of nitrogen. The number of oxygen particles in a given volume of
air is, therefore, much less than in a given volume of pure oxygen.
The nitrogen serves to dilute the oxygen. When combustion takes
place iu pure oxygen, the heat which is lit»erated is expended in
raising the temperature of the oxygen alone^ and the rapidity of the
combustion depends chiefly upon the temperature of the oxygen gas*

When the oxygen is diluted with an inert gas like nitrogen, much
of the heat which is set free during the combustion is expended in
raising the temperature of the nitrogen^ which takes no part in the
combustion, and as far as accelerating the combustioa is concerned
is, therefore, lost.

COMBUSTION

CombEstlon. — The subject of combustion, or burning, is one which
has attracted the attention of chemists from very early times. This
would be expected, since combustion is among tlie must familiar of
chemical phenomena. There is evidence that fire was known very-
early in the development of the human race, and its economic im-
portance cannot of course be easily overestimated. When combus-
tion was tirst observed, chemical knowledge, if such it may be called,
was of the very crudest sort. The conception of elements did not
exist, still less the conception of the element oxygen. Tliey observed
that substances apparently disappeared either wholly or in [>art when
burned* and they saw the iire or flame escape from the burning mass.

The FMogiston Theory of Combustion. — The tendency of the
human mind in time past was the same in one respect as it is to-day.
It was not content with simply observing facts ; it wished to account
for them and explain them, hence the origin of theories. From all
the facts which were early observed concerning combustion, esiiecially
the disap|>earance of the substances as they burned and the es(!ai>e

OXYGEN 19

of flame, it seemed evident that in combustion something escaped.
Although they could not discover what this suhstance was they
applied a name to it. It was termed pfdoyiston, and the theory, the
pfdogiston theory of combustion.

According to this theory when a substance burned it gave off
phlogiston, and the products of combustion dififered from the sub-
stance before it was burned in that they had lost phlogiston.

This theory of combustion held sway until oxygen was dis<;overed
by Priestley and Scheele about 1774-1775. The study of oxy^^en and
the part it played in combustion entirely overthrew the phlogiston
theory of combustion.

The Sdle of Oxygen in Combnition. — It was shown by the
Swedish chemist Scheele, that atmospheric air in which a substance
has been burning for a time is no longer able to supiK>rt combustion.
This made it probable that there was something in the air which
had disappeared during combustion. Scheele and also Priestley
showed how a gas could be obtained which supiKjrted combustion
far better than atmospheric air. The former obtained his gas by
heating saltpetre, the latter by heating oxide of mercury.

It remained, however, for the French chemist Lavoisier to show
the real significance of oxygen in all ordinary cases of combustion.
When a substance burned it united with oxygen, and combustion
consists in the union of the substance burned with oxygen. This is
the conception of combustion which we hold at the present da^-, and
is diametrically opposed to the theory of phlogiston. According to
the phlogiston theory of combustion something es(ai)es when a sul>-
stance is burned; according to the present theory an elt^nient, oxygen,
is added to the substance which is undergoing combustion.

Increaie in Weight in Combustion. — If combustion consists in
the union of oxygen with the sul)stance burned, then the weight of
the products of combustion must be greater than the weight of the
substance which has been burned. This alone would seem to be a
crucial experiment to decide between the i)lilogiston theory and the
oxygen addition theory of combustion. It would only be necessary
to weigh the body which is to be burned, and to weigh the products
of combustion, and see which is the heavier.

The phlogistonists, however, would not admit that this was any
test of their theory. Indeed, in the later period of the theory they
knew very well that the products of combustion are heavier than
the substance before it was burned. This fact they easily reconciled
to their theory. They said that phlogiston has negative weight —
weighs less than nothing — and when it escapes from a substance as

20

PRINCIPLES OF INORGAKIC CHEMISTRY

in combustion, the substance becomes heavier* This line of argn-
ment wauld hardly appeal to any one at the present day^ and 13
given simply on account of its historical interest.

That the products of combustion weigh jnore than the substance
before it was burned can be readily shown by the following experi-
ment (Fig. 2) ; Two pieces of candle of equal length are placed, one

upon each pan of a
large balance. A lamp
chimney is suspended
from each end of the
arm of the balance. A
piece of wire ganze
which fits the chimney
tightly is introduced
into each chimney, and
some coarse pieces of
caustic sotia added.
Caustic soda is now
added to the lighter
side until the pointer
stands exactly in the
middle of the scale*

^

¥r.. ±

V

One of the candles is now lighted, and
the products of the combustion, carbon
dioxide and water, are caught by the caus-
tic soda- After the candle has burned for
a time this arm of tlie balance will begin
to sink, showing that the products of com-
bustion of the candle are heavier than the
unburned candle.

Oxygen ^sed up ia Cornhnstian. — That
oxygen is actually used up in combustion
can be readily shown by the following
exj>eriment. Fill a glass tube with air
as shown in Fig. 3. Introduce a piece of
phosphorus. This will undergo slow com-
bustion and the oxygen will be used upi as
13 shown by the fact that the water will
rise steadily in the tube.

Exjieriment 2 shows that the products
of combustion are heavier than the substance before it la burned^ and
experiment 3 that oxygen is used up in combustion. It is oxygen

■^

Fm. X

OXYGEN 21

iprhich adds itself to the burning substance, and combustion is noth-
ing but oxidation.

Bapid and Slow Oxidation. — Combustion, as we ordinarily observe
it, is a comparatively rapid process. The substance bums up, as we
say, in a few minutes, and there is usually a larj^e evolution of lifat,
and in many cases a marked production of light. This is known as
rapid oxidation.

We know oxidation processes, however, which take jJace slowly
and extend over long periods of time, even years. Kxani])It>s are the
oxidation of metals, the decaying or slow oxidation of wood, and
the like.

When the oxidation proceeds slowly, as in these cases, there is no
apparent evolution of heat and no evolution of light. The question
arises, Are we justified in concluding that there is actually no evo-
lution of heat when slow oxidation takes place ? We cannot detect
any heat set free, but it might readily be that there is a slow evolu-
tion of heat, but so slow that it escapes before it can Ix? (h'tr(!ted.

While we cannot prove directly, unless large nia-sses of substances
are employed, that heat is set free in slow oxidation, it can be ])roved
indirectly. The products of slow oxidation are in many cases the
same as the products of rapid oxidation where much heat is evolved.
Since the original substances which combine are the same whether
the oxidation is slow or rax)id, and since the products are the same,
the same energy relations obtain whether the oxidation is rapid or
slow. From the conservation of energy, then, we know that heat is
evolved in slow oxidation as well as in rapid oxidation, and further,
that exactly the same amount of heat is evolved wlien a given quan-
tity of any substance is oxidized to a given oxide, whether the oxida-
tion takes place slowly or proceeds rapidly to the end. This necessary
consequence of the law of the conservation of energy is of more than
ordinary interest

Keasurement of the Heat of Combustion. — The measurement of
the amount of heat which is set free when combustion t^ikes place
is not a simple operation. Indeed, the accurate measurement of the
amount of heat is always more or less difficult, on account of the fact
that heat always flows from the warmer to the colder body, and so
many substances are comparatively good conductors of heat.

To measure the amount of heat set free in any chemical reaction,
such as combustion, the reaction must be carried out in a vessel sur-
rounded by a poor conductor of heat, so that the loss in heat will be
reduced to a minimum. The heat which is produced is allowed to
warm a known weight of water, and the temperature of the water

22 PRINCIPLES OF INORGANIC CHEMISTRY

in uoUhI l)oforo and after the experiment. The apparatus which is
UHnd for measuring quantity of heat is known as a calorimeter. It
conHiMts of an innermost vessel into which a weighed amount of
waUir is introduced. The reaction takes place in this vessel or in a
voHHid which is immersed in the water. The vessel containing the
waU'r iH surrounded by some poor conductor of heat, such as felt or
indtir-ilown. This is then surrounded by one or two layers of air,
which is a i)oor conductor of heat. Even when all of these precau-
tions are taken to prevent loss of heat, the rate at which the calorim-
eter loses heat must be determined, and a corresponding correction
introdu(?ed. If we know the amount of water used in the calorimeter
and the rise in temperature produced, we know the amount of heat
set free as the result of the reaction.

Some unit must be adopted for expressing the results of calori-
metric measurements. Whatever unit we select would be purely
arbitrary. The amount of heat which is required to raise one gram
of water one degree in temperature has been proposed as the unit of
quantity of heat. Since this quantity depends upon the temperature
of the water, and varies quite appreciably with the temperature,
it is necessary to define the temperature. The amount of heat
required to raise one gram of water from 0® to V C. is taken as
the unit, and is called the calorie, and written cod. The calorie is
sometimes defined as one one-hundredth of the amount of heat
required to raise one gram of water from zero to one hundred
degrees. The two definitions are for all practical purposes essen-
tially the same.

Sometimes it is more convenient to use a larger unit, and two
»uoh have been proposed and adopted. One is one hundred times
\\\^ ?^wallor calorie, and is written A'a/. The other is one thousand
Itm^yt th<* snmller calorie, and is written CaL The relations which
v\»** lvtwiH»n the three units is, then, 1 Ccd = 10 Kal = 1000 col,

b» v»t\U»r that the heats of combustion of substances may be com-
ji,%w*Mi>, n\* uuiHt use comparable quantities. We might take any
.^ili,t.»H»v ^jn.^utity of different substances, say ten grams of each.
hMi. t.U»MO »p».uauio8 would not be comparable, since they would not
ur4k^*.ui. iUi> tpi.uUitioH of the different substances which would com-
Uiu,. ^ i\U tmo (4U\»thor. It is best to take quantities of the different
vikti.M***^ hUioU tm* prt>iH>rtional to their combining weights, but

Hv4V A \tv\mk%iim Md of Decomposition.— We have just seen
iU*k\ uii. u uw» «M iiioiti ttuKntances unite and form a third sub-
mU^u. ., i»..a i-i o\^»1>\kI Further, a definite amount of heat is

oxy(;ex 23

set free when a given amount of any substance is formed. This
amount is known as the heat of formation of the sul>stance.

Given a substance already formed by the union (»f two or more
substances. A certain amount of heat must Ik* addtnl to it to de-
compose it into its elements. This is known as th«» heat of d»'rom-
position of the substance.

A very beautiful relation has been established lM»tw«»en thf hrat
of formation of a substance and its heat of decomposition. Th- tuo
are equal This will be s(^en at once to Ik? a n«»(essary conseiiuence
of the law of the conservation of enerjxy. Starting with any siil»-
stanceSy we allow them to combine. If now we (h»eonipose the eom-
pound formed into the original substanres, we come b;uk to exactly
the same condition under which we starttMl, and tlie sann' energy
relations must obtain at the end as at the beginning <»f the pnK-i-ss.
Exactly the same amount of heat which was set free during tli»* for-
mation of the compound must be added to the comijound to decom-
pose it again into its elements.

Vamet of the Compounds formed with Oxygen. — Tlie c<)iii])oun(ls
of oxygen with the other elements are tennecl oxidrs. Whvu sul-
phur was burned in oxygen, the gastH)us compound fornietl is known
as oxide of sulphur. The name actually used, however, is even
more explicit. One atom of sulphur combines with two atoms of
oxygen, giving the comjwund 8(V To indicate the i)resence (jf two
oxygen atoms in the molecule the conii)ound is termed sulphur
dioxide.

The compound formed when carlxui burns in oxygen is known as
oxide of carlx>n. There are, however, two oxides of carbon, one con-
taining one atom of oxygen to one of carbon (CO), and the other, two
atoms of oxygen to one of carbon (CO^). The one formecl in our
earlier experiment, where carbon was burned in pure oxygen, con-
tains two atoms of oxygen to one of carbon and is known as carbon
dioxide. The oxide of carbcm containing one atom of oxygen to one
of carbon is known as carlxni monoxide.

When phosphorus is burned in pure oxygen, tlie resulting com-
pound has the composition re])resented by the formula l^Oj. This
is known as the pentoxide of phosphorus. There is another oxide (»f
phosphorus having the composition P^Oj, and this is known as the
trioxide of phosphorus.

When iron is burned in oxygen the resulting compound has the
composition Fe304 and is known as ferrous ferric oxide, while the
compound FeO would be known as ferrous oxide. The compound
FcgOa is ferric oxide. The terms "ic" and **ous" have come to have

24

rRlNCIPLES OF INORGANIC CHEMISTRY

a generic significance j " ic " is applied to tlie oxide richer in oxygen,
and ** 0U9 " to the oxide which contains less oxygen.

PHYSICAL TROPE RTIES OF OXYGEN

Certain PhyBical Properties of the Element Oxygen. — Oxygen

under ordinary conditions is a tmns parent, colorless, odorless gas.
It is somewhat heavier than air, having a specific gravity in terms
of air as nnity of 1.1056, A litre of oxygen under normal con-
ditions of temperature and pressure, i.e. at 0^ and 760 millimetrea
pressurej weighs 1,4296 grams. In terms of hydrogen as the unit
the specific gmvity of oxygen is 15.88. This we shall leavn is the
ratio between the i-ehttive weights of the atom of hydrogen and the
atom of oxygen. Oxygen is only slightly soluble in water. At 0''
-100 volumes of water dissolve 4 volumes of oxygen. At 15^, 100
volumes of water dissolve 3.4 volumes of oxygen. Oxygen is much
more soluble in alcohol than in water^ 100 volumes of alcohol dis-
solving about 28 volumes of oxygen.

The Pressure of Oxygen varies with the Conditions. — We have
referred to the weight of a litre of oxygen under normal conditions
of temperature and pressure. This would imply that the weight of
a litre of oxygen would change if we changed temperature or press-
ure, and such is the fact. If we have a litre of oxygen at any given
pressure and subject the gas to a greater pressure, the volume would
be less than a litre, and, consequently, the density of tlie gas would
be increased and the weight of a given volume of the gas increased.
Similarly, diminution in pressure would cause increase in volume,
and, consequently, diminution in the weight of a given volume of
the gas*

If instead of varying the pressure we vary the temperature to which
the oxygen gaa is subjected, we would also produce change in volume.
If the temi>erature of the gas is increased and the pressure kept con-
stant, the volume of the gas would increase. If, on the other hand,
the temperature of the gas is lowered, the pressure being kept con-
stant, the volume of the gas would be diminished. Certain qnantit>a-
tive relations between the pressure and volume, and the temperature
and volume of not only oxygen gas, but of gases in general, have been
established, and these will now be briefly considered.

The Law of Boyle for Oases. — As already statedj the volume of a
gas becomes smaller with increase in pressure, and with increase in
pressure the density of a gas l>ecomes greater. The relation con-
necting these properties is very simple, The pressure of a gas is

OXYGEX 25

proportional to its density, and both are inversely proportional to

the volume. If we represent the pressure by p and the density by

d, we have — .

/> = crf,

where c is a constant for a given temi^erature. If r is the volume
and m the mass of the gas, Boyle's law may be expressed thus : —

2)v = cuu

If p is the pressure and r the volume of a pven mass of gas
under one set of conditions, and ;), and 1*1 the pressure and volume
of the same mass of gas under other conditions, Boyle's law may be
expressed thus:— ;>r=;vv

The product of the pressure and volume of a given mass of gas at
constant temperature is a constant.

While there ai-e many exceptions known to the law of Boyle,
especially when the gas is under either very slight or very great
pressure, it holds approximately in the great majority of cases, and
is one of the two fundamental laws of gas-pressure.

The Law of Oay-Lossac for Oases. — If a gas is kept under con-
stant pressure and its temperature raised, the volume will increase.
If the volume is kept constant as the temperature rises, the pressure
will increase. The remarkable fact has been discovered that the
increase in the volume of the gas for a given rise in temperature is
a constant, independent of the nature of the gas. All gases increase
about yfy (= 0.003C65) of their volume at 0°C., for every rise of one
degree in temperature. Gay-Lussac's law states that this teraperar
ture coefficient is constant for all gases.

If we keep the volume constant and warm the gas to f*, the press-
ure P at this temperature is calculated from the pressure po at 0^ as
follows : — P = p^, (1 + 0.003665 0-

If, on the other hand, the pressure is kept constant and the volume
allowed to increase with rise in temperature, the volume at ^, V, is
calculated from the volume at 0^, r©, as follows : —

r=t'b (1 + 0.003665 0.
If both pressure and volume are allowed to change when the gas is
heated, the pressure and volume at ^, ;; and v, are calculated from
the pressure and volume at 0° as follows : —

pv = poVo (1 + 0.003665 1),
from which, vo =^^ (i ^ aiio3665 O"

20 PRINCIPLES OF INORGANIC CHEMISTRY

This is the expression generally employed for reducing a gas to
what are termed normal conditions. If the volume v of the gas
is read at a given i)re8sure p and temperature t, we can calculate at
once the vohime at 0®, Vq, and normal pressure p^ which is taken
AM 7(\{) millimotreH of mercury. Exceptions are known to the law
of (lay.IiUMMuo as to the law of Boyle, but we have here a law which
appliim to gamm in general.

The Detarmixiation of the Absolute Zero of Temperature. — The
value of the constant 0.003665 is determined either by keeping the
pnmMure (jonstant and measuring the increase in volume with rise
ill teiii|ieratiiro, or by keeping the volume constant and measuring
thtt liMJniaMo in pressure as the temperature rises. The values found
by the two methods differ only slightly, and we take 0.003665 as
very nearly the true value of the temperature coefficient of a gas.

Thin \h very nearly ^fy, which means that if a gas is cooled
down to — 273** C, its volume would become zero if the law of Gay-
liUMN(U) heUl down to the limit. This temperature, termed the
uhMolula x**roy has now been nearly realized experimentally. It is
(|uite eertaiu that temperatures have been produced which are
within twenty degrees of this temperature, as we shall see. It is,
however, very probable that the laws of gas-pressure do not hold
at thene extremely low temperatures.

The Combined Expression of the Laws of Boyle and Oay-Lussao. —
TheHo two fundamental laws of gas-pressure can be combined in
one ex press ion.

If we reprewmt temperature as measured from the absolute zero
by y, the combined expression of the laws of Boyle and Gay-
LuHsae is: —

7^',^* is usually represented by R, when the above expression becomes,

pv=ET.

The Liquefaction of Oxygen. — Although oxygen is a gas under
atmosphttrici prcHsure and at all ordinary temperatures, it does not
follow that it is a gas at all temperatures and pressures. If we
look into the history of the liquefaction of gases, we find, however,
that oxygen resisted for a long time all efforts to liquefy it, and
was placed among the so-called permanent gases.

The oarly work on the liquefaction of gases made it obvious
that two conditions were necessary in order that a gas may be
liquefied. It must be subjected to a high pressure and to a low

OXYGEN 27

temperature. By fulfilling these conditions the English physicist
Faraday was able to liquefy many of the more common gases.
There were several, however, which resisted all efforts to liquefy
them, and among these was oxygen. Natterer subjected oxygen
to a pressure of between 3000 and 4000 atmospheres, at the same
time cooling it far below the ordinary temperatures, but was not
able to obtain it in the liquid form.

It is quite certain that oxygen would have been known only in
the gaseous state for a much longer period of time, had not the
discovery been made which we owe to Andrews. He pointed out
that there is a temperature above which a gas cannot be liquefied
no matter how great the pressure to which it is subjected. This
temperature he called the critical temperature, and for oxygen this
is now known to be — 119**.

This explains why Natterer was unable to liquefy oxygen when
he subjected it to a pressure of more than 3000 atmospheres. The
gas was not sufficiently cooled. It was above its critical tempera-
ture.

When oxygen was cooled down to its critical temperature, the
pressure required to liquefy it was only 50 atmospheres, slu^ this is
known as the critical pressure of oxygen.

Oxygen was first liquefied in 1877, simultaneously by two experi-
menters, Pictet and Cailletet. The method of Pictet is based upon
the fact that when a low-boiling liquid evaporates, especially when
under diminished pressure, a temperature much lower than its own
boiling-point is produced. By surrounding oxygen with liquid
carbon dioxide which boils at — 78®, and allowing the liquid to
evaporate under low pressure, a temperature is produced (—140®)
which is below the critical temperature of oxygen. At this tem-
perature the oxygen liquefies at a pressure below its critical
pressure.

The method of Cailletet is based upon a different principle.
When a gas is strongly compressed and then suddenly allowed to
expand, it cools itself enormously. Cailletet subjected oxygen to a
pressure of about 300 atmospheres and then allowed it to expand
suddenly. Drops of liquid oxygen were obtained.

A method of obtaining liquid oxygen in quantity, which is greatly to
be preferred to either of the above, is that of Linde, who has done so
much toward the liquefaction of the more resistant gases. The method
is based upon the cooling of a strongly compressed gas on expanding.
Air is compressed, then allowed to cool itself by expanding. This
is made to cool another quantity of compressed air, which in turn

28

PUlNCirLE5> OK IXORGAXIC CUEMISTRY

IS :illi>wtHl U» i'xi>iuul ami ostablish a still lower temperature. This
4't»liK'i- iiir is Ihi'ii allowed lo OiK>l slill auotlier jK.»rtioii of compressed
jiir, iiiid so iMi until a ti'inporatuiv is reachod at which air liqueties.

W'r lia\t^ MTiu lio\vc\cr, that air is a mixture chietiy of oxygen
and nilmvi'M. It now ivmaius to separate the liquid oxyvreu from
llii^ li«|ind iiitro^ru. N ilri>j:i*u, as wo shall learn. l>oils lower than
o\V).'.«*n. W 1 1 I'll a nu\tnro of liquid oxyiren and li«[uid nitrogren is
««\|m»jimI (i> tndiiiary li'nnH»raturcs, the uitror^'u, Winj^ the lower boil-
ing, liipiid, will luiil olT tirst« and finally leaves Whiud comparatively
|iiirn ltt|iii(l (i\\ v;;t'ii.

Till 1 1 1 .1 iiirtlmd of olilainiuiT o\y i^»n fivm the air in comparative
pill lis iiinl Ml iMittrniuns ipiantitics* whcivvcr a liipiid air plant is
av.iil.iiilo It rilionld Im) added to the metlunls discussed at the be-
(jihiiiiii; «>l )lii:i rliaptiM' lor ubtainini; oxy^'u.

I*io|i0itl0ii of Liquid Oxygen. - We owe our knowledge of the
pi«ipt«ilhM (il liquid uwm'ii almost entirely to the liussians, Wro-
liU'w-iKt iiiid Ulri/(i\srtlvi, iiiid to t lic Kii^lisliman. l>ewar.

NN iiibU'w ilii and Ul-i/t^w.'dvi have contributeil much to our knowl-
i<(b;it III llio rliMiintl o\m;*mi wIumi in the liquid condition. Dewar,
hii\tiiv. id hit dt qui lid Im liqnrfyin>; ^ases the enonnous plant of
thn l!o\al institution of (ireat l^ritain, has
ol it a inrdo\v^cn and ot her lowdK>iling liquids,
111 wiihliall s(M\ in quantities never apiu'oached
l»\ iin\ mie rise. Oewar has devised a form
tii lip pit I a I IIS I'tir ] unserving liquid oxygen
and ollii-r luw boiling liquids, which deserves
hpt«('iiil nut ici*.

Il hi Will Known that a vacuum is a very
poor roiidmlnr of heat. The rate at which
i( liquid will i'va|K)rate depends primarily
upon lh«» iiiln at which it can secure heat,
whieh In iiliMoliili'ly m-crssary in onler that
(ho liquid ntii\ pass ovi*r into vapor. Dewar
I'oUhhueled doiiblf wallrd, glass vessels and
pmupiMl iMil Ihn ail' bi'tween the walls. The
iiiKMmomeiil \ti nhtiw'n ill Fig. 4. The air is
ptimpod \\\\\ I It nil Ihii Hp.'in* lH*tween the two

' ^t"n Ihu i'ii|||irit(|ii|| \M(ll llii« pilinp himIimI otT.

I ' > ' / <v,i h piiti 1 1| hi nunh il *' \ui«uiiiti jiiflict.cfl ** apparatus will

■ \:>M:kii,i ly fliml\, luul iMMi bi* pi'CNiTved for quite a

" '>( ttt'i>4ii»luii hiivo Ki'iMilly I'acilil.ated the study of

< \

OXYGEX 29

Liquid oxygen is light blue in color, boils at — 181®, and at its
boiling-point has a specific gravity of 1.135, water being taken as the
unit. The specific gravity varies so greatly with the temperature
that at the critical temperature of oxygen, — 119®, it is only 0.65.

It is obvious that liquid oxygen furnishes us with an excellent
means of obtaining very low temperatures. While under atmos-
pheric pressure it boils at — 181®, when allowed to boil under a
pressure of a few millimetres of mercury a temperature as low as
— 225® can be obtained.

Power of Oxygen to enter into Chemical Combination. — The
element oxygen has rather remarkable power of entering into com-
bination with other elements. It combines with all of the more
common elements with the exception of fluorine. Of the rarer
elements it forms compounds with all except those recently dis-
covered by Ramsay. These elements, argon, helium, neon, krypton,
and xenon, do not combine with oxygen, but it should be said that
thus far they have not been made to combine with any other sub-
stance or with one another. No chemical element combines more
generally with other elements than the element oxygen.

OZONE

Allotropic Kodifioation of Oxygen. — We have dealt thus far with
the element oxygen in the condition in which it is ordinarily known
to us. Oxygen can, however, occur with very different properties
from ordinary oxygen. The second modification of oxygen is known
as ozone. The property of an element to occur in two different
modifications is known as allotropy, and ozone is spoken of as an
allotropic modification of oxygen.

Preparation of Ozone. — Every one has noticed the peculiar odor
about an electrical machine which has been in operation for a time.
The same odor was observed by the Dutchman, Van Marum, as early
as 1785, when an electric spark was passed through oxygen. This
gave the key to the preparation of the substance, which was dis-
covered in 1840 by Schonbein. When an electric spark is passed
through oxygen the volume of the gas diminishes and the result is
a mixture of oxygen and ozone.

Ozone is formed in larger or smaller quantities under a number
of conditions. When phosphorus is exposed to the air it underg6es
slow oxidation, and at the same time some of the oxygen of the air
is converted into ozone.

Ozone is also formed in small quantities in certain reactions where

30

raiNCIPLES OF INORGANIC CHEMISTRY

oxygen is Hbemted, The oxygen set free when sulphuric aeid acta
on manganese dioxide contains a detectable amount of ozone.

When water atndLdated with sulphuric acnd is electrolyzed, the
oxygen liberated at the anode contains an appreciable amount of
ozone.

When barium dioxide is treated with sulphuric acidj the oxygen
set free contains a very considerable aiwount of ozone.

The best metiiotl^ however, of obtaining 02one in quantity is by
passing electricity through oxygen.

A convenient form of apparatus for preparing ozone is the follow-
ing (Fig. 5) : Into the glass tube GO an iron tube // is inserted. The
glass tube is surrounded for a part of its length by tin-foil. Oxygen

k e

Fj =, o.

Is intTniliiced into the glass tube through the tube A and escapes
thr*Migli ti* A current of water is passed through the tube CQ
to kcn|* flir apparatus cool. The tin-foil, ou the one hand, and Ibe
tube (\ ou tlic other, are connected with the poles of an induction
nuichiue. Silent discharges take place l>etween the tin-foil and the
Irnrj, panning through the oxygen. Under these conditions a part of
tin* fixyj^i'ri i.H rTin verted into ozone.

Propertiai of Oione. — The property by which ozone is most easily
ri*iMigid/.*^i! li* lt« irritating odor, whence the name. Ozone, like oxy-
Ijmi, iN n \i,m urjtb'r urdinary corjditions, but can be converted into a
diirk \\\\w liipiiil It cat! lie detected moet easily chemically by its
m^ibin iipim *i ci>U>rlcj4ii mdntion of potasBium iodide, A dark brown
iiidor a|i|wuM'ii \n Hiich a »ohition wiien ozone is passed through it,
*X\\i^ Wi« m\v\\\\ ii^arn irt due to the oxidizing action of the ozone,
lllmrnUiiil tint Inn, Thin method of detecting ozone was regarded for
a lt*nw tiMH* ui fiundnliiriK evidence that it exists in the atmosphere,
(^ Imv, lM»vv*^viir» \ms\\ fnnnd tliat there are other substances which
IHdMV li iuhitioii of piiiUMMium iodide as well as ozone, and we are
%\{\\ Ih d^uibt iii to whr^tlitM^ oj^imt^ exists in the atmosphere. If it is
|H0*iiut ut ull lu thi» nlnuMplu^rc^ it is quite certain that it exists only

OXYGEN 31

Ozone is, in general, a much more active substance chemically
than oxygen. It has, therefoi*e, come to be known as "active"
oxygen. It has much greater oxidizing power than oxygen, espe-
cially at ordinary temperatures. It will effect oxidations which, at
the same temperature, oxygen is entirely incapable of producing.
Thus, ozone will oxidize a piece of metallic silver at ordinary tempera-
tures, covering it with a layer of brown oxide, while under similar
conditions oxygen is not able to effect such an oxidation.

Transformation of Ozone into Oxygen. — We have seen that
oxygen is transformed into ozone under the influence of the silent
electrical discharge. The question naturally arises, Can ozone once
formed be transformed again into oxygen ? The answer is it can.
When ozone is heated to 300**, it passes back into ordinary oxygen.
We can thus pass either from oxygen to ozone or from ozone to
oxygen.

This raises the important question. What is the cause of the differ-
ence in properties between the two modifications of oxygen as it is
usually stated ? Since either modification can be transformed into
the other, it is obvious that there is some close connection between
them.

The Difference between Oxygen and Ozone. — It is obvious from
what has been stated that there is a marked difference between the
properties of oxygen and ozone, yet ozone materially considered is
oxygen and nothing but oxygen. How can we account for the differ-
ence in the properties of these two substances ?

In dealing with the external universe we must not confine our
attention to what we are pleased to call matter, which is pure theory
and cannot be perceived as such by our senses, but must take into
account especially the various manifestations of energy; since all
that we can learn through our senses are changes in energy or energy
differences. In thinking of element or compound we are liable to
lay too much stress upon the material side, because we fancy that it
is this side which appeals to our senses, and to overlook or deal
lightly with the chemical or intrinsic energy which is stored up in
the substance. Every chemical compound has a great-er or less
amount of intrinsic energy stored up within it, and its chemical
properties are largely conditioned by this intrinsic energy. With
this conception clearly in mind we may approach the problem of the
difference between oxygen and ozone.

The Same Kind of Matter but Different Amounts of Energy. —We
have already seen that oxygen and ozone are made up of the same
kind of matter, since each is transformable into the other. If we

PRINCIPLES OF INORGANIC CHEMISTRY

study this side of the problem quantitativelyi we shall find that when
three voluraes of oxygen are converted into ozone, the resulting gas
occupies only two volumes* Thus, if three litres of oxygen were
convei-ted into ozone, only two litres of ozone would be formed. On
the other hand, if two litres of ozone were decomposed by heat, three
litres of oxygen would be formed.

To anticipate what we shall understand more clearly later, the
atom of oxygen cannot exist by itself in the free state, biit two atoms
of oxygen always unite and form what is called the molecule of oxy-
gen. In oxygen gas, as we ordinarily know it, we do not have the
atoms of oxygen uncombined with one another, but the molecules
which are formed by the union of two atoms.

It has been shown that in the molecule of ozone there are three
atoms of oxygen, while in the molecule of oxygen there are only two.
It is obvious, however, that the difference in the number of atoms in
the molecule, alone considered, is not sufficient to account for such
differences in properties as exist between oxygen and ozone. Indeed,
it is difficult to see how this would produce a difference in any prop-
erty other than the mass of the molecule,

27je real dijffenmce in the properties of oxygen and ozone is due to
the different amounts of intrhufic energy pj'^sent in their molecutes.
This statement is not raa^le dogmatically, but can be demonstrated
exi>enmentally in the following manner: —

When carbou is burned in oxygen the product is carbon dioxide.
When carbon is burned in ozone the product is carbon dioxide. We
start in both cases, with the same substance, carbon, and we end
with the same protlvict^ carbon dioxide. Any differences in the two
reactions must be due to the differences between the oxygen and the
ozone.

If we measure tlie amoimts of heat liberated in the two reactions,
we find that they are very different indeed. Considerably more
heat is evolved when carbon is burned in ozone than when carbon is
bnrned in oxygen, Tliis shows that there is wore intrinsic energy
pre^'ient in the molecule of ozone than in the moleade of oxygen.

This result is just what we would expect from the chemical
behavior of the two mmlifications of oxygen. Ozone is the more
active chemicidly, and ozone contains the larger amount of intrinsic
energy. This alone serves to show the importance of energy rela-
tions in dealing with chemical phenomena*

CHAPTER IV

HTDROOEN (At. Wt = 1.008)

Occurrence. — Hydrogen, which was discovered by Cavendish in
1766, is apparently the most widely distributed of all the elements.
It occurs in the earth's atmosphere in very small quantities. It oc-
curs in the sun, especially in the prominences seen during solar
eclipses, in the stars, and even in the nebulous masses scattered
throughout the universe. It has been found in certain great salt
deposits, as those of Salzburg, Germany, in meteoric iron, and in
connection with natural petroleums.

The greatest amount of hydrogen on the earth, by far, is in water,
whence the name {hydor, water, and gennao, to produce). All water con-
tains 11.19 per cent of hydrogen, and when we consider the amount
of water upon the earth, we get some idea of the amount of hydrogen
present on our globe. It also occurs in most forms of living matter.

Preparation of the Element Hydrogen. — To obtain the element
hydrogen, we would naturally turn to water as the largest source.
Hydrogen can be obtained from water by several means. When a
little acid is added to water and the electric current passed through
the acidified water, hydrogen gas is liberated at one of the i)oles, and
can be easily collected.

Hydrogen can also be obtained from water by purely chemical
means. When water-vapor is passed over highly heated iron, the
iron combines with the oxygen in the water-vapor, and hydrogen is
set free. The equation expressing this reaction is —

3Fe-f-4H20 = Fe304 + 4H2.

There are certain elements which combine with the oxygen of water
even at ordinary temperatures, liberating the hydrogen. Such an
element is metallic sodium. When metallic sodium is brought in
contact with water at ordinary temperatures, a violent reaction takes
place, in the sense of the following equation : —

2 Na + 2 H,0 = 2 NaOH + Ha.
D 33

3i

PRINCIPLES OF INORGANIC CHEMISTIir

Wlien potassium is used instead of sodium, a still more yiolent reac-
tion takes place : —

2K-h2H,0 = 2KOH + H^

111 practice we sel^lom use any of the above methods, since we have
means of preparing hydrogen on a large scale which are far more
convenient than any of these. When zino is treated with a strong
a<^id, such as hydrochloric or sulphuric j the metal passes into
solution and the hydrogen from the acid escapes. In the ease of
hydrochloric acid and zinc, this is represented by the following
equation ; —

In the case of zinc and sulphuric acid by the following equation: —

Zn + H,S04 = ZnS0, + H^

Hydrogen is readily prepared as follows : Introdnce some pieces of

zinc into a glass flask, -^i, as shown in the figure (Fig, 6), and pour dilute

hydrochloric acid into the flask through the funnel-tube Bj until the

g end of the tn\m dips

beneath the acid. Hy-
drogen gas will be liber-
ated and escape through
the side tube C,

The gas can then be
passed through a wash-
bottle filled with water
to remove any trace of
acidj and afterwards
dried by passing through
tubes containing calcium
chloride, sulphuric acid,
or phosphorus i>entoxide.
If it is desired to
prepare hydrogen on a
still larger scale, a form
of apparatus devised by
Kipp is very convenient.
From this apparatus hydrogen is obtained by simply turning a stop-
cock. When no more gas is desired the stop-cock is closed, and
the pressure of the hydrogen generated, automatically drives the
acid away from the zinc and causes the further liberation of gas to
cease.

Fig. 0*

HYDROGEN 85

Combination of Hydrogen with Oxygen.— Hydrogen, a colorless
and odorless gas, combines readily with oxygen at elevated tempera-
tures. A mixture of hydrogen and oxygen can be kept for an indefi-
nite time, provided the mixture is not heated. If the temperature
is raised sufficiently, the two combine with the greatest ease, pro-
ducing a violent explosion.

That hydrogen can be burned in the presence of oxygen without
any explosion taking place can be shown by the following exi>eri-
ment : Attach a rubber tube to the end of the small glass tube 0
(Fig. 6), and insert into the other end of the rubber tube a metallic
tube with a very fine opening. The small tube at the end of a mouth-
blowpipe works very well. Allow the hydrogen to escape f ronj the
apparatus through the metallic tube until every trace of air has been
removed from the apparatus. Then ignite the hydrogen at the end
of the metal tube. It will bum with a flame which is nearly color-
less, but which is intensely hot, as can be shown by inserting a piece
of metal into the flame.

The reaction which takes place between the hydrogen and the
oxygen of the air is represented by the following equation : —

2H2-f Oj = 2H20;

the product formed is ordinary water.

That water is formed in this process can be readily demonstrated
as follows : Bring a cold, dry, glass cylinder over the flame of burn-
ing hydrogen, and hold it in position for a few moments. The inner
wall of the cylinder will quickly become covered with moisture,
and after a short time drops of water will form on the walls of the
cylinder and drop from the mouth.

The explosive nature of the mixture of hydrogen and oxygen can
be readily demonstrated by the following experiment; Mix two
volumes of hydrogen gas with one volume of oxygen gas, and con-
duct some of the mixture through a solution of soap until a mass of
soap bubbles has been formed. The solution of soap should be
placed in a thick-walled, porcelain, evaporating dish. Place the dish
containing the soap bubbles in a protected place, such as under the
hood, and bring the flame of a gas-lighter carefully up to the bubbles
filled with the mixture of hydrogen and oxygen. An explosion will
take place whose violence depends on the size and number of bubbles
present. It is well, therefore, not to have any great amount of the
mixed gases present when the flame is applied.

This mixture of the two gases containing two volumes of hydro-
gen to one of oxygen is known as electrolytic gas or detonating gas,

86

PUINCIPLES OF INORGANIC CHEMISTRY

iinee it is the aame mixture which is obtained when an electric cur-
rent IB passed through acidulated water and the gases liberated at
the two poles allowed to tiiix.

Mixture of Hydrogen and Oxygen affected by the Presence of
Certam Substances. — We have seen that hydrogen and oxygen will
remain in tlie presence of each other nnconiUned for an unlimited
time, provided the temperature to which the mixture is subjected is
not too high. Such a mixture is very materially affected by the
presence of certain substances. If a piece of ordinary platinum foil
is introduced into a mixture of hydrogen and oxygen, the volume of
the mixed gases rapidly diminishes, showing that combination has
taken place» and water is formed* Platinum sponge acta still more
effectively than platinum foilj probably on account of the much
larger surface which it ex|>oses.

One peculiarity of the above reaction is that the j^latinum does
not undergo any change, itself not entering into the i-eaction ; and
further, that a very small amount of platinum may cause an enor-
mous quantity of hydrogen and uxygen to combine.

Other metals produce the same effect, although to a less extent
than platinum^ and often require a higher temperature to cause any
appreciable amount of combination between the two gases.

A special name has been applied to reactions brought about by
the simple contact with some foreign substance. They have been
termed ^ntnbjlk.

Catalytic Eeactions and Catalyzers. — The above is far from being
an isolated exami^le iu the lield of chemistry* On the contrarVj it
is a type of a large nvmaber of reactions. Such reactions^ however,
have certain features in common wliich enable us to classify them.
In all catalytic reactions the substance which effects the reaction —
the catalyzer — does not enter into the reaction. Secondly, a very
small amount of the catalyzer can effect relatively an enormous
amount of chemical combination. As the subject develops we shall
encounter a number of catalytic reactions, and the whole subject of
catalysis and catalyzers has come very much to the front in the last
few years. The opinion is rapidly growing that catalysis plays a
very important part in connection with the life processes, and under-
lies many of the chemical transformations which are taking place in
the living body.

Relations by Volume In wliich Hydrogen and Oxygen Combine* —
It was discovered early in the nineteenth century that hydrogen and
oxygen combine in simple volume relations. Xo mutter in what
proportions the gases hydrogen and oxygen are mixed» for every

HYDROGEN 37

Tolume of oxygen which disappears when combination takes place
two volumes of hydrogen disappear. The ratio of the volumes
which combine is, therefore, one to two.

The further question which remains is what relation exists be-
tween the volumes of the gases which combine and the volume of
the water-vapor formed. The simplest relation would Xye that the
volume of the water-vapor would be equal to the sum of the volumes
of the oxygen and hydrogen which have entered into combination.
Such a relation, however, does not exist. The volume of the water-
vapor formed is less than the sum of the volumes of the gases which
have combined. This is the same as to say that when hydrogen and
oxygen combine there is a contraction in volume.

The relation which actually exists is, however, comparatively sim-
ple. Two volumes of hydrogen gas combine with one volume of oxygen
gas and form two volumes of water-vapor. Three volumes of the con-
stituent gases have disappeared, and two volumes of the product have
been formed. There has been a contraction in volume of one-third.

We shall learn from a study of other cases that this is a general
relation. In the first place, gases combine in simple volume rela-
tions, and in the second, there is a simple relation between the
volumes of the gases which enter into combination and the volume
of the product formed.

Heat Energy produced when Oxygen and Hydrogen Combine. —
That there is a large amount of heat energy proiluced when oxygen
combines with hydrogen is shown by the fact that the vessel which
contains the gases becomes appreciably heated. The amount of
heat which is produced in this reaction has been carefully measured.
When 2 grams of hydrogen combine with 15.88 grams of oxygen,
the heat set free is 68.360 calories.

This is an unusually large quantity of heat to be ])roduced by
such small quantities of substances entering into chemical reaction.
It has been utilized as a source of very high temperature in a form
of lamp which we shall now describe.

The Oxyhydrogen Blowpipe. — The oxyhydrogen blowpipe is a
form of apparatus in which hydrogen is so burned in oxygen as to
concentrate the heat in a small space. The apparatus is represented
in Fig. 7. The hydrogen enters through the side-tube //, and is lit
at E. Oxygen enters through the tul)e 0 and does not mix with
the hydrogen until the flame is reached. If it mixed with the
hydrogen before reaching the flame, we would have electrolytic gas,
or detonating gas as it is sometimes called, and it would explode
violently when a flame was applied to it.

88

PRINCIPLES OF INOEGAKIC CUEAIISTRY

The flame of the oxyhydrogen blowpipe gives very little lights
but is intensely hot It will give some idea of the temperature of
the flame to state that platinum can be easily melted iji it.

While the flame of the oxy hydrogen blowpipe ia itself only
slightly luminotiSj an intense light can be produced by allowing
it to fall upon certain substances which can be heated to a high
temperature without fusion. Such a substance is ordinary lime.
Whea the oxyhydrogen flame is allowed to fall upon a cylinder of

^-^

^^,

Fio. 7.

lime» an intense wlute light \% produced. This m the Drummond
light The light is so intense that it can be used where high illu-
mination is required, as in projecting lanterns and the like.

Dry Hydrogen will act combine with Dry Oxygea. — It would
be gathered froju what has been said thus fur that hydrogen and
oxygen always combine if the temperatui'e to which they are sub-
je^vted is sufficiently high. This is not the case. If very great
precautions are takeji to remove every ti-ace of moisture from both
the oxygen and the hydrogen, the mixture of the two gases may
be hf*rited above TOO*— far above their ignition temperature — with-
out the slightest combination taking place. The significance of this
fact cannot be seen at present, but w* ill become obvious as the sub-
jtiet develops. It lies at the fouuflation of what we believe to be
the true explanation of the cause of chemical action.

The Reducing Power of Hydrogen — ^The tendency of hydrogen
to combine witli oxygen manifests itself, not only when the oxygen is
in the free state, but even when it is combined with other elements,
Hydrogen has the pow^cr of removing oxygen from its compounds
with other elements, especially at somewhat elevated temi>eratures.
The removal of oxygen from a compound is known as redttclhn^ and
tlie substance which can remove the oxygen as a reducing agent

HYDROGEN 89

Take the oxide of zinc, which has the composition ZnO. When
hydrogen is passed over this substance at an elevated temperature,
it combines with the oxygen and leaves the zinc reduced to the
elementary condition.

ZnO-f-H3=HjO-f-Zn.

Similarly, when oxide of iron is heated in the presence of hydrogen
gas, the oxygen combines with the hydrogen, forming water, and
leaves the iron in the free condition.

Yefi, -h 4 Hj = 4 H2O -f- 3 Fe.

This reaction may occasion some surprise when it is recalled
that one of the methods described for making hydrogen was to pass
water-vapor over highly heated iron. The iron combined with
the oxygen of the water and set hydrogen free. Now we have
exactly the reverse taking place, hydrogen combining with the
oxygen of iron oxide setting iron free. Reactions of this kind,
which can proceed either way, — either from left to right, as we
write our chemical equations, or from right to left, — are known as
reversible. The way in which the reaction will proceed is condi-
tioned solely by the relative quantities of the substances present.
If there is a large amount of water-vapor present, the reaction will
proceed thus : —

3 Fe -f- 4 H2O =Fe304 -h 4 Hg.

If, on the contrary, there is a large amount of hydrogen present,
thus: —

Fe304 -f- 4 H2 = 3 Fe -f- 4 H2O.

This it the first time that we have encountered the effect of quantity
or mass on chemical activity.

We shall learn that reversible reactions are the rule and not the
exception in chemistry, and that the effect of mass or mass action
has been formulated algebraically, and is one of the fundamental
generalizations upon which the science of chemistry rests.

Compounds of Hydrogen with Other Metals. — Hydrogen forms
compounds with a number of other elements, and some of these are
among the most important compounds known to the chemist. Thus,
hydrogen combines readily with sulphur and analogous elements,
forming with sulphur the compound HgS, with selenium HjSe, and
with tellurium HaTe. It combines with chlorine and allied elements,
forming one of the most important classes of acids ; the best known
member of which is hydrochloric acid. It combines with nitrogen,

40

ntlNCtPLES OF INORGANIC CnE.>iISTRY

furiiiiu^ thi4 wuU'knowii base ammonia (^Hj), Hydrogen also com-
Hnm ili reedy with a uu tuber of tUe inetals ami furma definite
iHHiipniirnlN with thege substances. These are known as ht/drtdes.
Tij*» (*oni|KUimi witli palladium is eapeeially well known, haviog the
foniiitmitiun l*djH. Hydrogen also eombinea with sodium and
(H!turtftinmi fnnuing XaH and KH, and with calcium, stroutiuinj and
barium^ forming CaH^, SrH^, and BaH^ These nomponnds will, liow-
evett he diaoussed in detail under the several elements in question.

Hydrogen praeent in All Acidi. — We shall learn tliat the element
bydrog<*n is present in every niember of that enormously large class
of compounds known as acids. And, fui-ther, that it is the hydrogen
which gives to these com|>ounds their acid properties. This fact has
come to l>e recognized recently in its full significance through the
investigations of physical chemistry. It was thought for a long
time that oxygen is the element fundamental to acidity. Indeed, the
name means acid former. ComiM>unds w^ere, however, discovered
which are the very strongest acids and which contain no oxygen
whatever, Tlie attempt which was made to fit these cases in with
the oxygen theory of acids will be considered when hydrochloric
acid is taken up.

This role of hydrogen, where it gives acidity to all compounds
possessing it, is by far the most important which it plays in the
whole f^eld <^f t^heuii^stry.

Haseent Hydrogen. — When hydrogen is first liberated by the
action of a metal on an acidj it has very different properties from
those which it possesses after it has once been formetL While
hydrogen gas as we ordinarily know it must be heated to an elevated
temperature before it will reduce the oxides of most metal a, hydro-
gen which is just being formed will reduce many snch subatances
even at ordinary temiHBratures. Maiiy other reactions wdiich hydro-
gen gas will either not effei^t at all, or effect only at elevated
temperatures, will be produced readily at ordinary temperatures by
hydrogen which is just being formed,

Hydrtigen which is just being formed has acquired a specific
name to distinguish it from hydrogen which has been formed for an
appreciable time. It is knowTi as nascent hfdfttgen. This condition
of the nascent state we shall leani is not peculiar to hydrogenj but is
possessed by other elements as well.

The explanation which has been offered to aj(?count for the prop-
erties of substances in the nascent state is based upon the atomic
theory, The hydrogen molecule, like the molecule of oxygen, has
been shown by methods which we shall study later, to consist of

HYDROGEN 41

Vo atoms. When hydrogen is first set free, it is very probable that
it is in what we at present must call the atomic condition — one
atom by itself. This is supposed to be the condition in the nascent
state. After the hydrogen atoms have time to come in contact with
one another, the atoms combine in groups of two, and we have molec-
ular hydrogen as we ordinarily know it.

PHYSICAL PROPERTIES OF HYDROGEN

Certain Physical Properties of the Element Hydrogen. — Hydro-
gen is a transparent, colorless, gas, without taste or odor. It is the
lightest of all known substances, being nearly sixteen times lighter
than oxygen. One litre of hydrogen at normal temperature and
pressure weighs only 0.08995 gram. The relative lightness or small
density of hydrogen can be shown in a number of ways.

If a small balloon or light sack of any kind which will hold a
gas is filled with hydrogen, the mouth tied, and the balloon set free,
it will rise rapidly in the air, showing that hydrogen is considerably
lighter than air. This is made use of on a large scale by aeronauts
for ascending to considerable heights in the atmosphere. A large
silk balloon is filled with hydrogen, and it will not only rise in the
atmosphere but will carry considerable weight with it. When it is
desired to descend, the hydrogen is allowed to escape through a
valve into the air.

Another method of demonstrating the small density of hydrogen
is the following : Fill a cylinder with hydrogen by displacement of
water and cover the cylinder with a glass plate. Place a second
cylinder filled with air just over the first and remove the plate of
glass. The hydrogen will rise from the lower into the upper cylin-
der and displace the heavier air, which will fall into the lower cylin-
der in which the hydrogen was originally present. This can be
proved by touching a match to the mouths of the two cylinders after
they have been separated. A small explosion in the upper cylinder
will show that it contains the hydrogen, while the absence of any
appreciable hydrogen in the lower cylinder is shown by the fact that
it will not take fire and bum.

A still more striking illustration of the small density of hydro-
gen is shown by an experiment based upon the rate at which hydro-
gen gas diffuses. There is a well-known law connecting the rates at
which gases diffuse with their densities.

Oases diffuse with velocities which are inversely proportional to the
square roots of their densities.

42

PRmCIPLES OF INORGANIC CHEMISTRY

The lighter the gas, the more rapidly, then, will it diffuse. That
hydrogen diffuses rapidly can be shown by the following experiment
(Fig. 8) : A hollowj porous cup C is fastened to a glass tube M,
which extends into the flask Fj passing through a stopper which
tightly closes the mouth of the flask, A second glass tube T, drawn
out to a fine opening, passes through a second hole in the stopper and
dips beneath the water in the flask* A large glass vessel V is now

filled with hydrogen and placed over the porous porcelain cup.
Hydrogen diffuses rapidly in through the cup, due to the small den-
sity of the gaft, proilaces a pressure inside the apparatus, and this
forces the water up into the glass tube T^ and out through the small
opening. Iti this way quite a fountain can l>e produced.

Hydrogen is only slightly soluble in water, 100 volumes of water
at 15^ dissolving only 1.9 volumes of hydrogen,

^e Liquefaction of Hydrogen. — Hydrogen like oxygen was one
of the few gasea which resisted liquefaction until quite recently.

HYDROGEN 43

It was, therefore, placed by Faraday and the earlier investigators
among the "i>ermanent gases." Like oxygen it was subjected to
enormous pressures by Natterer and others, but they were not able
to liquefy it because it was not cooled to its critical temperature.
The critical temperature of hydrogen is very low indeed, — 242**,
and at first sight it is not easy to see how such a temperature can be
reached. When liquid oxygen is allowed to evaporate under small
pressure, a temperature of — 210° to — 220® can be secured, but this
is still above the critical temperature of hydrogen. If, however,
hydrogen under a pressure of several hundred atmospheres is cooled
to — 200° or — 220° and then is suddenly allowed to expand, it will
in expanding cool itself to its point of liquefaction. The critical
pressure of hydrogen is less than 20 atmospheres.

The liquefaction of hydrogen in appreciable quantities we owe
almost entirely to Dewar. He has established its boiling-point to be
— 252°. It is, however, possible to reach a still lower temperature
by a method which has now become familiar to us. By allowing
liquid hydrogen to boil under greatly diminished pressure, still fur-
ther cooling is produced, and a temperature as low as — 258° has
been realized. Under these conditions the hydrogen solidified.
The freezing-point or the melting-point of hydrogen has been shown
by Dewar to be — 258°. It should be observed that this is only 15°
above the absolute zero.

We naturally ask the question. How can such loio temperatrires hi
measured f All ordinary forms of thermometers are, of course, use
less long before any such temperatures are reached, alcohol solidify-
ing easily in liquid oxygen. Even the air thermometer, based upon
the change in volume of air with change in temperature, fails at
such low temperatures, because the laws of gas-pressure do not hold
near the point of liquefaction of a gas, and air is easily liquefied by
contact with liquid hydrogen. •

The best form of thermometer for measuring such low tempera-
tures is what is known as the platinum thermometer. This is based
upon the fact that the resistance to the passage of an electric current
offered by a metal wire changes with change in temperature. The
lower the temperature the less the resistance offered to the passage
of the current. A platinum wire is used for several reasons, one of
them being that platinum is not acted upon by many substances.

Even the platinum thermometer at such low temperatures is not
capable of measuring the temperature very accurately, since relations
which obtain at higher temperatures probably do not hold accurately
in the regions of such extreme cold. It is^ however, probable, all

44

PRlKCtPLES OF INORGANIC CHEMISTEY

things ponsidered, that temperatures such as those of liquid hydrogen
can he measured to within a few- degrees.

Can the Absolute Zero be realized Experimentally. — The further

queatiou arises^ H there any possibility of reachiug the supposed

I absolute zero? ^Ye are compelled to auswer that as far as we can

I see at present there is no such possibility. There is only one sub-

letance known (helium) winch boils lower than hydrogen, and this,

probably t only slightly lower. Further^ the amount of helium which

can be obtained is apparently so small that we can scarcely hope to

Tiae it for obtaining much lower temperatures than those alrearlj

realized.

If helium existed in sufficient quantities, we could liquefy it,
allow the liquid to boil under diminished pressure^ and in this way
secure a temperature several degrees below the boiling-point of this
element From what is known at present of the probable boiling-
point of helium, it is safe to say that even under these conditions a
temperature as low as — 273^ could not be realized. In order that
this temperature should be reached, some substance must be dis-
covered whose boiling-point is considerably below that of hydrogen.
Until such Is obtained it is idle to predict the realization experi-
mentally of the supposed absolute zero of temi>erature, — 273°*

Properties of Liquid Hydrogen, ^ — Liquid hydrogen is colorless
and traiispareut and has small viscosity. The supposed blue color
of liquid hydrogen is due to impurities. It has a density of 0.07,
water being unity. By contact with liquid hydrogen, oxygen (and
as we shall learn also air) is converted first into a liquid and then
into a solid^ or is frozen^ as we say.

A beautiful and thrilling experiment has been performed by
Bewar, who has liquefied hydrogen by the litre. Liquid h^'drogen
WMB poured into a test-tube and the tube exposed to the air* Liquid
too began to stream off the test*tube, and finally the tube became
md. with frozen air. The remarkable character of this ex peri-
L is evident to any one*

Speotmm. — The spectrum lines of hydrogen are

ic. A capillary glass tube is enlarged at Ix^th ends

^ exliatisted. It is then filled with hydrogen gas at a

i mled off. An electric discharge is parsed through

1. tte two platinum terminals fused into the two

B. ne light emitted when the discharge is passed

I {Rtrplish red. When this light is analyzed

two very bright lines and one fsunt

HYDROGEN 45

The spectroscope consists of a prism through which the light is
passed. Light of different wave-lengths is refracted differently, and
we have a separation of the several wave-lengths from one another.
When white light is viewed through a spectroscope, it is broken up
into the spectrum colors. When the light emitted by a gas is passed
through a spectroscope, bright lines appear, and not a continuous
spectrum. The light emitted by hydrogen when analyzed spectro-
scopically shows a bright green and a bright red line, and a faint
line in the violet. Lines in exactly these positions are shown by
no other substance, ^\^len sunlight is analyzed spectroscopically,
we find dark lines in exactly the positions occupied by these bright
lines of hydrogen. These are also due to hydrogen, and illustrate the
general principle that a gas absorbs exactly the same waveAengths which
it can itself emit. This is the law of Bunsen and Kirchhoff. White
light from the interior of the sun, passing through hydrogen in the
exterior, has those wave-lengths absorbed which the hydrogen itself
can emit. By means of the spectroscope, and by this alone, are we
able to prove the presence of hydrogen and other terrestrial elements
in the sun.

Another form of spectroscope should be referred to. When white
light is thrown upon a metallic surface containing a great number of
parallel lines, it is dispersed to even a much greater extent than
when passed through a prism. The concave-grating spectroscope of
Rowland has proved of incalculable service in spectrum analysis.

ElectrolyuB of Hydrogen. — J. J. Thomson, by means of electroly-
sis, separated hydrogen into a positively and a negatively charged
constituent A glass tube across whose centre was placed a loosely
fitting aluminium septum was filled with hydrogen. This was
subjected to an electrical discharge from the two platinum electrodes
fused into the two ends of the tube. After a time the spectrum of
the hydrogen on the two sides of the septum was observed. On the
one side the green hydrogen line was very prominent and the red
faint, while on the other side the red line was very prominent and
the green light faint.

The hydrogen gas was thus electrolyzed into a positive and a
negative constituent, the one being characterized by the strong green
line, and the other by the bright red line. The hydrogen molecule,
like the molecules of chemical compounds, is, therefore, made up of
a positive and a negative constituent.

CHAPTER V

WATER AND H7DROOEN DIOXZDB

Oocnrrence of Water. — Water is probably the best known chemi-
cal compound, on account of its very wide distribution over the
surface of the earth. In the free condition it covers about three-
fourths of the surface of the earth. Further, it is widely distributed
through the rocks over the surface of the earth, each cubic metre
of rock containing on the average about one litre of water. It exists
in large quantities in the atmosphere, in the form of water-vapor.
It also exists in combination with a large number of substances as
water of crystalUzcUion, or water of hydration. Its presence is not
limited to inorganic or inanimate nature. It forms an essential part
of all living matter. If living matter, animal or vegetable, is heated
above one hundred degrees, there is an enormous loss in weight, and
this is mainly due to loss in water which is driven off. The main
constituent of living matter as far as mass is concerned is water.
The human body is more than two-thirds water, and the animal and
vegetable food which we eat contains scarcely less water in proportion
to solid matter. We can thus see why animals can live without
food much longer than without water, and why water is absolutely
essential to vegetable and animal life.

Water as it occurs in Nature is Impnre. — It is safe to say that
all natural water contains impurities. This does not refer to impuri-
ties which are thrown into water artificially, as by the drainage of
human habitations, but to impurities which we may call natural.
The water of the sea is very impure because of matter dissolved
from the soil and rocks by the waters before they reach the sea, and
after they have been poured into it. The waters of small streams
and rivers are impure for the same reason. They dissolve a part
of the solid matter over which they run and with which they may
otherwise come in contact, and carry it along in solution as they
make their way down to larger bodies and ultimately to the sea.
It is, then, obvious that any water which has come in contact with
the earth would be impure. There are, however, very different

46

WATER AND HYDROGEN DIOXIDE 47

degrees of purity represented by terrestrial waters. If the water
has come in contact with certain substances, it will be very much
more impure than by contact with other substances. If water has
come in contact with soil containing a large amount of limestone,
and especially if there is much organic matter in the soil, it will
dissolve large quantities of the limestone and is then what we call
hard water.

If, on the other hand, the water has fallen upon a region which
contains mainly sandstone or other difl&cultly soluble rocks, but little
of the solid matter will dissolve, and we have then comparatively
pure water. This is the reason why water from mountains com-
posed of sandstone is relatively pure. Further, water from mountains
comes in contact with relatively little soil, and rocks in general are
much less soluble than soils.

While it is obvious that water which has once fallen upon the
earth must be more or less impure, the question might reasonably
be asked, Is not rain-water which has never come in contact with the
soil fairly pure ? This question is the more reasonable since it is
known that when water is evaporated, as by the heat of the sun, from
the sea and land, most of the impurities remain behind.

Rain-water would undoubtedly be fairly pure were it not con-
taminated while in the atmosphere. However, while it exists in the
atmosphere in the form of vapor it takes up many kinds of impuri-
ties, and especially after it is formed into drops and falls through
the atmosphere, some foreign matter is dissolved by it. Indeed,
according to a well-established theory raindrops form around dust
particles in the atmosphere and carry these particles down with
them as they fall. This is undoubtedly the reason why the atmos-
phere seems so pure after a heavy rainstorm.

From the above it is then obvious that all natural waters contain
impurities, but that the amount of impurity varies greatly from one
sample of water to another.

Mineral Waters. — In certain localities minerals exist which are
more or less soluble in water. Eain-water or water from other
sources dissolves these substances and holds them in solution.
Such waters are known in general as mineral waters, the nature
of the water depending upon the nature of the mineral in solution.
At the great sources of mineral waters, such as Saratoga, New York,
there are beds of various salts deposited beneath the surface of the
earth, some of them nearer the surface, others at much greater
depths. When the waters percolate through such regions, they
come in contact with these various deposits, and dissolve more or

48

PRINCIPLES OF INORGANIC CHEMISTRY

less of the different substanoes, the amouBts dissolved depending
upon the relative solubilities.

These substances give characteristic tastes and other properties
to the water in wlueh they are present, and thus we have the various
mineral waters which are so well kaow*n and so anuch sought after,

If the water percolates through a soil eontiiiuing large amounts
if carbon dioxide set free from decomposing vegetable matter or
from other sources, and espeeially if it comes in contact with carbon
dioxide under high pressure, large amounts of the gas may dissolve
ill the water and be carried by it to the surface of the earth or to a
mineral well wliich has been bored l>elow the surface. Stich waters
are known as effervescentj siuee they give off a paii; of the dissolved
carbon dioxide when exposed to the air. In other cases the water
dissolves considerable hydrogen sulphide and gives us what ia
kuown as sulphur ivater. When other substances are dissolved in
the water, they give their characteristic properties to it, and thus we
have the almost endless variety of mineral waters which are present
upon the market w^ith their ludicrously pedantic names,

Fnrification of Water — Water is usually rendered impure in the
way described above, by carrying \^nth it in soKition dissolved sub-
stances. It may, however, be rendered impure by matter which is
not in solution, but simply in a state of mechanical suspension.
This latter condition is illustrated by small streams after a heavy
rain. The finely divided soil is carried along with the water in a
state of fine suspension, and we have mudd^ water.

When the impurity is in a state of mechanical suspension and is
not in solution, it can be removed Yyy JUlnilhn, Filtration consists
in passing water through a substance with very fine openings or
pores, so ixne, iudeed, tliat the particles of water can pass, but not
the particles held in mechanical suspension. This is effected on a
small scale in the laboratory by means of certain varieties of paper
kuown as ** tilter paper." If it is desired to purify on a large scale
water which contains foreign matter in suspension, some other
device must be resorted to. It Is sometimes passed through a thick
layer of very fine sand, and other substances have been used.

Filtration is of fundamental importance to the chemist, especially
in connection with analytical operations. Quantitative analysis de-
pends largely upon the precipitation of the substance whose quantity
w^e wish to determine in the form of a solid. This solid must then
\}e filtered off from the liquid which is present and carefully washed
before it can be dried or ignited and weighed.

If the impurity in the water is in solutioiii it is obvious that we

WATER AND HYDROGEN DIOXIDE

49

cannot separate it by any mechanical process such as filtration.
Some other principle must be utilized. When water containing non-
volatile impurities is boiled, the vapor which escapes is practically
pure. If this vapor is condensed again, we have practically pure
water. This process of converting a liquid into vapor and recon-
densing the vapor is known as distillation^ and the apparatus in
which a distillation is carried on as a still.

Distillation, like filtration, is of fundamental importance in the
chemical laboratory. The water which is furnished a chemical
laboratory from natural sources is not sufl&ciently pure to be em-
ployed in any chemical operation. Before it can be used it must
be distilled, and in all chemical work only distilled water is
employed.

The form of still which is used when only a small amount of
liquid is to be distilled is shown in Fig. 9. Into the glass flask

Fig. 9.

F, the liquid to be distilled is introduced. This is heated and eon-
verted into vapor by a burner placed beneath the flask. C is the
condenser, consisting of a small inner glass tube surrounded by a
much larger glass jacket. Cold water is passed into the jacket at a,
and out at h. The vapor in the inner tube is condensed to a liquid
as it passes through the condenser, and flows into a receiver, R,

If it is desired to distil a liquid on a large scale, the form of the
apparatus is greatly modified, but the principle is exactly the same
as in the apparatus described above.

Another method of purifying water is by freezing it. Just as
the vapor which separates from impure water is pure, just so the ice
which freezes out of impure water is practically pure. When impure
water is partly frozen in a quiet place, the ice which separates con-

nUNCirLES OF INORGAXIC CUEMISTUY

tAtni iiuioh Ir»i! iiupiirit^ than the water from which it separated,
lUn iiHUHvily ifMmiining for tJie most part iu the unfrozen water.

Tliiw liHithdil of frmximj, ur cttfslttUization, is far hns efficient and
luiii^ll mImwoi' Iu carry out than the method of distillation^ but has
vwluti ill Ih** hilKvrnttiry in certain connections.

Water not an Element, but a Compoiind, — Water is the first snb-
niarH'n wliu ii we hiive thus far studied wliich is not an element, hut
li coni|»oHi'(l of more than one element. For a long time in the early
*liiyH uf chemistry water waa regarded as an element, and together
with uirt earth, and fire constituted the four chemical elements.

Tliat water is not an element is obvious from our studies of ox}--
p*n and iiydrogen. We have seen that by electrolysis both oxygen
and hydrogen can be obtained from water j an<i an element, by defini-
tititi, is a substance which cannot be decomposed Into any other
gubstances.

Camposition of Water — We have seen that oxygen and hydrogen
can Ije obtained from water^ but this does not show that water con-^
tains only these two elements. To answer this question two general
methods are available. First, decompose water, and kcc whether
anything bnt hydrogen and oxygen is obtaine<l Second, cause oxy-
gen and hydrogen to combine, and see whether water is formed.

The most convenient means of decomposing water is the electric
current. When a little acid is added to waU»r to iiijuinish its resist-
auce to tlie flow of the current, and an electric cyrrent is passed
through it, it is decomposed. This process of effecting decomposi-
tions by means of the current is known as ehvtroitfin'ii. The metallic
terminals, or polesj have specific names with whicli it is ijnjiortant
to be familiar. The pole which the current leaves and enters the
solution is known as the mtode ; the pole which receives the current
from the solntion, as the cathode.

The only products obtained by the electrolysia of water, are the
two gasesj oxygen and hydrogen, oxygen being set free at the anode
and hydrogen at the cathode. That these are oxygen and hydrogen,
reajieetively, can lie shown by the fact that the former will ignite a
match which has just been extinguished, and the latter will burn
with the characteristic hydrogen fiame.

If we wish to know the relative volumes of the two gases set free
from water, we must collect and measure them.

A conveuient form of apparatus for effecting the electrolysis of
water and collecting the gases set free is the frtllowing: —

Into the two arms ^1 and B (Fig. 10) of the U-tube are inserted two
platinum electrodes. These tubes are completely filled with acidu-

WATER AND HYDROGEN DIOXIDE

51

lated water by filling the reservoir R to the desired height, and
opening the two stop-cocks at the ends of A and B, The current is
passed into the solution through the electrode in B and out through
the electrode in A, The
stop-cocks are closed be-
fore the current is passed,
and oxygen collects in B,
and hydrogen in A When
the current has been flow-
ing for a short time it will
be observed that the gas
is collecting in A faster
than in B. The tubes A
and B are graduated so
that at any moment the
amounts of gases set free
can be read off at once.
After an appreciable
amount of gas has col-
lected in B, interrupt the
current and read the vol-
umes of the gases in the
two tubes. It will be
found that there is just
tioice the volume of gas in
A that there is in B.
Close the circuit, and allow
the electric current to flow
until a considerably larger
volume of the two gases
has been set free. Inter-
rupt the current again and measure the volumes of the two gases.
It will be found that the volume of the hydrogen is again exactly
double that of the oxygen.

No matter how long the current is allowed to flow, nor how much
water is decomposed, we would always find that the volume of the
hydrogen set free was exactly double that of the oxygen. From the
decomposition of water, or by the analytical method^ we are therefore
led to the conclusion that water is made up by the union of two vol-
umes of hydrogen with one volume of oxygen. Since one volume of
oxygen weighs 15.88 times one volume of hydrogen, the proportions
of hydrogen and oxygen by weight in water are 1 : 7.94. To

Fio. 10.

PRlNCirLES OF INORGANIC CHEMISTRY

determine the compos Itioa of water^ however, we are not dependent
solely upon the analytical method. We can use also the si/ntheticuL
If water is composed of two volumes of hydrogen to one volume
of oxygen, then, when we mix two volumes of hydrogen with one
volume of oxygen and pass an electric spark through the mixture,
which causes the gaaes to combine, all the hydrogen should combine
with all the oxygen and form water, Tliis is exactly what takes
place. Whenever two volumes of hydrogen are mixed with one
vohune of oxygen and the gases made to combine by means of an
electric sjiark, or by rise m temperature, all the hydrogen and all
the oxygen are used op and water is formed. If more than two
volumes of hydrogen are used, all the oxygen will be used up and
the excess of hydrogen will remain uneombiued. If less than two
vohunes of hydrogen are used, all the hydrogen will be used up and
the excess of oxygen will remain.

Tho results of synthesis confirm those of analysis; viz. that
water is formed by the union of two volumes of hydrogen with one
volume of oxygen.

Chemical Behavior of Water* — All things considered, water is
probably the must important chemical compound known. It is
formm]^ as we have seen, by the union of hydrogen and oxygen* It
is also fn^quently formed, as we shall learn, by the union of hydrogen
with the group OH, this group representing already a combination
b<^twecn oxygen and hydrogen. The possibility of the union of
hydrogen in one cQm|>ound with the group OH, known as hydroxyl,
in another comi>oumi, conditions a large number of chemical reac-
tions*, Imleed, were it not for this fact, the science of chemistry
would \m wtty different from what it is to-<lay.

This naturally raises the question, whence this tendency of
Ikj^rogteti to uuite with hydroxyl ■' The answer is to be foutid in the
IMC|;5 tit*Utions which obtain in hydrogen and oxygen on the one
kHd«lxid iu water, on the other. We have seen that when hydrogen
milk oxygen an enormous amount of beat is liberated.
i$ approximately a measure of the diifereoce lietween the
Krdt^^^u and oxygen, and in water. This dilference is
^ ^sittft^ % liife amount of intrinsic energy being converted into
ii^iimlqJU^fPi tad o&ygen combine.

^^%fni^ fmeiple that whenever a chemical reaction which
wii^>ja\,iy ^M«A rf heat can take place, it does so. This is
li^ m^ ^ ^ iBT ^irit diti« is a strong tendency on the part of
ai^^ i^ni^ m f«» lyi^r into heat, and this is, doubtless, the
^i^^^CGoi^^nkKy of hydrogen and oxygen to combine

WATER AND HYDROGEN DIOXIDE 63

whenever an opportunity presents itself. The importance of this
fact for the whole science of chemistry will become apparent as the
subject develops.

Water has the power of combining with a certain class of chemi-
cal compounds known as the oxides, converting some of them into
the important class of compounds known as the bases, and others
into the very important class of compounds known as the acids.

Take the well-known substance lime, CaO. When this is treated
with water it combines with it, forming calcium hydroxide : —

CaO-|-H,0=Ca(OH)2.

As an example of water combining with an oxide, forming an
acid, take the trioxide of sulphur, SOj. When this is treated with
water, the following reaction takes place : —

S03-hHjO=H2S04.

The subject of acids and bases will be considered more in detail
when the elements which form these substances are studied.

Water a Stable Compound. — Few compounds known are more
stable than water. If we try to decompose it into its elements, we
will appreciate what this means. It is true, as we have seen, that it
can be decomposed into its elements by means of the electric current,
but unless an acid is added to it a current of high voltage is required
to effect any appreciable amount of decomposition.

If we try to decompose water into its elements by heat, enormous
temperatures are required. In order to effect even slight decompo-
sition a temperature of 1000® or higher is necessary, and a consid-
erable amount of decomposition is effected only when temperatures
between 2000° and 3000° are employed.

Stress is laid upon these facts for the purpose of illustrating a
general principle. WJien a chemical compound is formed with great
evolution of heat, it is almost always a very stable substance.
Indeed, the degree of stability can usually be measured by the
amount of heat set free during the formation of the substance. We
have seen that the heat set free is an approximate measure of the
difference between the intrinsic energy of the substances before they
unite and the products of the reaction. When there is large heat
evolution during a reaction it means, other things being equal, that
the products contain relatively little intrinsic energy; and since
intrinsic energy is the cause of chemical activity, we should expect
those substances with a small amount of such energy to have little
chemical activity. To say that a substance is relatively inactive

PRINCIPLES OF INORGANIC CHEMISTHV

chemically is, in geiieml, the same as to aay that the substance la
stable, eince stability in the last few years has come to mean
chemical inactivity*

PHYSICAL PROrERTIES OF WATER

Physical Propertiei of Water; Boiling-poiiLt — Water at ordi-
nary tenipeiutures and pressures Is a colorless liquid- In very
thick layers it has a bluish tint Under a pressure of 7G0 miUi-
metres of mercury it boils at 100° C. j Le. at this temperature the
tenHion of the aqueous vapor is just sufficient to ovei^come the
pressure of the atmosphere* As the pressure to which the water
is Bubjected increaaesj its boiling-point rises. Under a presaure
of five atmospberes the boiling-point of water is 152°. Under a
pressure of ten atmospheres water boils at 180.3°j while under a
jiressure of twenty atmospheres it does not boil until a temperature
of 213* is reached.

These are the conditions which obtain in a steam-engine. The
water is under the pressure of its own vapor, which amounts to from
five, to ten or twelve, atmosphereSj and consequently its boiliug-
potnt is very greatly raised. From the data given above, we can
form a pretty close approximation as to the temperature of the
water in the boiler of a steam-engine.

Just aa water must be heated much higher in the boiler of a
fiteam-engjue than in the air in order to obtain boilings just so it
boils much lower on a high mountain than in a valley. Aa we
futeend a mountain the pressure of the atmosphere becomes less,
and, consequently, the pressure which the tension of the water*
vapor must overc4^me* On the top of Mont Blanc water boils at
about 84'.

As we ascend a mountain the pressure of the atmosphere
becomes gradually less, according to a well-known law, so that if
we know the exact temperature at which water Iwiled at the sea-
level at any given time, w^e could use the temperature at which
wat4^r boiled at any higher altitude to calculate approximately the
he i gilt which yve bad reached.

Heat of Vaporiiatlon, — Any one who has ever observed water
boil must have been impressed by the enormous amount of heat
which is required to convert the liquid into vapor. He must have
been further impressed by the fact that the temperature of the
vapor is practically the same as that of the liquid from which it
was formed. The amount of heat required to convert one gram of

■ WATER AND HYDROGEN DIOXIDE 65

water at 100° into vapor at 100° is 540 calories, i.e. the same amount
of heat which would be required to raise 540 grams of water 1°
in temperature. This is known as the heat of vaporization of water.

The Freezing of Water. — When water at ordinary pressure is
cooled to 0^ it freezes, as we say, or passes into ice. As we cool
water down toward its freezing-point, it contracts in volume until
a temperature of 4° is reached. As the temperature is further
lowered from this point, the water begins to expand and continues
to do so until the freezing-point is reached.

The importance of this apparently insignificant fact is very great
indeed, from the standpoint of the economy of nature. Since water
expands from 4° to the freezing-point, ice is lighter than water and
floats upon it. The importance of ice floating upon water and not
sinking to the bottom is twofold. In the first place, ice floating
upon water protects it from the extreme cold of the atmosphere,
since ice is relatively a poor conductor of heat, and in the second
place, if ice sank to the bottom of our streams as fast as it was
formed, this would continually expose a fresh surface of the water
to the cold and our streams might be frozen solid, which would
mean the extermination of all living things within them.

Again, if water became continually heavier as- its freezing tem-
perature was reached, the coldest water would constantly settle to
the bottom, and this would tend to make the stream begin to freeze
at the bottom and finally become solid ice.

The fact that water contracts to 4° and then begins to expand
again, causing the ice to be formed on the Surface of the water and
to float there, explains why the rivers and lakes in cold climates
are frozen only a few feet in depth, and why the forms of life
which inhabit them are not exterminated in one cold winter.

Just as the boiling-point of water is raised by increase in press-
ure, just 80 the freezing-point is lowered when the pressure is
increased. It may not be obvious at first sight why the freezing-
point of water is lowered as the pressure is increased. This can be
seen from the following considerations. Water expands in volume
from 4° to the freezing-point. Anything which will oppose the
expansion will hinder the water from freezing, and under such
conditions it will require a lower temperature to freeze the water.
The effect of pressure on the freezing-point is, however, small.
An increase in pressure of a whole atmosphere lowers the freezing-
point of water only about 0°.007.

A pretty and interesting experiment is based upon the above-
described fact. Ice at 0® can be melted by simply subjecting it

m

PRINCIPLES OF INORGANIC CHEMISTRY

to pressure* This experiment wiis shown bj TyndaJl in his popular
lectures as follows, A ray of light was paased through a block of
icp whieh was kept at 0^ The ice was then subjected to pressurej
and melting began to take place. The drops of water formed in the
interior of the block of ice^ having different refract ivity from the
ice^ could be readily seen when the light was thrown on a screen.

lee has been subjected to such high pressures that it could be
melted at - 18^

Heat of Pnaion of Ice. — We ha^e just learned what an enormous
amount of heat energy must be expended to convert water into
vapor. We shall now see that a large amount of heat is required
to convert ice into water. The amount of heat required to convert
one gram of ice at 0" into water at 0* is 80 calories, i.e. the amount
of heat which would raise one gram of water from 0" to SO", This is
known as the heat of fusion of ice.

Heat of Condensation of Steam and of Solidification of Water. —
We have spoken of the heat of vaporization of w^ater being such a
large quantity. Just as we have to add a large amount of heat
energy to water to convert it into vapor, just so when we reeoudeuse
the vapor to liquid an enormous amount of heat is set free. This is
knnwu as tlie h^at of condenmtion of water. We would naturally
Sink what relation exists between the heat of vaporization and the
heat of condensation of a substance ? 77* e two have been fomid by
tjcperhiient to be eqnaL That they must be equal follows from the
law of tlie conservation of energy* Starting with water we convert
it into vapor and then recondense the vapor to water. The initial
Htiil thukl stages are the same, and the amount of energy added to
vft%M^\ umi part of the transformation is given up again when the
irvnrwv triiiJKriirination takes place.

\Vm Imvt^ mwu that considerable heat energy must be added to ice
Ui hipirfy it When liquid water passes over into ice a large amount
id oiutrjfy IN given up in the form of heat. This is known as the
kmt nf mitit({ffaftioiu The heat of solidification of water is exactly
ih|UmI to tin* lu'at uf fusion of ice, as has been shown e x peri men tally »
HUd rt« mth he Mhuwii from the conservation of energy by reasoning
t»liiiilly MinJiHtoui to that used above in the case of heat of conden*
inliun

MupiihtaUuif nud Supercooling of Water. — If water is heated as
U \^*m\\y tluufs witlHMit Uikiug any precautions to remove air and
|ilh<4i UnpuntU>M, it brgins to boil at 100" if the barometer stands at
Tihi iHillhiii^lvt^v of nvHTUry, or slightly above or below this tempera-
iU(ii| dM|xiuUii!g upon whether the barometer is higher or lower than

WATER AND HYDROGEN DIOXIDE 57

the normal pressure. If, however, suitable precaution is taken to
purify the water, and to remove all air from it by warming it in a
vacuum, it may be heated considerably above 100° without boiling.
If water which has been very carefully purified is warmed in a new
glass flask, which itself has been carefully cleaned and is free from
scratches or any irregularities, it may be heated many degrees above
its normal boiling-point without boiling. Water in this condition is
said to be superheated, Wlien boiling once begins it is liable to take
place with explosive violence, as we would expect.

In order that water may be superheated to any appreciable extent
it is necessary to take the precautions indicated above, since if there
are any impurities present or any irregularities in the vessel used,
these will serve as points from which the boiling will begin as soon
as the temperature of 100® is reached. The vapor forms readily
around such impurities, especially if they are gaseous, as air, or on
irregularities, and it is impossible to superheat the water to an
appreciable extent. Superheated water is an unstable condition of
this substance, and readily passes over into the condition stable at
this temperature, t.e. vapor. Just as water may be heated above
its boiling-point without ebullition taking place, just so it may be
cooled below its freezing-point without the separation of ice. When
pure water is cooled below (f without any ice separating, it is said
to be supercooled. It is much easier to supercool than to super-
heat water. If ordinary distilled water is placed in a smooth glass
tube, which is inserted in a mixture of ice and salt (a freezing
mixture which may have a temperature as low as — 20°), and gently
stirred as it cools, a temperature of —4° to —5° may be reached
without any ice separating. If to such supercooled water a fragment
of ice as small as can be seen is added, more ice will at once begin
to separate, and by vigorously stirring the mixture of ice and water,
ice will continue to separate until heat enough is liberated to warm
the remaining water up to its true freezing-point, all supercooling
being thus removed.

The amount of the solid substance necessary to cause more of the
solid to form is so small that we cannot conceive of it, as the German
physical chemist, Ostwald, has shown. It is worth noting that in
order to remove the supercooling of a liquid it is necessary to use
some of the same substance in a solid state. If any other solid is
used the freezing is not liable to begin and the supercooling may re-
main ; supercooled, like superheated water is an unstable condition.
Water does not readily remain in these conditions, but passes over
under slight provocation into the condition which is stable at the

58

PRINCIPLES OF INORGANIC CHEMISTRY

temperature in question. From supercooled water ice thus readily
separates, just as vapor forms with the greatest ease from super-
heated water.

The Yapor-tention of Water in its Different States of Aggregir
lion. — The tension of water-vapor is, as we would expect, greatest
when the water is in the form of liquid, and least when the water
is in the solid state. Ice, however, has a vapor-tension at all
ordinary temperatures, and water always has an appreciable vapor-
tension at all temperatures at which it remains in the liquid state.
The vapor-tension of water at different temperatures and in dif-
ferent states of aggregation has been carefully measured, since the
results frequently come into play in many scientific investigations.
When these results are plotted as ordinates against temperatures as
abscissas, curves are obtained which are of great scientific value,
as we shall now see.

The Temperature-pressure Diagram of Water. — In the following
diagram (Fig. 11) the ordinate represents the pressure of water-vai>or

and the abscissa tern-

B

perature. Water ex-

ists, as we have seen,

\

as a solid, a liquid, or a

\

gas,dependingchiefly

\ LIQUID

upon the tempera-

\

ture, and also upon

SCUD \

the pressure. These

\

states of aggregation

\ ^^

are known as pivasesy

\^^,^X^^

and water is said tx)

Pr --^/T?'^^

exist in three phases.

1 ^"^ VAPOR

If we draw the

temperature-pressure

r»iirvp» ppnrftSPiit.incr

TEMPLRATURC

the conditions of

FlO. 11.

equilibrium between

the different jiliases of water, the curves will have the form seen
in Fig. 11.

The curve PA represents the condition of equilibrium between
liquid water and water- vai)or. Below this curve the vapor is the
stable phase, above it the liquid. The curve PB is the line of equi-
librium l)etween the liquid and the solid phases of water, the liquid
being the stable j)hase to the right of this curve and above the curve
jPJ, while the solid is the stiible phase to the left of PB and above

WATER AND HYDROGEN DIOXIDE 69

PC. The curve PC is the line of equilibrium between the solid
phase of water and water-vapor ; above this curve and to the left of
PB ice is the stable condition, while below this curve and PA water-
vapor is the stable phase.

It will be observed that the three curves intersect in a point
which we have called P. This point has properties which make it
of special interest. Since it is common to all three curves it means
that at this temperature all three phases of water have exactly the
same vapor-pressure. That such is the case can be shown by the
following considerations. Take the liquid and solid phases. The
point P represents the temperature at which ice and water are in
equilibrium under their own vapor-tension. Since this is much less
than an atmosphere, being in fact about four millimetres, the tem-
perature of the point P is slightly above zero, since pressure lowers
the freezing-point of water. If the vapor-tension of the ice is not
the same as that of the water, it must be either greater or less. If
it is greater, the ice will vaporize and the vapor condense as liquid ;
if it is less, the water will vaporize and the vapor freeze to ice.
Since, however, by hypothesis this point represents a condition of
equilibrium between these phases, where neither can increase at the
expense of the other, we could not have either of the above condi-
tions realized. Therefore, since the vapor-pressure of the ice cannot
be greater than that of the wat«r at this temperature, and cannot be
less, it must be equal to it A special name has been given to the
point P. Since it represents a condition of equilibrium between
three phases, it is known as a triple point The curves PA^ PB, and
PC represent conditions of equilibrium between two phases, and the
areas PAB, PBC, and PCA represent conditions under which only
one phase is stable.

The Phase Bnle of Willard Ctibbs. — We can now state and apply
a generalization of wide-reaching significance and of great importance,
which holds for conditions of equilibrium such as those with which
we are now dealing. This generalization, which was discovered by
J. Willard Gibbs, is known as the Phase Rule,

We are dealing with one component, water, and three phases, — the
solid, liquid, and gaseous. If the number of phases exceeds (he number
of components by two, the system is non-variant, or has no degree of
freedom. This means that none of the conditions can be varied
without destroying the equilibrium. The triple point P is a non-
variant system. The number of phases is three and the number of
components one, and we cannot vary either the temperature or the
pressure without disturbing the equilibrium between the three phases.

»1 •«>■•»"

"-^il^w- .It* "iiTllaltj.f

. — -r -i'l: :-

. ■. .. ■i.r-i; :.:; :i ".::»• .iiiil'L*? .f

- • ■ ^.- ■.-. •: ■ . :--rL'v[ . -l:e
., -^^ ■. - .'. ■■■::' -i* ■:;:i:i;.>:ri-

». ■ ■■. "^ N ■■..■'. li^rrtllTS

I ... v.»v.. :» u^iiiw>. .\.i;vL- j..i.> vr^.L ably the hij^hesl

WATER AND HYDROGEN DIOXIDE 61

specific heat. By the specific heat of water is meant the amount of
heat required to raise a given amount of water, say a gram, one
degree in temperature. This amount of heat, as we have seen, is
one calorie. One calorie of heat will raise a gram of any other
known substance more than on© degree in temperature. Water thus
stands at the head of the list as far as specific heats are concerned.

If we examine the specific inductive capacity or the dielectric con-
stants of liquids, we find water either at the very extreme or very
neaily the extreme. It seems probable that there is one liquid with
a higher dielectric constant, but if so, its dielectric constant is not
much higher than that of water.

Similar results would be obtained if we ran through the whole
list of properties. Water would stand in practically every case either
at the top or bottom of the list of substances. Its properties are,
therefore, distinctly extreme. They are either a maximum or a
minimum, and usually a maximum.

If we take all the properties of water into account, we shall see
that we are easily justified in regarding it as the most remarkable
chemical compound known.

Solvent Power of Water. — Water has a remarkable power to
dissolve other substances which are brought in contact with it. In-
deed, of all known substances it is the best solvent, and with respect
to this property it also stands at the very head of the list of chemical
substances. The importance of solution for chemistry cannot be
overestimated. This becomes obvious when we consider that most
chemical reactions take place in solution. Indeed, comparatively
few solid substances are capable of reacting with other substances
in the solid state. Were it not for solution the whole science of
chemistry would be very different from what it is to-day, and far
less interesting. Three-fourths, and probably a much larger propor-
tion, of the chemical reactions with which we are now familiar would
not take place at all. The reason for this we shall learn a little
later.

Water dissolves to a greater or less extent not only most solid
substances which are brought in contact with it, but also most
liquids and gases.

TJnsaturated, Saturated, and Snpersatorated Solutions. — When
water has dissolved a certain amount of a given substance, but is
still capable of taking up more of it, the solution is said to be un-
saturated. When water has dissolved all of a given substance
which, at the temperature in question, it can take into solution, the
solution is said to be saturated. When water contains more of a

62 PRINCIPLES OF INORGANIC CHEMISTRY

given substance than it can hold in a stable condition, the solution
is said to be supersaturated.

A saturated solution can be prepared by two methods. First;
bring the substance to be dissolved, in excess, in contact with the
solvent, and agitate the liquid until it will take up no more of the
dissolved substance. This method is slow, and requires a long time
for equilibrium to be reached. The second method consists in
warming the solvent to a considerably higher temperature, and
agitating it at the more elevated temperature with an excess of the
substance to be dissolved. It is a general rule to which only a few
exceptions are known, that substances are more soluble at higher
temperatures than at lower. The solvent at the higher temperature
takes up more of the substance than is sufEcient to saturate it at
the lower temperature. When the solution is cooled down to the
temperature desired in the presence of some of the solid substance,
the excess of the dissolved substance separates in solid form, and
the solution is saturated at the temperature in question.

The results obtained by the second method are always a little
higher than those obtained by the first, and it is well to use both
methods and take the mean of the results obtained from the two.

To prepare a solution supersaturated at any given temperature
we warm the solvent in contact with the substance to be dissolved
to a somewhat higher temperature, and allow it to take up all the
substance that it can. Every trace of the excess of solid, undis-
solved substance is then filtered off, and the solution practically
saturated at the higher temperature cooled down to the desired
temperature. If we are careful to avoid agitating the solution
during cooling, we will have a solution supersaturated at the lower
temperature.

To determine whether a solution is supersaturated, add a few
fragments of the solid phase of the dissolved substance. If there is
supersaturation more of the dissolved substance will separate in
solid form, until all supersaturation is removed. This explains why
it is necessary to filter off all the solid matter in preparing a super-
saturated solution.

Limited and Unlimited Solubility. — We have every degree of
solubility represented. Many solids are soluble in water to only a
very slight extent, while other solids dissolve in much less than their
own weight of water. There is no solid known which dissolves in
water to an unlimited extent.

Some liquids are scarcely soluble in water at all, while others
are miscible with water in all proportions. Thus, the oils are very

WATER AND HYDROGEN DIOXIDE 63

slightly soluble in water, while the alcohols are miscible with, or
what is the same thing, dissolve in water in all proportions.

Gases show very different degrees of solubility in water. Some
are only slightly soluble, while others dissolve in very considerable
quantities. For any given gas the solubility varies with the
pressure — the greater the pressure the greater the solubility. A
simple relation was discovered by Henry connecting the solubility
of a gas with the pressure, and which has come to be known as
Henry* 8 Law, Hie amount of a gas dissolved by a liquid is propor-
tional to the press^ire to tchich the gas is subjected. Henry's law has
stood in general the test of experiment, but there are exceptions
known to it, especially when the gas is quite soluble.

No gas is soluble in water to an unlimited extent.

These relations which have been applied to water as a solvent
hold also for other liquids.

Properties of Water affected by Dissolved Substances. — Certain
properties of the solvent are very greatly affected by dissolved
substances. The freezing-jmnt of water is lowered by dissolved sub-
stances, and this is perfectly general no matter what the nature of
the substance which is dissolved in the water. Similarly, the
boiling-point of water is raised by the presence of substances dissolved
in it. Since boiling-point varies inversely as vapor-tension, this is
the same as to say that the dissolved substance lowers the vapor-
tension of the solvent. This is also a general effect, independent of
the nature of the dissolved substance.

The power of water to conduct tJie electric current is also greatly
affected by the presence of certain dissolved substances. Pure water
is almost a non-conductor of the current. A cubic millimetre of the
purest water which has been thus far prepared, offers the same resistance
to the passage of the electric current as a copper wii'e whose cross-section
is one square millimetre, wrapped around the earth one thousand tim£8.
When certain classes of substances are dissolved in water the solu-
tion becomes a good conductor, while other substances do not impart
this property to water. Those substances whose solutions conduct
the current, which we shall learn to know as acids, bases, and salts,
are called electrolytes; while those substances whose solutions do not
conduct the current, including all except the above compounds, are
known as non -electrolytes.

The Dissociating Power of Water. — The question naturally arises.
Why do solutions of some substances conduct the current, and solu-
tions of other substances not conduct ? This question is much more
easily asked than it is answered. It has been found by elaborate

64

PRINCIPLES OF LX0RGAI5IC CHEMISTRY

experimental investigations tbat all those substances which, when in
solution, conduct the current, and only those, produce greater lower-
ing of the freuziug-point and greater lowering of the vapou-t^nsion
of water than the substances whose solutions do not conduct the
current. We shall learn that the amount by wliieh the freezing-
point or the vapor-tension of a solvent is lowered, depends only
upon the ratio bet\veen the number of parts of the solvent and the
number of parts of the dissolved eubstance. If one substance pro-
duces a greater lowering of freezing-point or of vapur*teiisioa than
another at equal concentration, it means that its solution contains
•a larger number of parts. From this and otlier lines of reasoning
which will be considered a little later^ we are forced to the conclu-
sion that water (and also other solvents) has the power of breaking
down the molecules of certain substances which we have called elec-
trolytes into i>arts. The parts are^ however, not simply the atoms,
but the atoms or groups of atoms charged with electricity. These
charged parts are known as mis^ and the breaking down of molecules
into ions as dis^iociaiion. Since the ions are the carriers of electricity
and separate at the poles in electrolysis, this kind of dissociation ia
kno\vn as eiedmhfth dissocialion.

When a molecule is electrolytically dissociated by a solvent like
water, one of the ions ia always charged positively and the other
negatively. The positively charged ion is called the eatiouj and the
negatively charged ion the anion.

HYDROGEN DIOXIDE

Hydrogen Dioxide. — One compound of hydrogen and oxygen
other than wati-r calls for special comment This is the compound
hi/drofjen dioxide^ winch, as the name implies, contains more oxygen
than water. It has the comi>ositiou expresssed by the formula II^Oj,
and is, therefore, an oxidized water. It probably occurs in the
atmosphere, but only in very small quantities.

Preparation and Puriflcatioa. — Hydrogen dioxide is most readily
prepared in atiy quantity by treating a compound with which we
had to deal when we were studying oxygen, called barium dioxide,
with an acid. If we use hydrochloric acid, the equation is expressed
thus : —

BaO, + 2 HCl = BaCl, + H,0^

If we employ sulphuric acid^ thus : —

BaO, -h H,S04 ^ BaSO* + H,0^

WATER AND HYDROGEN DIOXIDE 65

The cold dioxide is introduced slowly into the cold, dilute acid,
when the reaction indicated by the equations takes place. In this
manner the commercial product is prepared, and this contains about
three per cent of the dioxide.

The commercial product neai-ly always contains a trace of hydro-
chloric or sulphuric acid. This can best be removed by treating the
dioxide with a little zinc oxide, frequently shaking it and allowing
it to stand for several hours. The solution of the dioxide is then
filtered to remove any excess of the oxide of zinc, and distilled
under diminished pressure to free it from any zinc which has dis-
solved.

If it is desired to prepare a more concentrated solution of hydro-
gen dioxide, this is effected by slowly evaporating the more dilute,
purified solution on a water-bath at a temperature below 75®. Water
boils lower than hydrogen dioxide, and under these conditions
distils out of the aqueous solution of the dioxide, and for the most
part leaves the dioxide behind. The distillation must take place far
below the boiling-point of water, since at 100® hydrogen dioxide
undergoes marked decomposition. Indeed, some of the dioxide
decomposes at 75®, esi>ecially if there is any appreciable amount of
impurity present. A fairly concentrated solution of the dioxide can
be obtained in this manner.

If this solution is now distilled under a pressure of from ten
to twenty millimetres of mercury, a fairly pure hydrogen dioxide
can be obtained. The object of the diminished pressure is to lower
the temperature at which the solution will boil, as we have already
seen to be the result in the case of water. The liquid which first
passes over is almost pure water, containing only a little of the
dioxide, since water boils at all pressures lower than the dioxide.
The liquid which comes over later is almost pure hydrogen dioxide.

This process of separating liquids, based upon differences in
their boiling-points, Ls known as fractional distillation^ and is an
important operation in chemistry, as we shall learn.

Properties of Hydrogen Dioxide. — Hydrogen dioxide is a color-
less, viscous liquid, much heavier than water, having a specific
gravity of 1.4996. We saw that water in thick layers has a
markedly bluish tint. Hydrogen dioxide is still deeper blue when
observed in thick layers.

One of the characteristic properties of hydrogen dioxide is the
ease with which it decomposes into water and oxygen.

2 HA = 2 H,0 -h 0^

QG

PROTCIPLES OF INORGANIC CHEMISTRY

This is the reason why such precautions have to be taken in purify-
ing it to prevent decomposition.

The eoQcenfcrated hydro|jen dioxide is very explosive, on Eujcount
of the ease with which the above deeomposition takes plac^e. When
it is brought in contact with certain solid substancesj such as the
metals or metal oxides, the decomposition is so rapid that violent
explosions result.

Hydrogfen DioKide a Good Oxidizing^ Agent — On account of the
ease with wliioh hydrogen dioxide gives up oxygen it is a good
oxidising agent. If brought in contact with substances which can
take up oxygen it parts with it readily, and such substances are
oxidij:ed. Upon this fact is based its value as a drnftfed-auL When
brought in contact with organic matter it oxidizes it and destroys
its viL'tlitj'. Bacteria and other germs which produce disease are
thus destroyed by hydrogen dioxide.

As an oxidizing agents it has various applications in the field
of chemistry. This property enables it to be readily detected,
Wheu hydrogen dioxide is brought in contact with a solution of
potassium iodide^ oxidation takes place, converting the potassium
into potassium hydroxide and liberating the iodine.

2KI + HA = 2KOH4-2L

The iodine is recognized at once by its brown color, or by its power
to color starch paste blue.

Hydrogen Dioxide also a Reducing Agent. — We have seen that
a reducing agent is one that adds hydrogen to a compound* We
must now add that it is also one that removes oxygen from a
compound. Hydrogen dioxide has the remarkable property of being
not only a good oxidizing agent, as we have jiist seen, but also of
being a good reducing agent in a number of cases. When hydrogen
dioxide is brought in contact with metal oxides rich in oxygen, both
the dioxide and the metal oxide give up their excess of oxygen.
When hydrogen dioxide is brought in contact with manganese
dioxide or lead dioxide in the presence of an acid, the following
reactions take place: —

HjOj + MnOj = H,0 + MnO + 0^ ;
HjOa + PbOi = H,0 + PbO + 0,.

It should be observed that exactly one-half of the oxygen set
free comes from the hydrogen dioxide and one-half from the oxide
of the metah Each substance being rich in oxygen gives up a part
of its oxygen and is reduced ; the metal forming the salt of the acid

WATER AND HYDROGEN DIOXIDE 67

Catal3rtic Becompotition of Hydrogen Dioxide. — It was stated
above that when fairly concentrated hydrogen dioxide is brought
in contact with certain metals like platinum, it is decomposed with
explosive violence. Metallic platinum, especially in the finely
divided condition, can also decompose dilute hydrogen dioxide.
This can be illustrated very readily by simply bringing a piece
of platinum sponge, or some platinum black obtained by depositing
platinum electrolytically, in contact with hydrogen dioxide. The
platinum will become covered at once with a layer of gas, which
is oxygen resulting from the decomposition of the dioxide.

Many other metals and certain minerals, such as pyrolusite,
effect the same decomposition by simple contact with hydrogen
dioxide.

One other characteristic of this reaction should be observed. A
very small amount of the solid substance can decompose a large
amount of the dioxide, and further, the platinum or other solid is
unchanged by the reaction — it does not enter into the reaction.

We have already become acquainted with a similar reaction in
the combination of oxygen and hydrogen, as effected by contact with
metallic platinum. Such reactions were termed catalytic. We
have here another example of a catalytic reaction. The platinum
does not enter into the reaction, and a small amount decomposes a
large amount of the dioxide. These are the conditions which must
be fulfilled in order that a reaction may be termed catalytic — in
order that we may have catalysis.

Balations of Water and Hydrogen Dioxide. — We have seen that
as far as composition is concerned hydrogen dioxide is simply oxi-
dized water. We have also seen that the properties of the two
substances are very different. Water is a very stable chemical com-
pound, undergoing decomposition only with the greatest diflBculty.
To decompose pure water an electric current of high voltage must be
used, or it must be subjected to an enormous temperature. Hydrogen
dioxide, on the other hand, is a very unstable substance, undergoing
decomposition under the slightest provocation, or even spontane-
ously. The presence of the extra oxygen atom in water has thus
apparently changed it from one of the most stable to a very unstable
substance.

This is a mere statement of the facts observed, but the reasoning
mind is never content to stop here. Why does the presence of one
more oxygen atom in water give it such different properties ? This
is the question which every thinking person must ask. It can never
be answered by simply studying the composition of the two sub-

68 PRINCIPLES OF INORGANIC CHEMISTRY

stances — the material side of the problem. We must go deeper
into the problem and see what are the energy relations which obtain
in the different substances.

When we were studying oxygen and ozone we saw that the latter
differed from the former essentially only in the amount of intrinsic
energy present in its molecule. In the case of water and hydrogen
dioxide we have already discovered a difference in composition. Is
there any marked difference in the amounts of energy present in the
two molecules ? We can answer this question by converting hydro-
gen dioxide into water and measuring the amount of heat liberated.
Suffice it to say here that a large amount of heat is set free when
hydrogen dioxide passes over into water. This shows that hydrogen
dioxide contains more intrinsic energy in the molecule than water,
and this is the real cause of the marked difference in properties
between the two substances. Hydrogen dioxide containing the
larger amount of intrinsic energy is the less stable substance, and
this is strictly analogous to what we observed in the comparison of
oxygen and ozone. Ozone, containing the larger amount of intrin-
sic energy in its molecule, is far more unstable than oxygen.

We shall learn that this is a general relation. The more intrinsic
energy there is present in a substance, other things being equal, the
less its stability. This may be accounted for on the basis of intrinsic
energy tending to pass over into heat energy. In order that this
may occur, chemical transformation must take place. Indeed, this
probably lies very close to the foundation of all chemical reaction.

CHAPTER VI

DETSRICINATION OF RELATIVE ATOHIC 'WISLOBTB

Combining Vnmbers and Atomio Weights. — We saw in the second
chapter that the atomic theory was proposed to account for certain
well-established laws of chemical combination — the laws of definite
and multiple proportions and of combining weights. If atoms are
the ultimate units of matter, these must have definite weights, and
it is obviously of great importance for chemistry to determine the
relative weights of the atoms of different substances.

If the same number of atoms of any two substances combine,
the combining numbers or relative weights of the substances
that combine represent the relative weights of the atoms which
enter into combination. This apparently furnishes a means of
determining relative atomic weights. It is only necessary to deter-
mine the relative weights of substances which combine — the com-
bining numbers — in order to ascertain the relative weights of the
atoms of these substances. This would be true if a given number
of atoms of one substance always combined with an equal number
of atoms of another. But we know that this is not the case, since
it often happens that two elementary substances combine in several
proportions. To determine the relative atomic weights of the ele-
ments, we must, therefore, know the combining numbers of the
elements, and also the number of atoms of the different elements
which combine with one another. We will take up first the method
of determining the combining numbers of the elements.

Chemical Methods of determining Combining Numbers. — The
simplest method would be to take some element as our standard, and
call its combining number one. Then allow all of the other elements
to combine with this one, and determine the weights of the different
elements which combined with unit weight of our standard element.
Since hydrogen has the smallest combining number, it would natu-
rally be chosen as the unit. The problem then would be to determine,
say, the number of grams of the different elements which combine
with one gram of hydrogen, and these figures would represent the
combining weights of the elements in terms of hydrogen as unity.
Since it is true that comparatively few of the elements combine

70

PRINCIPLES OF INORGANIC CHEMISTRY

directly with hydrogen, the direct comparisou with hydrogen canDot
be made in many eases.

A large niiml>er of the elements, however, combine directly with
oxygen. We can determine the ratio between the combining miiiibera
of these elements and oxygen, and then the ratio between the com-
bining number of oxygen and that of hydrogen, and thus calculate
the combining numbers of the elemeuts in terms of our unit hydrogen*

We might thus work out a table of the combiuing numbers of all
of the elements in terms of hydrogen aia unity. This part of the
problem is, however, not as simple as would be indicated from the
above. Many of the elements combine in more than one proportion.
Take the case of hydrogen and carbon. The combining number of
carbon iu terms of hydrogen as unity would be 3, if determined by
the analysis of marsh gas. From the anal3^sis of ethylene we would
conclude that it was 6, while from the analysis of acetylene it would
appear to be 12, A similar complexity would result in the case of
carbon and oxygen. If we take oxygen as 16 in terms of hydrogen
1, the combining number of carbon, as determined from carbon
monoxide, would be 12, while as determined from carbon dioxide it
would be 6. We would thus obtain different combining numbers
for the same element, depending upon which of its compounda we
selected.

It is perfectly clear that neitlier^the chemical analysis of the
compound, nor its synthesis from the elements, throws any light on
the problem as to the number of atoms of one substance combined
with one atom of the other. Berzelius attempted to solve this pai*t
of the problem of atomic weights by means of ceilain dogmatic rules,
which have only this value, that they brought out a large amount
of experimental work which resulted in new and improved metho<la
of analysis. Chemical methods alone can lea^d only to the combin-
ing weights or numbers of the eleiuenta, and, as already stated, in
many cases more than one combining weight for an element would
be obtained. Other methods must be employed in order to deter-
mine the number of atoms of the one element which have combined
with one atom of the other To these we shall now turn.

Holecular Weights determined from the DenBitiet of Gases. — Gay-
Lussac shiDwed in 180S that the densities of gases are prupoit tonal
to their combining weigh ts^^ or to simple rational multiples of them.
If two gases react chemically, the volumes which react are either
equal, or bear a simple rational relation to one another. And,
further, if the product formed is a gas^ its volume bears a simple
rational relation to the volumes of the gases from which it waa

DETERMINATION OF RELATIVE ATOMIC WEIGHTS 71

formed. Thus, one volume of hydrogen combines with one volume
of chlorine, and forms two volumes of hydrochloric acid gas. One
volume of oxygen combines with two volumes of hydrogen, forming
two volumes of water-vapor. One volume of nitrogen combines with
three volumes of hydrogen, forming two volumes of ammonia.

From the laws of definite and multiple proportions, the law of
combining numbers, and the atomic theory which was proposed to
account for these, we see that every chemical reaction takes place
between a definite number of atoms, and the number is usually
small. Therefore, the discovery of Gay-Lussac leads to the con-
clusion that —

The number of atoms contained in a given volume of any gas must
bear a simple, rational relation to the number of atoms contained in an
equal volume {at the same temperature and pressure) of any other gas.

We have thus far, however, no means of determining the numeri-
cal value of this relation, and, therefore, cannot use the discovery of
Gay-Lussac alone to determine relative atomic weights.

Avogadro's Hypothesis. — Avogadro, in 1811, taking into account
all of the facts known, advanced the hypothesis that —

In equal volumes of all gases, at the same temperature and pressure,
there is an equal number of ultimate parts or molecules.

Avogadro extended his hypothesis to all gases, including even
the elementary gases, and regarded the molecules of these substances
as made up of atoms of the same kind, which had united with one
another. This was a necessary consequence of his hypothesis.
One volume of hydrogen gas combines with one volume of chlorine
gas, and forms two volumes of hydrochloric acid gas. If there are
the same number of molecules in equal volumes of all gases, there
would be twice as many in the two volumes of hydrochloric acid as
in the one volume of hydrogen, or the one volume of chlorine. Since
each molecule of hydrochloric acid must contain at least one atom
of hydrogen and one atom of chlorine, the molecule of hydrogen
and of chlorine must be made up of at least two atoms. Ampere,
in 1814, advanced essentially the same hypothesis as had been pro-
posed three years before by Avogadro. The hypothesis of Avogadro
has been confirmed by such an abundance of subsequent work, in
80 many directions, that it is now placed among the well-established
laws of nature. It points out distinctly the difference between
atoms and molecules, and rationally explains why different gases
should obey the same law of volume and of pressure, and have the

72

FHrXCIPLES OF INDUGANIC CHEMISTRY

same temperature coefficient of expansion. It has been tested from
both the physical and mathematical standpoints, and now lies at the
basis of HHieh of our knowledge of gases,

Avogadros Hypothesis and Uroleciilar Weights, — Given the
hyiK>thesis of Avogadro, the determination of the relative molecular
weights of gases is very sim]>le. If there is an equal number of
molecules contained in equal volumes of the different gases, the
relative weights of equal volumes of these ga&es give at once the
relative weights of the molecules contained in them. It is only
necessary to choose some substance as our standard, and express the
molecular weights of other substances in terms of this standard.
We would naturally select as the unit that substance which has
the smallest density, and this is hydrogen. From what has been
saidj liowever, in reference to the union of hydrogen and chlorine,
forming hydrochloric acid, it is certain that the molecule of hydrogen
contains at least two atoms. We will, therefore, call the molecular
weight of hydrogen two, and calculate the molecular weights of
other elements in terras of this sUmdard, The densities of sub-
stances are usually determined in terms of air as the unit. It is a
simple matter to recalculate these in teniis of hydrogen as two*
The density of hydrogen in terms ot air as the unit is 0.0606.
We must midtiply this by 28.73 to obtain our new unit 2
(2 -t- O.Or>96 — 28. T3), Similarly, for other substances whose densities
are known with reference to air; and these densities must be multi-
plied by the constant 28,73 to tmnsform them into densities in terms
of hydrogen = 2, These latter values are the relative molecular
weights of t!ie substances in the form of gas, referred to the molec-
ular w^mght of hydrogen as two, A few results are given in the
following table, showing in coluian 1 the derjsities in terms of air as
the unit; in column II the densities or relative molecular weights in
terms of hydrogen = 2, The results in column II are obtained by
multiplying the results in column I by 28,73,

Hydrogen, 0° C,
Oxyg«n, 0* C. ,

Kltrogen, 0° C. .
Sulphur, 1400" r.
Chlorine, 200^ C,
BTOtnhKs mf i\
Mercury, HOO^ C,
IcKiiiK', Uli}^ a ,

I

0,0000

1J0563

0.0713

2.17
2.45
5.54

5,et

8J3

x28Ja

2

31.79

27.flO

62.M

70.38

150.16

105,05

2rj0.52

DETERMIXATION OF RELATIVE ATOMIC WEIGHTS 73

The molecular weights of compounds can be determined in exactly
the same manner from the densities of their vapors. If these have
been detennined on the basis of air as unity, we must multiply by
28.73 to obtain the molecular weight referred to hydrogen as two.
The molecular weights of compounds, thus obtained, must bear a
rational relation to the combining weights of the elements which
enter into the compound. The molecular weights as obtained from
vapor-densities can, therefore, be corrected by the most careful
analytical or synthetical determination of the combining weights
of the elements which enter into the compounds.

Atomic Weights from Molecular Weights. — If we knew the num-
ber of atoms contained in the molecule of elements in the gaseous
state, the problem of relative atomic weights would be solved at once
by dividing the molecular weight of the gas by the number of atoms
in the molecule. The problem is, however, not as simple as this,
since we do not know a priori the number of atoms in the molecules
of elements. Other lines of thought have enabled us to solve this,
the second part of our problem.

The definition of an atom as an indivisible particle of matter
shows that fractions of atoms cannot exist. No molecule can con-
tain a fraction of any atom. The quantity of any substance which
enters into a molecule must be at least one atom. It may be more
than one, but it cannot be less. This is the key to the problem.
Suppose we wish to determine the number of hydrogen atoms in a
molecule of hydrogen. We must examine compounds into which
hydrogen enters, and find out what is the smallest quantity of
hydrogen which enters into the molecule of the compound. Let
us take hydrochloric acid, whose molecular weight is 36.45. This
is shown by analysis to be composed of 1 part of hydrogen and 35.45
parts of chlorine. This 1 part of hydrogen is at least one atom ;
it may be more, but it cannot be less. By examining a large num-
ber of compounds into which hydrogen enters, it has been found
that hydrogen never enters into a molecule of any substance in a
smaller quantity than in hydrochloric acid. This is, therefore, for
us the atom of hydrogen, but it may in reality be composed of a
great number of smaller parts. The hydrogen which enters into the
molecule of hydrochloric acid is just half the quantity which form?
the molecule of hydrogen gas, since one volume of hydrogen com-
bining with one volume of chlorine yields two volumes of hydro-
chloric acid gas. The molecule of hydrogen, therefore, contains at
least two atoms, and since there is no experimental reason for
assuming that it contains more than two, we say that the molecule

74

PRINCIPLES OF INORGANIC CHEMISTRY

of hydrogen is made up by the union of two hydrogen atoms.
Knowing the number of atoms in the molecule, the atomic weight
follows at once from the molecular weight determined by vapor-
density, and corrected by the most refined methods of chemical
analysis.

By methods similar to the above the molecules of many elements
have been shown to be composed of two atoms. But this by no
means applies to all elementary substances. The molecules of some
'.elementary substances contain more than two atoms, and in a very
ifew cases the molecule and atom seem to be identical. And, further,
rthe number of atoms contained in the molecule has been shown to
vary in some cases with change in conditions, especially with change
of temperature. But by studying a large number of compounds of
an element, and ascertaining what is the smallest quantity of the
element which ever ent-ers into a compound, we can determine the
number of atoms contained in a molecule of the element itself.
Knowing the number of atoms in the molecule of the element, and
the weight of the molecule, we can determine relative atomic
weights. The relations between the molecular weights of a few of
the elements and their atomic weights are given in the following
table: —

Elemknts

Hydrogen
Nitrogen .
Oxygen .
Phosphorus

Sulphur .

Chlorine .
Arsenic .
Selenium
Bromine .
Cadmium
Tellurium
Iodine
Mercury .

Atoxic Wkiouto

MOLKCVLAft WkIQBTB

1

2

14.01

28.02

16.88

81.76

30.96

123.84

81.98

r

I

03.96 above 800* C.

191.88 at 600*^ C.

86.18

70.36

74.9

299.6

78.9

167.8

79.84

158.68

111.7

111.7

126.8

252.6

126.89

251.78 under 600° C.

199.8

199.8

This table brings out a number of facts to which reference has
already been made. The molecular weight of a number of the
elements is twice as great as the atomic weight. In some cases, as
with sulphur, the molecular weight is twice the atomic weight at
a given temperature, and then varies with the temperature. In the

DETEIIMINATIOX OF RELATIVE ATOMIC WEKIHTS 76

eases of cadmium and mercury the molecular weights are apparently
identical with the atomic weights. This matter will be taken up
later in other connectiona.

It frequently hapjiens that an element boils at such a high
temperature that we cannot determine accurately its vapor-density.
In such cases volatile compounds of the element are used, and
their molecular weights determined. These com pounds are then
analyzed^ and the one containing tlie smallest quantity of the
given element in its molecule is said to contain one atom of the
element The real atom of the element may be a fraction of this
quantity, but this is for all chemical or physical chemical purposes
the atom^ and its relative weight is the atomic weight of the
element in question.

Atomic Weight! from Specific Heats. — Dulong and Petit in 1819
showed that a verj' simple relation exists between the specific heats
of elements in the solid state and their atomic weights- They found
that the specific heats varied inversely as the atomic weights^ and,
consequently, that the product of the specific heats ami atomic
weights of the elements is a constant. This will be seen from the
following data: —

Sricirrc HiEAt

Atomic Wmanr

PsODtJOT

Lithium

0.041

7.01

OJ

ficdium.

» * *

O.90S

22.00

6.7

« t *

0.2^jO

2S.94

6.0

Fotaidum

0.166

30.03

6.6

Calcium

0.170

.mm

6,0

Iron

0.112

55. W)

6.S

Cobalt .

0.107

68,00

6.3

Kickel .

0.108

58,60

6.4

Zinc

0.0032

04.90

a I

Prom these and similar facts Dnlong and Petit announced their
law; —

The atoms of all elements ham the same mpmit^for heat energy.

After the discovery of this law it was a comparatively simple
matter to determine the atomic weights of solid elements from their
specific heats. If specific heat multiplied by atomic weight is a
constant, the atomic weight is equal to the constant divided by the
specific heat. The numerical value of the constant, taken as tlie
average for a number of elements, is about 6.25.

74

PRINCIPLES OF INORGANIC CHEMISTRY

of liydrogen is made up by the union of two hydrogen atoms.
Knowing the number of atoms in the molecule, the atomic weight
follows at once from the molecular weight determined by vapor-
density, and corrected by the most refined methods of chemical
analysis.

By methods similar to the above the molecules of many elements
have been shown to be composed of two atoms. But this by no
ipeans applies to all elementary substances. The molecules of some
«elementiiry substivnces contain more than two atoms, and in a very
ifew cases the nioUn^ule and atom seem to be identical. And, further,
ttbe luimbor of atoms contained in the molecule has been shown to
vary in some oases with change in conditions, especially with change
of teniperaturo. But by studying a large number of compounds of
an element, and ascertaining what is the smallest quantity of the
element which ever enters into a compound, we can determine the
number of atoms contiiined in a molecule of the element itself.
Knowing the numWr of atoms in the molecule of the element, and
the weight of t\u) molecule, we can determine relative atomic
weights. Tlie relati(ms l)otween the molecular weights of a few of
the elements and their atomic weights are given in the following
table : —

Kl-KMIfNTH

Atomio Wkigiits

MoLRCVLAR Wrights

Hydrogen

Nitrogen

Oiygeu

Vboapbonw

Sulpbur

CWorlue

AwduU?

S^Wuiutti

1|IV4UUM»

C^iuuuw

Xtawxuuu

V^iittc

1

14.01
16.88
30.96

81.98

86.18

74.9

78.9

79.34

111.7

126.3

126.89

199.8

2

28.02
31.76
123.84
r 63.96 above 800° C.
I 101.88 at 600° C.
70.36
299.6
167.8
168.68
111.7
262.6

261.78 under 600° C.
199.8

ru.^ -ftttiK^ Vfi'swp* vH^^ * uumln^r of facts to which reference has
<«ijMJk \fcAie. Tb<» TOoltK^ilar weight of a number of the

*i»^k«^ V ^vsx le^ ^w^ A* t>^^ »^^'"^'® weight. In some cases, as
^i*.w. -^ ^i^^fciM ^•^VK'ht 18 twice the atomic weight at
w^.^MMMMtiztt. w*. Aim ^vi«« ^'ith the temperature.

Tisa

In the

DETERMINATION OF RELATIVE ATOMIC WEIGHTS 75

cases of cadmium and mercury the molecular weights are apparently
identical with the atomic weights. This matter will be taken up
later in other connections.

Jt frequently happens that an element boils at such a high
temperature that we cannot determine accurately its vapor-density.
In such cases volatile compounds of the element are used, and
their molecular weights determined. These compounds are then
analyzed, and the one containing the smallest quantity of the
given element in its molecule is said to contain one atom of the
element The real atom of the element may be a fraction of this
quantity, but this is for all chemical or physical chemical purposes
the atom, and its relative weight is the atomic weight of the
element in question.

Atomic Weights from 3pecific Heats. — Dulong and Petit in 1819
showed that a very simple relation exists between the specific heats
of elements in the solid state and their atomic weights. They found
that the specific heats varied inversely as the atomic weights, and,
consequently, that the product of the specific heats and atomic
weights of the elements is a constant. This will be seen from the
following data : —

SPKCiric Hkat

Atomic Wkioiit

Peoduot

Lithium

0.041

7.01

Q.Q

Sodium

0.203

22.00

6.7

Magnesium

0.250

23.04

6.0

Potassium

o.iee

30.03

6.6

Calcium

0.170

30.01

6.8

Iron

0.112

66.00

6.3

Cobalt

0.107

68.60

6.3

Nickel

0.108

68.60

6.4

Zinc

0.0032

64.00

6.1

From these and similar facts Dulong and Petit announced their
law: —

The atoms of all elements have the same capacity for heat energy.

After the discovery of this law it was a comparatively simple
matter to determine the atomic weights of solid elements from their
specific heats. If specific heat multiplied by atomic weight is a
constant, the atomic weight is equal to the constant divided by the
specific heat. The numerical value of the constant, taken as the
average for a number of elements^ is about 6.25.

76

PRINCIPLES OF INORGANIC CHEMISTRY

Exceptions to the law of Dulong and Petit were early recognized,
Weber determined the specific heats of the elements carbon, boroD^
and silicon, at tem]_>eraturea between 0° and 100° C, and obtained
much smaller values than would be expected from the law of Dulong
and Petit, nsing the atomic weights of these elements as determined
from Avogadro*s law. He found, however^ that the Bi>ecific heats of
these elements varied widely with change in temperature, and that
above a certain temperature the specific heats became constant At
these elevated temj>eraturea, where the specific heats became con-
stant^ they conformed to the law of Dulong and Petit, These
constant specific beats were obtained only at comparatively high
temperatures ; for silicon at about 200^ C, for the different modifica-
tions of carbon at about GOO" C, for boron at about 500° C. The
different m'odifications of carbon had different specific heats at low
temperatures, but at elevated temperatures tins difference also was
found to vanish, the different varieties of carbon at red beat show-
ing tlie same specific heats. Similar observations were made on
gluciuum by Xilson and Pettersson.

The law of Dulong and Petit is, in general, only approximately
true, and holds only within certain limits of temperature.

The relation between the specific heats of compounds and
the specific heats of their constituents was next investigated*
Neumann showed that equivalent quantities of analogous com-
pounds have the same capacity for heat, and Regnault^ Xopp^ and
others pointed out the following relation between the specific lieata
of compounds and the specific heats of their constituents* 77i€
cajHicif^ of the atoms for heat energf^ i» not appreciably changed when
they unite and fonu mmponnds. In a word, the capacity of the
molecule for heat is the sum of the capacities of the atoms in the
molecule.

The recognition of this relation makes it possible to greatly
extend the method of determining aton^ic weights by specific heats.
Many of the elements are solids only at temperatures which are too
low to be dealt with by the methods of measuring specific heats.
Put these elements form solid compounds with other elements whose
specific heats and atomic weights can be determined. Let ua take
an example.

Chlorine is an element whose specific lieat in the solid state
would be very difficult to determine. Chlorine, however, forms a solid
compound witli tlie element lead. The specific heat of lead chloride
has been found by Regnault to be 0,CX)C4; 206,4 parts of lead yield
277,1 parts of lead chloride, Multiplying this number by the spe-

DETERMINATION OF RELATIVE ATOMIC WEIGHTS 77

cific heat of lead chloride, we obtain the molecular heat. 277.1 x
0.0664 = 18.4. Subtracting the atomic heat of lead, 6.5, we have 11.9
as the atomic heat, corresponding to 70.7 parts of chlorine. Since
the atomic heat of the elements is about 6, we have in 70.7 twice
the atomic weight of chlorine, or the atomic weight of chlorine =
35.35. This agrees very closely with the atomic weight of chlorine
determined by the vapor-density method, based upon the law of
Avogadro, and by analysis.

The above example serves to illustrate the way in which the spe-
cific heats of compounds are used to determine atomic weights. The
method has been widely applied, and it may be said in general, that
the atomic weights determined from the law of Dulong and Petit
agree with those obtained from the law of Avogadro, although some
discrepancies exist.

Isomorphiam an Aid in determining Atomic Weights. — It was
recognized even in the eighteenth century that substances of different
composition often have the same, or very nearly the same crystal form.
This was at first explained by assuming that certain substances have
the power of forcing other substances to take their own crystal form.
Mitscherlich interpreted this fact quite differently. He studied
the salts of arsenic and phosphoric acids, and found that those which
contained an equal number of atoms in the molecule had the same,
or very similar crystal forms. Mitscherlich concluded at first that
it was only the number and not the nature of the atom which condi-
tioned the crystal form. Later, he recognized that the way in which
the atoms were united in the compound was an important factor in
determining its crystal form, and then arrived at the generalization
that, ^^ An equal number of cUoms combined in the same tvay produce
the same crystal form^ and that the same crystal form is independent of
the chemical nature of the alom^, but depends only on their number and
position,^*

If this relation was true, it would throw much light on the num-
ber of atoms in a compound, and, therefore, be of service in deter-
mining atomic weights. Given two isomorphous substances such as
BaClj 2 HjO and BaBrj 2 HjO; from the law of Mitscherlich their
molecules must contain the same number of atoms. If we know
the atomic weights of all the elements in the former compound,
we can find the atomic weight of the bromine in the latter sub-
stance.

This relation pointed out by Mitscherlich was accepted at once
by Berzelius, who made it the basis of atomic weight determina-
tions. The law, however, did not long remain without exceptions.

..\* Ml. I.N ^'t- =\nUi;.\NlC niEMISTUY

Ml lo . ».uii|MMiinis RaMii.O^ y;uS4?^, and

....... 1.1 • ;i«\\ t\ nu* lit iy mi mill . I '.'in'v .Liferent

I- ...*».*.*. \ii ALti'iupc viis ■Jiiidt* ro <.)VJin.'onie

. . • M-.M' I i»mp*»uiius "he rormiiliLi. BaMn«<\

\ .. . .. i. ^^ Aiit* ^o -irfiii:';" ic T-iruiai.-i* vi::ii all

i.i .» 'I- u'u;:iLi'iie'i. i2.«i a z-ziber of

• , I. ■> \ :!'.:uv ri,: :^i* sL-iie '--xber of

^t ■>». ^. • T ^ •?i«>T mi' ' iz. a;: yr.xtniarioa
\ »..:.iiv j.M, .♦ s. ip'ra-v. .:: i.f :<::::: in in ;^
.. . •» » .•• •» ■ t.» *■' 'i*.

. • .*-;.». ■ ••!• -'it- ;••« :z :*: as a zieans

-Mo- >.*;.■■■»- ila: e:'j.alitv

* . - ^ '.i -•■^•■ssjLTv •.- :r.:er :hac

\ •%,.■.•■ :l'r :c7r.: i>:Eior-

. »•■* • « ' .*! ■'•'-*>>■■..::■. -5. The

V.-...*-. ■■.' X-: -.* -**r= -. r-'V'.eci of

•* » < >*:-^ -v^". .-s-vc-.ally in the

y^ V, . . Xv%'Uv. '^ 'itwiMfcJUui^ JLttfouc Wei^liti. — The

»,%;... I ..; o -v'.i::ve aromic

.» . ..*.i,.^ :4 u-' f ; i: \c .uv:;raoy. Of

..»..• *', •»..'. Ill- •» 1,^ : 1*0 cvHistitvieiits

u ....... 4. K.i: r\v l**Ji-t\i» the other

». X ,,A'; iii.'^'s ino:hiH.U ami the

».-.'.i\. I'.ia u'.vu isomorphism,

, v,A. ,jA-.i ■ Iio \ Itv'Jiiua't methihls. l^y

\ . .^. .:>...:», \w» Jcteniiim* with the

, . »?, !i., «i'i ;!as of elements, ami

.1 A s *• i\\ ;*Io \\ he I her we are dealing

I , xw Ml J. I kh^K»se some element

. .. .k '\ '.I'vo I he lij^hlest element,

lit . ■».*. -sx-J* ih'iie, aiul all atomic

\ \. '.•.!':. imtoiiiuiately true, as

.» ■ ^ ^ ^ ^^,, K'l- v»»inl»ine iliivetly with

\ ■ V. ..*k*^v *\»juiK»uiuU whieh can he

I '»

\v.. v„.V?ie vMlh a hiri-e number of
'^ ^^' ' ' ^ ^^, ^;^^;»;vvvmi»o\imls with whieh

DETERMINATION OF RELATIVE ATOMIC WEIGHTS 79

we are acquainted. It therefore seemed best to compare the atomic
weights of the elements directly with the atomic weight of oxygen,
and then compare oxygen with hydrogen, with which it forms the
very stable compound, water. It should be stated, however, that
this method is by no means free from objections, and many prefer
retaining hydrogen as the unit The atomic weight of oxygen, in
terms of hydrogen as the unit, was supposed for a long time to be
the whole number 16. If this was true, it would obviously make no
difference whether we called hydrogen 1 or oxygen 16, and then
compare all other atomic weights with these standards. It has
recently been shown beyond question that when hydrogen is 1, oxy-
gen is not 16, but considerably less (15.88). We must, therefore,
choose between these two substances as the basis of the system of
atomic weights. The majority of investigators at present seem
inclined to select oxygen as the standard, taking its atomic weight
as 16, and referring the atomic weights of all the other elements to
this basis.

The most direct method of determining the combining weight of
an element, in terms of oxygen as our standard, would be to deter-
mine the weight of the element which would combine with a known
weight of oxygen. The combining weight of the element would
then be calculated by the simple proportion, —

Wt. oxygen : wt. element = at. wt. oxygen : combining wt. element.

We should then have to determine, by some of the methods already
referred to, how many atoms of the element in question combined
with one atom of oxygen.

W^hile it is true that oxygen combines directly with many of the
elements, forming stable compounds, it is by no means true that it
forms such compounds with all of the elements. And fui-ther, some
of the elements forqj compounds with oxygen which are gaseous or
liquid at ordinary temperatures, and for these or other reasons are
not adapted to atomic weight determinations. In such cases the
atomic weight of the element must be compared with that of some
element other than oxygen, which in turn has been compared with
oxygen. Thus, the atomic weights of the halogens have been deter-
mined in terms of the atomic weight of silver, and the latter then
determined in terras of oxygen. Even more complex cases may
arise, where it is necessary to compare the atomic weight of an ele-
ment with the sum of the atomic weights of two or more elements,
each of which has been determined in terms of oxygen.

It is evident that the more direct the comparison of the atomic

80

PRINCIPLES OF INORGANIC CHEMISTRY

weight of the element with that of oxjgeii, the better; since the
accumitlatioji of experimental errors, resulting from indirect com-
parisons, is avoided.

Some of the moat refined experimental work which has ever been
done has had to do with the problem of relative atomic weights.
It is obviously necessary that these constants should be determined
with the very greatest degree of accuracy, since all chemical analysis
and much of the most refined M^ork in physical chemistry and in
physica depends upon them. In this connection we should men-
tiouj especially among the earlier workj that of Stas and Marignac,
and among the more recent investigations those of Morley and
Richards,

The work of Stas had to do more especially witli the relations
between silver nnd the halogens, but ineludedj also, a large number
of other elementSj especially lithiumj sodium^ potassium j sulphuTj
lead, and nitrogen. The work of Stas, as a whole, has become a
model for refinement and accuracy, and is simpjly wonderful, when
we consider the comparatively crude apparatus with which it was
carried out.

Marignae has done an enormous amount of woi-k on the problem
of atomic weights. He has determined the atomic weights not only
of chlorine, bromine, and iodine, but of carbon and nitrogen, calcium,
barium, magnesium, zinc, manganese, nickel, cobalt, lead, bismuth,
and many of the rarer elements.

The comparatively recent work of Morley on the ratio between
the atomic weights of oxygen and hydrogen is one of the finest
pieces of scientific work in modern times. He has established this
ratio by different methods, with an unusual concordance in the re-
sults, to be 1 : 15.879,

The work of T. W, Richards on the atomic weights of a large
num\>er of the metals should receive special mention. He has im^
proved old methods, devised new ones, and applied them with a skill
which is rare. His determinations are to be ranked among the very
best wiiich have ever been made.

Table of Atomic Weights- — The most probable atomic weights
of the elements, based upon tlie best determinations, are given in
the following table. In preparing this table the tables of Clarkei
of Richards, and of the committee of the German Chemical Society
have all been carefully considered j also the original determinations
themselves, wherever there were appreciable differences between the
values chosen by the different authorities* The basis uf this table
is oxygen « IQ,

DETERMINATION OF RELATIVE ATOMIC WEIGHTS 81

Elbmsxt

AniMic WkiuuT

EtUfKKT

Atomic Wbiuht

Alaminiiun •

Antimony . .

Argon • . .

Arsenic . •

Barium . .

BLsmath . .

Boron . . .

Bromine . .

Cadmium . .

Csesium . .

Calcium . .

Carbon . .

Cerium . .

Chlorine . .

Chromium .

Cobalt . . .

Columbium .

Copper . .

Erbium . .

Fluorine . .

Gadolinium .

Gallium . .

Germanium .

Gluclnum . .

Gold . . .

Helium . .

Hydrogen . .

Indium . .

Iodine . . .

Iridium . .

Iron . . .

Krypton . .

Lanthanum .

Lead . . .

Lithium . .

Magnesium .

Manganese .

Mercury . .
Molybdenum

27.1
120.2

30.0

76.0
137.4
208.5

11.0

70.06
112.4
132.0

40.1

12.0
140.25

35.45

52.1

59.0

04.0

03.6
166.0

19.0
156.0

70.0

72.5

0.1

107.2

4.0

1.008

115.0

126.97

193.0

55.9

81.8
138.9
206.9
7.03

24.36

55.0
200 0

06.0

Neodymium

Neon . .

Nickel .

Nitrogen .

Osmium .

Oxygen .

Palladium

Phosphorus

Platinum

Potassium

Praseodymium

Radium .

Rhodium

Rubidium

Ruthenium

Samarium

Scandium

Selenium

Silicon .

Silver . .

Sodium .

Strontium

Sulphur .

Tantalum

Tellurium

Terbium

Thallium

Thorium .

Thulium .

Tin . .

Titanium

Tungsten

Uranium

Vanadium

Xenon

Ytterbium

Yttrium .

Zinc . .

Zirconium

143.6

20.0

58.7

14.04
101.0

16.0

ior>.5

31.0
194.8

39.15
140.5
225.0 (?)
103.0

85.5
101.7
150.3

44.1

70.2

28 4
107.93

23.05

87.6

32.06
183.0
127.6
160.0
204.1
232.5
171.0 (?)
119.0

48.1
184.0
238.5

51.2
128.0
173.0

89.0

05.4

90.6

CHAPTER Vir

DETERMINATION OF THE MOLECULAH WEIGHTS OF OASES
AND OF DISSOLVED SUBSTANCES

DENSITIES AKD MOLECULAR WEIGHTS OF GASES

Beniities and MolecElar Weiglita. — The det€rmiQatio]i of the rela^
tive deusitiea of gases coiisista in deter mining the relative weights
of equal volumes of gases at the same temperature and pressure*
Since equal volumes of gases under the same conditions contain an
equal numher of molecules, the densities stand in the same relation
as the molecular weights. Thus, by means of Avogadro's law we
can determine the molecular weights of substances in the gaseoua
state.

Some substance must be taken as the unit in determining the
densities in gases. Air has generally been selected as the unit, and
the weights of equal volumes of other gases, at the same temperature
and pressure, compared with that of air. Hydrogen has also been
used as the unit, and is to be preferred to air, since the composition
of the latter varies slightly from time to time and from place to place.
The density of air is 14.37 times the density of hydrogen, and since
the molecular weight of hydrogen is 2, we must multiply the density
referred to air as the unit by 28.73 to obtain the molecular weight
of the gaa. If we represent the molecular weight of the gaa by m,
and the density referred to air as the unit by d^

m = d K 28.73.

In this way the molecular weights of gases can be calculated from
their densities.

A number of methods, and a large number of modifications of
methods have been proposed for determining the densities of gases.
The more important will be briefly considered.

Method of Bumai. — The method of Dumas consists m deter-
mining the amount of substance which, in the form of vapor, at a
given temperature, just tills a flask whose volume is afterwards

m

MOLECULAR WEIGHTS OF GASES 83

determined. The flask is weighed full of air. Knowing the volume
of the flask, we know the weight of air contained in it ; therefore, we
know the weight of the empty flask. The weight of the flask being
known, and the weight of the flask plus the substance which just
filled it with vapor, we know the weight of the substance. By deter-
mining the weights of the vapors of different substances which fill
a flask of given volume, we have the relative densities of the
vapors.

The apparatus used is a balloon flask (Fig. 12) holding from 100
to250cc.

The flask is carefully dried and
weighed, using as a tare another
flask of very nearly the same size.
We are in this way made indepen-
dent of the conditions of tempera-
ture, moisture, etc, under which the
weighing is made.

A few grams of the substance
whose vapor-density is to be deter-
mined are introduced into the flask,
the neck drawn out to a capillary,
and the flask placed in a bath which °' '

is at least ten or fifteen degrees above the boiling-point of the
substance. The substance vaporizes, drives out the air, and when
the vapor of the substance ceases to escape, the capillary is fused
shut. The flask after cooling is weighed. The fine point is then
cut off under mercury and the flask filled with mercury. The flask
may then be weighed again, or the mercury poured out and measured,
giving the volume of the flask.

The method of Dumas is not as well adapted to higher tempera-
tures as other methods to be considered later. In the first place, it
is difficult to measure high temperatures accurately ; and, further,
the amount of substance contained in the bulb at high tempera-
tures is so small that relatively large errors result from this source.
Deville and Troost have used this method at fairly high tempera-
tures, employing porcelain balloons, but their results are not very
accurate. The method of Dumas cannot be used with even a fair
degree of accuracy above 600** to 700** C.

An attempt has been made to use the Dumas method under
diminished pressure. Habermann has so arranged the bulb that a
low pressure can be maintained constant, and the pressure read on a
manometer. Larger bulbs are required for work under diminished

84

PRINCIPLES OF INORGANIC CirEMlSTEY

preBsure^ and even then the quantity of substance is so small that
considerable errors are introduced.

A large number of modifications of the method of Dumas have
been proposed, but that of Bunsen should be especially mentioned.
He used three vessels of the same volume and weight One was
empty, one was filled with air at a given temperature and preasnrej
and the third was filled with the vai_>or at the same temperature and
pressure. If we represent by Wi the weight of the vessel filled with
the vapor, by W^ the weight of the vessel filled with air, and by Wg
the weight of the vessel in which there is a vacuum, tlie relative
density of the vapor and air is expressed thus : —

HI - Tn

After vessels of the same volume and weight have once been
prepared, this method of procedure is more convenient and far more
rapid than that originally described by Dumas.

The method of Dumas is used less to-day than it was formerly,
having been largely supplanted by better methods, espectally at
elevated temperatures. The apparatus used in this met^jod is, how-
ever, exceedingly simple, and even at present the Dumas method is
employed in certain cases where the pmsence of a foreign gas in the
vapor must be avoided.

The Method of Gay-LoBsac. — The method devised by Gay-Lussac
for determiuing the densities of vapors is based u^wn a principle
which is quite different from that which we have just considered,
lu the method of Dumas the vapor required to fill a given volume
was weighed. In the method of Gay-Lussac a weighed amount of
substance is converted into vapor, and the volume of the vapor
measured. The method as originally proposed by Gay-Lussac con-
sists in placing a known weight of liquid in a calibrated vessel over
mercury* The whole is then warmed until the liquid is converted
into vapor. The tempemture is noted, also the volume of the vapor.
The latter ia reduced to standard couditioiis, a correetion being in-
troduced for the tension of the mercury vapor* This method has
been so greatly improved that the oriffiual is no longer used,

Hofmann's Hodification of the day Lussae Method. —Tbe modifi-
cation of the Gay-Lussac' apparatus proposed by Hofmann, consists
in elongating the inner tube beyond the barometric height so that
a vacuum will exist in the top of the tube. The substance is intro-
duced into the tube over the merf^ury and volatilized under diminished
pressure. The apparatus is shown in Fig, 13.

MOLECULAR WEIGHTS OF GASES

85

q

|fe{=t"

The calibrated tube A rests in a mereaij reserroir R. and is more
than 76 cm. long. It is fastened into a vapor-jacket J, into which
yapor enters at a, and leaves at 6. ^

m is a bar of metal terminating
in an adjustable point, which is
brought down to the surface of the
mercury ; the cross-hairs attached to
the bar at h serving to read more
accuratdj the height of the mercury
in the tube A.

After the substance is converted
into vapor the volume of the vapor
is read and reduced to standard con-
ditions. Knowing the weight of the
substance and the volume of vapor,
the density of the vapor is calculated
at once. The advantage of the modi-
fication proposed by Hofmann is that
the substance is converted into vapor
at a temperature below its boiling-
point under atmospheric pressure.
Thus, the vapor-density of many sub-
stances which would decompose if
boiled under atmospheric pressure
can be determined. Indeed, Hofmann devised this method especially
for use with organic substances which would easily decompose.

The Oas4isplaoement Method of Yictor Meyer. — A method for
determining vapor-densities which has practically supplanted all
other methods, except in very special cases, was devised by Victor
Meyer in 1878. The method consists in volatilizing a small,
weighed portion of substance in a tube filled with air, and collect-
ing and measuring the volume of air which is displaced.

The apparatus used is seen in Fig. 14. The inner vessel A is
surrounded by a glass jacket J, in which is boiled some substance
which will heat ^ to a constant temperature, and at the same time
to the temperature desired. The tube A is closed above with a
stopper, and from the central tube a side tube runs over to, and
under a calibrated tube filled with water and clipping into a water
reservoir. The substance to be used is weighed in a weighing tul>e
which is closed loosely at the top, and introduced, when desired,
into the top of A. In carrying out a determination, a liquid which
has a higher boiling-point than the substance whose vapor-density is

Fig. 13.

PEIXCIPLES OP INORGANIC CHEMISTRY

to be determined is placed in the outer jacket. This liquid is boiled,
and a pftrt of the air in the inner yeseel ia driven out When no
more air escapes from the side-tube, the tube coutaining a weighed
amount of guLstance ia introduced into the top of A^ and rests on
^ the rod n When temperature equilibrium

has been perfectly established, the mouth of
the side-tube is platted under the measuring
tube in the water tank, the rod r drawn back,
L S - j and the small vesfiel containing a weighed

III r amount of the substance allowed to drop to

^^^vJL^^^ Z the bottom of A. The substance volatilizes,

' ^ drives out the loosely fitting cork from the

weighing tube, and then displaces air from
the tube A. The displaced air is received in
the measuring tube f, and its volume is equal
to the volume of vapor formed in the tul^ A
by the known weight of the substance intro-
duced. We know the amount of substance
ufledy also the volume of the air displaced^
which is equal to the volume of vaiwir formed ;
consequently, the density of the vapor of the
substance.

A very small amount of substance suffices
for determining vapor-density by this method,
and the method can be used at very high
tern peratu tea. At higher temperatures vessels
of glass cannot of course be employed, but
porcelain can be used. Berlin porcelain can
be employed up to 1600**, and other more
resistant forms of i>orcelain can be used up
to 1700*, or perhaps a little higher. Platinum
vessels can be used up to 1700°. There is no
material known which can be used above 1800°.
The great advantage of this method, in addition to the small
amount of auhstance required, is that the temperature of the experi-
ment does not need to be known. It is only necessary to keep the
temperature constant before and after the introduction of the sub-
stance. The gas-<lisplacement methofl is so far superior to all others
at high temperatures that it has practically" supplanted them alh

It is not necessary to fill the vessel A with air. This may be
replaced by an indifferent gas» in c^se the oxygen of the air would
act chemically upon the substance to be vaporized. Thus, if we

FtG. 11.

MOLECULAR WEIGHTS OF GASES 87

were determining the vapor-density of arsenic or sulphur, oxygen
must be excluded, and the vaporizing vessel could be filled with
nitrogen or hydrogen. If the vapor of magnesium was being studied,
nitrogen could not be used, since it would act chemically upon the
magnesium.

The gas-displacement method of Victor Meyer has also been used
under diminished pressure, and the vapor-densities of substances
determined considerably below their boiling-points. The advantage
of increased stability of the substance at the lower temperature has
already been mentioned.

Method of Bunsen. — Bunsen has devised a rough method of
determining the relative densities of gases. Gases under the same
pressure pass through a small opening with velocities which are in-
versely as the square roots of their densities. The method consists
in allowing equal volumes of different gases to pass through a very
fine hole in a platinum plate, which covers the top of the cylinder
containing the gas, and noting the time required. The cylinder is
immersed in mercury, which enters from below as the gas escapes at
the top. The method is not capable of any very great refinement,
and the results obtained by means of it are only close approximations.

Of the methods considered for determining the densities of
vajwrs, that of Meyer is by far the most generally applicable. The
method of Gay-Lussac and the modification proposed by Hofmann
are seldom used. The method of Dumas is used at present only in
special cases, to which reference will be made in detail a little later.

Sefolts of Yapor-density Measurements. — The vapor-densities of
elementary gases have shown many interesting and surprising rela-
tions between the number of atoms contained in the molecules of
these substances. The molecular weights of a number of elementary
gases, calculated from their densities, show that the molecule is
made up of two atoms. This applies to hydrogen, oxygen, nitrogen,
chlorine, bromine, and a number of others. The vapor-densities of
mercury, cadmium, and glucinum show that the molecule is mona-
tomic, or that the molecule and atom are identical. On the other
hand, the molecules of phosphorus, sulphur, etc., contain more than
two atoms, if the temperature to which they are heated is not too
high.

Abnormal Yapor-densities. Apparent Exceptions to the Law of
Avogadro. — The vapor-densities of the elementary substances men-
tioned above show that the molecules of some vapors contain a
number of atoms, the molecules of others two atoms, while in some
vapors at low temperatures^ and in others at higher temperatures,

88

PELNXIPLES OF INORGANIC CHEMISTRY

the molecule contains one atom, or the molecular weight is identical
with the atomic weight. In the case of no elementaty substancSj
howe^^er, was the molecular weight found from vapor-density less
than the atomic weight of the element, and in none of the com-
pounds thus far mentioned was the molecular weight less than the
sum of the atomic weights of the elements entering into the com-
pound. In a number of cases the molecular weights showed that
the molecule of the compound was the simplest possibloj but there
was nothing to indicate that the simplest molecule had in any case
broken dov^-^n into its constituents. We must now turn to another
class of phenomena. The molecular weights of subshmces like
ammonium chloride, phosphorus pentachloride, et^!.^ calcuhited from
their vapor-densities, were less than the sum of the atomic weights
of their constituent elements. Thus, the vapor-density of ammonium
chloride^ corresponding to the formula NH^Cl, must be 1,89, while
Bineau found the value 0.89* The vapor-density of phosphorus
pentachloride of the formula PClj must be 7.20. Neunjann found
by the method of J>umas at 182* the value 5.08. Tliis decreased
with rise in temperature up to 290^, where it became constant at 3.7*
Similar results were found by Cahours. A number of other ex-
amples similar to the above were known, but these suflBce to illus-
trate the point. The explanation of these abnormal results was not
furnished at once, and for a time the hypothesis of Avogadro was
rather at a discount because of their existence* The explanation,
however, has been furnished^ as we shall now see, and the law of
Avogadro thoroughly substantiateth

Explanation of the Abnormal Yapor-denBitiea. — After Deville had
shown in 18.j7 that many chemical com]Kiuuds are broken down or
dissociated by heat, it occurred to Cannizxaro, Kopp, and others,
that the abnormal vapor-densities of substances like ammonium
chloride, pho.sphorus pentachloride, etc., might Ije due to the disso-
eiation of these substances by heat If a substance like ammonium
chloride was disso^'iated, one molecule would yield one molecule of
ammonia and one of hydrochloric acid. One molecule of phosphorus
pentachloride would break down into one molecule of phosphorus
trichloride and one molecule of chlorine. If such a dissociation did
take plac^, it would account for the abnormally small vapor-dens i ties
found, since the substances in the form of vapor would occupy a
greater space than if there waa no dissociation. But this did not
prove that such a dissociation actually took place. How could this
point be tested ?

Take the case of ammonium chloride ^ if it is dissociated by heat

MOLECULAR WEIGHTS OF GASES

89

it would yield ammonia and hydrochloric acid in equivalent quanti-
ties. It would, however, be exceedingly difficult, if not impossible,
to detect either ammonia or hydrochloric acid when the two gases
were mixed in equivalent quantities. This problem was solved by
Pebal. He made use of the different rates at which these two gases
diffuse to separate them, in part, in case they were present in the
vapor of ammonium chloride. The apparatus which he used is seen
in Fig. 15. The ammonium chloride d rests on a plug of asbestos c,
near the top of the inner tube, which is
open above. A stream of hydrogen is
passed through a into the outer part of the
apparatus, and another stream through b
into the inner part of the apparatus. The
whole is heated above the boiling-point of
ammonium chloride. If the salt is decom-
posed when it volatilizes, the ammonia
being lighter than the hydrochloric acid
would diffuse more rapidly through the
plug of asbestos. The vapor in the inner
tube below the plug would therefore con-
tain an excess of ammonia. This vapor is
swept out by means of the stream of hydro-
gen gas, and made to pass over a piece of
moist, red litmus paper in the vessel B.

It was found that this was colored blue, proving the presence of an
excess of ammonia.

The vapor remaining in the inner tube above the wad of asbestos
must contain an excess of hydrochloric acid, since more ammonia
has passed through the asbestos than hydrocliloric acid. This is
swept out by means of the stream of hydrogen in the outer vessel,
and passed over a piece of blue litmus in the vessel A, This turned
red at once, showing the presence of free hydrochloric acid in this
gas. It would seem, then, that Pebal had demonstrated beyond
doubt that the vapor of ammonium chloride contains both free
ammonia and free hydrochloric acid, and, therefore, that this sub-
stance is dissociated by heat.

The objection was, however, raised to the experiment of Pebal,
that a foreign substance, asbestos, had l)een used in contact with the
vapor of ammonium chloride, and that this might have caused the
vapor to dissociate, or at least might have facilitated the breaking
down of the salt by heat. This objection, while apparently having
but little foundation, could not be ignored. To test this point Than

90

PRINCIPLES OF LNORGANIG CHEMISTRY

I\a. 16.

devised the foUowing apparatus (Fig. 16): The tube AB, in which
the ammouiuui chloiide is eontaiiied, is placed horizontally, and the
septum is made out of ammunium chloride. Nitrogen is p^sed
through the tube, the ammonium chloride d heated with a lamp, and

the vapors in the two sides
passed over colored litmus, as
in the experiment of Pebal.
The vapor in the side next to
the amtuoiiiuju chloride was
found to contain free hydro-
chloric acid, and free ammonia
was shown to be present in
the vapor which hatl diffused
through the plug of ammonium
chloride.
It is thus shown beyond question that the vapor of ammonium
chloride is broken down, in part, into ammonia and hydrochloric
acid, by heat alone.

The work of Wanklyn and Robinson has shown that jihosphorus
pentachloride is dissociated by lieat into the trichloride and chlorine.
The penta^'hioride was placed iu a short-necked glass flask, in which
it was to \ye converted into vaiKJr, Over the neck of this flask a
wider gk3a tube was phiced, so that tlie two were separatetl by an
atr-space. Air was passed iu through the upper tube and escaped
throuj,'h the space between tlie two glass tubes. If the vapor of the
pentachloridf was dissociated by heat into the trichloride and chlo-
rine, Uiese would diffuse with different velocities into the upper
portion of the vessel, since they have different vapor-densities.
They would then be swept out by the curmnt of air in different
quantities, the chlorine being in excess since it is the lighter^ and
wouldj therefore, diffuse more rapidly into the upper poiiion of the
vessel.

Free chlorine was proved to be present in the vapors which
escaped^ and analysis showed an excess of phosphorus trichloride
remaining in the flask. Therefore, the phasiihonis penUichloride
was brt>ken down, in part at least, bj heat itito its cousfttuents.
This conclusion was confirmed by the observation that as the vapor
of phosphorus pentachloride is heated higher nnd higher it liecomes
colored more deeply greenish-yellow, — the chai^acteristic folor of
chlorine itselt

The same explanation undoubtedly applies to other sul^tances
whose vapor-densities are abuormally small They are more or less

LAW OF MASS ACTION 91

broken down by heat into their constituents ; the amount of the dis-
sociation increasing with the temperature.

Diliociation of 7apon diminished by an Excesi of One of the
Products of Dissociation. — A discovery was made iu connection
with the study of dissociating vapors, which has proved to be of the
very highest importance. If there is present an excess of either of
the products of dissociation, the amount of the substance decom-
posed is lessened. Thus, ammonium chloride is less dissociated if
there is present an excess of either ammonia or hydrochloric acid.
Similarly, phosphorus pentachloride is much less decomposed at a
given temperature if there is present an excess of either phosphorus
trichloride or chlorine, as Wttrtz has shown. Indeed, the vapor of
phosphorus pentachloride is scarcely dissociated at all by heat in
the presence of an atmosphere of phosphorus trichloride, or of chlo-
rine. The vapor-density of phosphorus pentachloride in an atmos-
phere of the trichloride was found to be about 209, while the
calculated vapor-density is 208.

This is a perfectly general principle, illustrated by phosphorus
pentachloride and ammonium chloride. The dissociation of sub-
stances in general by heat is driven back by an excess of any one of
the products of dissociation.

This is the second example thus far met with of the effect of
mass on chemical activity. The importance of the action of mass
will be more clearly seen as the subject develops. We shall now
take up the law of mass action.

THE LAW OF MASS ACTION

The Work of Onldberg and Waage. — Guldberg, who was later
professor of applied mathematics at the University of Christiania,
and Waage, professor of chemistry at the same institution, were the
first to mathematically formulate the effect of mass on chemical
activity. Their first preliminary paper was published iu Norwegian
in 1864. Their epoch-making paper appeared in 1867. In the first
part of their paper they review the theories of affinity which had
been held. The views of Bergman n and Berthollet are taken up, and
it is pointed out that neither is sufficient to account for all the facts
known. They attributed this to the lack of a suitable method for
determining the magnitude of affinity. They point out that the
method of Bergmann, based on the assumption that if the substance
B replaces C from a compound with A, giving the compound AB,
the affinity between A and B is greater than between B and C, is not

92 PRIKCIPLES OF lNOUG.iNIC CIIEMISTKY

satisfactory, since this assumption leaves out of account a large
number of conditions which affect the reaction. The attempt to
measure the magnitude of chemical affinity by the heat evolved dur-
ing the reaction was regarded as unsatisfact<.)ry, because it depends
in part upon the conditions under which the reaction takes place,

Guldberg and Waage point out that in chemistry, as in mechanics,
we must study forces by their effects, and the most natural method
is to determine forces in the condition of eqailibriura ; " that is to
saji we must study the chemical reactions in which the forces which
produce new compounds are held in equilibrium by other forces.
This is the case in the chemical reactions where the reaction is not
complete but partial, i.e. in the reactions where —

"(a) Addition and decomposition take place at the same time^

and where,
" (b) Substitution and reformation proceed simultaneously,"

The authors do not take up in this paper the case of addition
and decomposition, or dissociation, since the dat-a available are not
sufficient, but develop the law of mass action from a study of the
second class of reactions, viz. substitution.

In the development of the law their own words are given : —

"Let us assume that two substances, ^1 and Bj are transformed
by double substitution into two new substances. A* and B'j and
under the same conditions A* and B* can transform themselves into
A and B. Neither the forjnation of A^ and B* nor the reformation
of ^1 and B are comidete, and at the end of the reaction we have
the four substances present Af B, A\ and B*. The force which
causes the formation of A* and B' is in equilibrium with that which
causes the formation of A and B. The force which causes the
formation of ^4' and B* increases projKirtioual to the affinity coeffi-
cients of the reaction A-\- B=^ A* + B\ but it depends also on the
masses of ^4 and B.

** We have learned from our experiments that, the force is propor*
tional to the prodmt of the adim mas$es of the two .wbManc€» A aiid B.

"If we designate the active masses of A and B by p and q, and
the affinity coefficient by A' the force = A' . ;i . q.

** As we have often observed, the force Kpq^ or the force between
A and B, is not the only force which comes into play during the
reaction. Other forces tend to retard or accelerate the formation of
A' and B\ Let us, however, assume that other forces do not exist,
and let us see what formula is develop^ in this case. We believe
that the consideration of this ideal reaction, where only the forcea

MOLECULAR WEIGHTS OF DISSOLVED SUBSTANCES 93

between A and B, and between A* and J?' are taken into account,
will furnish the reader with a clear and distinct presentation of our
theory.

" Let the active masses of A and 5' be p^ and q\ and the affinity
coefficient of the reaction -4' 4- -B' = -4 4- A he A"' ; the force of the
reformation of A and B is equal to K'p'q'. This force is in equilib-
rium with the first force, consequently, —

Kpq = A>V. (1)

"By determining experimentally the active masses p, 7, ;>', and q',
we can find the relation between the affinity coefficients A' and A"'.

A''

On the other hand, if we have found this relation — , we can calcu-

A
late the result of the reaction for any original condition of the four
substances."

MOLECULAR WEIGHTS OF DISSOLVED SUBSTANCES

Determination of the Molecnlar Weights of Dissolved Substances
by the Freezing-point Method. — We have already learned that the
freezing-point of water is lowered by the presence of dissolved sub-
stances. The amount of the lowering has been shown by the French
physical chemist Raoult, to be proportional to the ratio between the
number of molecules of the solvent and the number of molecules of
the dissolved substance. If we know the lowering of the freezing-
point of any solvent produced by dissolving in a given volume of
that solvent a number of grams of any undissociated substance equal
to the molecular weight of the substance, we can then use the freez-
ing-point lowering to determine the molecular weight of any sub-
stance in that solvent. If we dissolve a gram-molecular weight of
an undissociated substance in a hundred grams of the solvent, the
resulting freezing-point lowering is known as the freezing-point con-
stant of the solvent. Knowing the freezing-point constant for any
solvent, it is a comparatively simple matter to determine the molecu-
lar weight of any substance in that solvent.

We must weigh the solvent, also the amount of substance to be
dissolved in the weighed amount of the solvent, and determine the
lowering the freezing-point produced. Let the weight of the solvent
be W, the weight of the dissolved substance ir, the lowering of the
freezing-point produced A, and the freezing-point constant C. The
molecular weight of the substance M is calculated as follows : —

94

PRINCIPLES OF INORGANIC CHEMISTRY

The freezing-point constants of some of the more common solvents
are given below : —

COKBTAWT

Acetic acid

39.0

Benzene

50.0

Ethylene bromide

117.9

Formic add

27.7

Nltro benzene

70.7

Water

18.6

Apparatus devised by Beckmann. — The

most convenient form of apparatus used
for determining the value of A is that
devised by Beckmann. It is shown in
Fig. 17. The glass vessel A is to receive
the solvent or solution whose freezing-
point is to be determined. The substance
can be introduced through the side-tube,
but the latter can be readily dispensed
with. The tube A passes through a cork
into the wider glass tube Aiy and an air-
space exists between the walls of the two
tubes. The thermometer T is inserted
into A, and fastened tightly in position
by means of a cork. The liquid in A is
stirred by means of a glass rod bent in a
circle of sufficient diameter to allow the
bulb of the thermometer to pass through.
The stirrer is attached to a vertical rod S,
and moved up and down by means of the
hand. ^ is a battery jar, which contains
the freezing-mixture. The substance used
iu the jar depends upon the freezing-point
of the solvent with which we are dealing.
If the solvent freezes appreciably above
IK« freeiing-point of water, it is only
«^C#tfliMy to use water and ice. If we are
^iMfkiug with water as the solvent, the
ffe^mttU^MftUture more commonly used is
i^* Jliiki iHiJtt Care must be taken that not

MOLECULAR WEIGHTS OF DISSOLVED SUBSTANCES 95

too mnch salt is used, since, when the mixture is too cold, the results
obtained are oSten not reliable.

The thermometer used by Beckmann requires spec^ial comment.
It is constructed on a different plan from that of any other ther-
mometer which has erer been employed. In the first place, the Imlb
is very large, and, consequently, the divisions on the scale correspond
to a very small range in temperature. The largest scale divisions
correspond to degrees. The total range of such a thermometer is
usually about 6^ The next smaller divisions correspond to tenths of
a degree, and the smallest divisions to hundredths of a degree. By
means of a small lens it is possible to read the scale to
thousandths of a degree.

The unique feature of the Beckmann thermometer is,
however, the arrangement at the top. This is seen in
Fig. 18.

The capillary terminates in a reservoir or cistern, into
which, by warming the bulb, mercury can be driven. The
mercury in this reservoir can be thrown either to the top
or bottom by holding the thermometer and tapping or
thrusting it By this means it is possible to inereuse or
decrease the amount of mercury in the bulb of the ther-
mometer, and to so adjust the amount that the top of the
column will come to rest at any desired point on the scale,
when the instrument is placed in the freezing solvent.
The freezing-point of any solvent or solution can, then, be
adjusted at any desired position on the scale, and the dif- ^^'

ference between the freezing-points of the solvent and solution deter-
mined. This differential thermometer of Beckmann has proved of
incalculable service to physical chemistry, and has contributed more
to our knowledge, in the field which we are now studying, than any
invention or device which has ever been proposed.

Determination of the Molecular Weights of Dissolved Substances
by the Boiling-point Method. — The determination of the molecular
weights of dissolved substances by the boiling-point method is strictly
analogous to the determinations by the freezing-point method. The
boiling-point of a solvent is raised by the presence of dissolved sub-
stances, and the rise in boiling-point has been shown to be propor-
tional to the lowering of the freezing-point. The rise in the boiling-
point, like the lowering of the freezing-point, depends upon the ratio
between the number of molecules of the dissolved substance and the
number of molecules of the solvent. If we know the rise in the
boiling-point of a solvent produced by a gram-molecular weight of

96

PRINCIPLES OF INORGANIC CHEMISTRY

an undissociated substance in a hundred grams of the solvent, —
the boiling-point constant of the solvent, — the determination of the
molecular weight of any substance in that solvent is a compara-
tively simple matter.

If we represent the weight of the solvent used by TT, the weight of
the dissolved substance by w, the rise in the boiling-point by By and
the boiling-point constant by C, the molecular weight of the sub-
stance M is calculated as follows : —

M=

Cw

The boiling-point constants of a few of the more common solvents
are given below : —

COXSTAMT

Constant

Acetone

Aniline

Benzene

Carbon disulphide .
Chloroform. . . .

17.1
32.0
26.1
23.6
35.9

Ether

Ethyl alcohol . . .
Methyl alcohol . .
Water

21.6

11.7

8.4

6.1

Boiling-point Method of Beokmann. — The rise in the boiling-point
of a solvent produced by a dissolved substance was determined for a
long time by the method of Beckmann. The apparatus which he
employed is shown in Fig. 19. The glass tube A contains the liquid
whose boiling-point is to be determined. Into this liquid the ther-
mometer dips as shown in the figure. In the bottom of the
tube are placed glass beads, garnets, or platinum scraps, so as to
secure a more uniform rate of boiling. A condenser is attached to
the tube A, as shown in the figure. This tube is surrounded by a
double-walled, glass jacket B, into which is introduced some of the
same liquid whose boiling-point is to be determined in A. This is
also provided with a return condenser. The liquid in B is boiled at
the same time as the liquid in Ay so that the innermost vessel is sur-
rounded by a layer of liquid having the same boiling-point. The
whole apparatus rests upon an asbestos box, and heat is supplied
by a flame placed beneath. Beckmann has devised a number of
modifications of this apparatus, but in the opinion of the writer
none of them represents any marked improvement on the form just
described.

MOLECULAR WEIGHTS OF DISSOLVED SUBSTANCES 97

The pure solvent is poured into the tube Ay the filling-material
(beads or genets) introduced, and the thermometer inserted so that
when the cork is forced into the top of tube A^ the bulb of the ther-
mometer is entirely covered by the liquid, but does not touch the
glass beads. The mercury in the
Beckmann thermometer is so ad-
justed that the top of the column
comes to rest between the divi-
sions 0^ and 1^ when the solvent
boils. The vessel A is then care-
fully cleaned and dried, and after
introducing tlie filling-material
a weighed amount of the solvent
is poured in. The thermometer
is inserted and the condenser
attached. Some of the pure sol-
vent is poured into the vapor-
jacket, and boiled simultaneously
with that in the tube A, The
position of the mercury is care-
fully noted on the thermometer,
after the solvent has boiled about
twenty minutes, and the barom-
eter is also very carefully read.
The flame is now removed and
the solvent allowed to cool.

The substance whose molecu-
lar weight is to be determined is
pressed into tablets, weighed, and
introduced into the solvent. The
boiling is renewed after all the
substance has dissolved, and the temperature at which the solution
boils carefully noted on the thermometer. The barometer is read
again, and if any change has occurred the proper correction is
introduced into the readings on the thermometer. Care must
always be taken to tap the thermometer before making a reading.
The difference between the boiling-point of the solvent and that
of the solution is the rise in boiling-point produced by the dissolved
substance.

Boiling-point Apparatus of Jones. — A number of attempts have
been made to improve the boiling-point apparatus of Beckmann.
The following form was devised and used by Jones : —

Fig. 19.

H ^8 PRINCIPLES OF INORGANIC CHEMISTRY ^^H

^^^P Into the glass tube A (Fig. 20) some glass beads or garnets are in- fl

^V trotiuced. To the side-tube A the condenser is attached. Into the

1 1 1

beads a cylinder of platinum /^

1 i

is inserted by placing the finger

♦

upon the top of the cylinder and

-

gently shaking the whole appa-

ratus* The liquid whose boil-

^^^^H

- *

^ ing'point is to be determined is

■

introduced into ^4 until the bulb

of the thermometer, placed as

- a

shown in the figure, is covered.

^

The liquid must not come within

a centimetre^ or a centimetre and

- 1

a half, of the top of the plati-

num cylinder. The tube -^l is

surrounded by a thick jacket

I Q

of asbestos J, and rests on an

-

^^^^^

^

lar hole is cut, and over which a

J H

piece of wire gauze is laid. Heat

^^^^H

is supplied by means of a very

^^^^^^B

\ y^y9 small flame iJ, placed beneath

^y^H the apparatus and protected by a

y^ metalUc screen as shown in the

^^^^H 1^ ' P

f^ drawing.

^M Hi

^B The essential difference Ije-

H tween this apparatus and other

1 ^Hj forms is the platinum cylinder

^E 3

IP ^^ '* which is introduced into the boil-

^f9

P^Hl ^^^ liquii The object of this

fe WM cylinder is twofold. It prevents
^BW the cooled, recondensed solvent
W^^ 1 feom coming in contact with the

^l^^^^Ly thermometer before it ja reheated

^^^1 ^t^^3 m^MH. w toe iK)iimg>pomt, xt reauces

^m ^f / jt^ \ ^m *^^ effect of radiation to a mini-

H ■ / © \ 1 J»«^ ^^ the bulb of the ther-

■ 1 JS

J \ H mometerissarroiindedonlybythe

m 1

M boiling liquid, or even if a layer

^^^ '

* of asbestos is wrapped around

^^^ ^'**' '** the glass tube, heat will be radi<

^^^ ated out from the hot bulb on to colder objects in the neighborhood.

MOLECULAR WEIGHTS OF DISSOLVED SUBSTANCES 99

The tempeiatoie of the bulb will always tend to be a little lower
than that of the boiling liquid in which it is immersed. By sur-
rounding the bulb with a piece of metal as nearly as possible at the
same temperature as the bulb itself, the effect of radiation is reduced
to a minimum.

The apparatus is exceedingly simple, and when applie<l to the
determination of molecular weights of dissolved substances, was
found to give good results in both low-boiling and high-boiling sol-
vents. Another application of this method will be considered a
little later.

CHAPTER VIII

OSMOTIC PREiaatTRE AND THE THEOBT OF ULECTROLTTIC

DISSOCIATION

Osmotic PresBnrd. — Having studied water as a sol vent, we can
turn to a class of phenomena which has come into great prominenee
in the last two years, and which directly and indirectly has thrown
mnch light on chemical phenomena in general. If a solution of a
substance in a solvent like water is placed in a vessel, and over this
solution the pure solvent poured^ we w^ould find after a time that the
substance is not all contained in that part of the solvent in which it
was originally present, but a part of it has passed into the layer of
the pure solvent which was poured upon the solution. This shows
that there is some force analogous to a pressure, driving the dis-
solved substance from one region to another, from the more con-
centrated to the less concentrated solution. This pressure has been
termed osmoU'c presmire.

Demonstration of Osmotic Pressure. — The existence of this press-
ure was early recognized. Abl>e Nolle t demonstrated its existence
about the middle of the eighteenth century. A glass tube closed at
the bottom with animal parchment w^as filled with ordinary alcohol,
and the tube then immersed iu w^ater. Water could pass in through
this parchment, but alcohol could not pass out. The contents of
such a tube gradually increased in volume, showing to the eye the
existence of osmotic pressure. During the first three-fourths of the
last century osmotic pressure was demonstrated by filling an animal
bladder with an aqueous solution of alcohol^ and immersing the
bladder in water. The water passed into the bladder and the alco-
hol could not pass out in any quantity. Hence, the bladder becaine
distended and finally burst. It will be observed that in all of these
experiments recourse waa had to animal membranes. A discovery
was subsequently made, which has entirely done away with the use
of natural membranes in demonstrating osmotic pressure.

These membranes, which have the property of allowing the sol-
vent to pass through them, and of preventing the dissolved suIj-
stance from passing, aie known as semi-p^rmeabk. It was M* Traube

100

OSMOTIC PHESSUKE 101

who first prepared such semi-permeable membranes artificially. Ho
found that certain precipitates, deposited in a suitable manner, have
the property of allowing the solvent to pass through them, but hold
back the dissolved substance. These precipitates in(*lude cop^n'r
ferrocyanide, and a number of similar gelatinous substunces. A
method of demonstrating osmotic pressure, now that we can jtroimrt)
artificial membranes, is the following : A glass tube almut 2 cm. in
diameter and 8 to 10 cm. long, is tightly closed at the bottom with
v^etable parchment This is soaked in water for some hours so as
to drive out air-bubbles. The top of the glass tulx) is tiglit ly closed
with a rubber stopper, through which is passed a fine capillary tuln)
about a metre in length. The end of the capillary should just pass
through the cork, but must not protrude beyond its lower surfiwu*.
The large glass tube is now immersed in a beaker, which is suf-
ficiently deep to receive the entire tube. Tlie tube is tln»n firmly
clamped in a vertical position. The beaker is filled with a three
per cent solution of copper sulphate. The cork is then removed
from the tube, and the latter completely filled with a three ptM- cent
solution of potassium ferrocyanide, to which enough potassium nitrate
has been added to make from a one to a two i)er cent solution. The
tube is then closed as tightly as possible with the cork through
which the capillary passes, care being taken that no air-bubble
remains beneath the cork. The apparatus is then set in a quiet
place for some days. After a day or two, if the experiment is suc-
cessful, the liquid will begin to rise in the capillary, and may reach
a height of from 40 to 50 centimetres.

The experience of the writer has been that not all such exjjeri-
ments succeed. Indeed, the number which give a good demonstra-
tion of osmotic pressure is only alK)ut one-third of the total attein])ts
which he has made. The frequent failure is doubtless due in part
to the nature- of the parchment used.

The method by which the semi-permeable membrane is formed in
this ca^ is almost self-evident. The copjwr sulphate from Ik»1ow
passes into the parchment, and the potassium ferrocyanide from
above also enters the parchment. The two meet right in the walls
of the vegetable parchment. At the surface of contact they form
the gelatinous precipitate of copper ferrocyanide in the walls of the
parchment. The precipitate, deposited in this manner, has the proj)-
erty of semi-permeability — it allows the water to pass through and
prevents the dissolved substances from passing. Since osmotic press-
ure always acts so that water passes from the more dilute to the
more concentrated solution, the flow of water in this case is from the

102

PRINCIPLES OF INORGANIC CHEMISTRY

copper sulphate on the outside to the potassium ferrocyanide and
potassium nitrate on the inside. The liquid rises in the capillarj
due to the inflow of water through the semi-permeable membrane*

Morsels Hethod of Preparing Semipermeable Hembranes. — The
demonstration of osmotic pressure has now Ijecome a very simple
matter, due to a method derised in this laboratory by Morse, and
developed by Morse^ Horn, and Frazer.

*' It occurred to the authors that if a solution of copper salt and
one of potassium ferrocyanide are separated by a porous wall which
is filled with water, and a current is passed from an electrode in the
former to another electrode in the latter solutionythe copper and the
ferrocyanogen ions must meet in the interior of the wall and sepa-
rate as copper feri'ocyanide at all points of meeting, so that in the
end there should be built up a continuous membrane well supported
on either side by the material of the wall/^

In order to remove the air contained in the walls of the cup they
made use " of the strong endosmose wluch appears when a current is
passed through a porous wall separating two portions of a dilute solu-
tion in which the two electrodes are immersed.'- A dilute, boiled
solution of potassium sulphate was used for this purpose, ** On pass-
ing the current between the electrodes in the direction of the one
within the cup, the liquid in the cup rises with a rapidity which
increases with the dilution of the solution, and with the intensity of
the current. The water, in passing through the wall, a|ipearfl to
sweep out the air in an effective manner.'^

Ha^'ing removed the air by means of endosmosis, the membrane
was formed by filling the cup with a tenth-normal solution of potas-
sium ferrocyanide, and Immersing it in a tenth-normal solution of
copper sulphate. One electrode of platinum wa5 inserted into the
cup, and the other of sheet copper completely sm-rounded the cup.
The current was passed from the copper to the platinum electrode.
As soon as the copper ions, moving with the current, came in con-
tact with the Fe(CN), ions moving against the current, a precipitate
of copper ferrocyanide waa formed in the wall of the cup. This
gradually became more compact, as was shown by the fact that the
resistance offered to the passage of the current rapidly increased.

The advantage of driving the ions into the wall by means of the
current is that the membrane can be formed much more compactly
than by simply allowing them to pass itito the wall by diffusion.
With such a cell it is possible to demonstrate osmotic pressure in a
most satisfactory manner. When tlie cell is filled with a normal
Bolution of cane sugar^ closed with a cork through which a capillary

OSMOTIC PRESSURE

103

monometer passes, and immersed in pure water, the liquid will rise in
the capillary at the rate of more than a foot an hour, and in two days
a pressure of thirty feet of the sugar solution is easily secured. This
so far surpasses all other demonstrations of osmotic pressure thus
far devised, that they become insignificant by comparison. The
demonstration of osmotic pressure on the lecture table by means
of this method has become as simple a matter as many of the daily
experiments in inorganic and organic chemistry.

This method promises much for the quantitative study of osmotic
pressure. The ease with which the cells can be prepared, using
suitable porous cups, and the great resistance offei*ed by the mem-
branes formed by the electrical method, bid fair to open up new
possibilities in connection with the direct measurement of
osmotic pressure. A number of measurements of the osmotic
pressure of solutions of cane sugar and glucose, of concen-
tration as great as normal, have already been made. Ftes-
sures as high as 31 A atmospheres have been measured.

Several other semi-permeable membranes have already
been prepared by Morse and his co-workers, using the electro-
lytic method. Of these perhaps the most important is
ferric hydroxide, since this substance can be employed in
investigating alkaline solutions.

Measurement of Osmotic Pressure.
— Certain measurements of osmotic
pressure were made by W. Pfeffer
twenty-seven years ago. He made
use of the artificial membranes which
had been discovered by Traube, and
deposited them upon a support which
was sufficiently resistant to enable
them to withstand considerable press-
ure. Unglazed porcelain cells were
injected with water and placed in a
solution of copper sulphate. After
a time they were filled with a solu-
tion of potassium ferrocyanide. The
two substances enter the walls, the
one from the inside, the other from
the outside, and form a precipitated
membrane of copper ferrocyanide.
This appears as a fine, reddish-brown
Fia. 21. lii^e ill the walls of the porcelain.

104

PRIXCIPLES OF INORGANIC CHEMISTRY

ZZH

The membrane ODce formed prerents either of the substances from
passing through, and hence it appears as a fine line. In Fig. 21 is
shown the apparatus used bj Ffeffer. The manometer m is one-half
natural size. The sketch is a longitudinal section of the apparatus*

The cell z used by Pfetfer waa
only about 4f} mm, high and
16 mm. internal diameter.

The measurements of os-
motic pressure were made by
means of these porcelain cells
lined with tlie precipitate,
which formed the semi-per-
meable membrane. After the
manometer was attaclietl to
the cell, the latter was filled
with the solution whose os-
motic iiressure was to be
measured. The cell was then
tightly closed and fastened to
a i^^lass rod as seen in Fig. 21*

The whole cellj including
the manometer, was introduced
into a bath as shown in Fig. 22.
The bath was tilled with pure
water, and the osmotic presB-
nre of the solution against
pure water measured on the
itLf iriiry manometer. Special
ptL-autiotis were taken to
keep the tenn>erature of the
whole apparatus constantj
since, as we shall see, there is
a Iftrge temperature coefficient of osmotic pressure. The temperature
of the experiment was accurately determined by means of carefully
standardized thermometers.

"1

/;<

Fia. 22.

BELATIONS BETWEEN OSMOTIC PRESSURE AND GAS-
PKESSURE

Pfeifet carried out the measurements already referred to, and
doubtless saw their physiological sigtuficance, but he did not |)oint
out any relations between osmotic pressure and gas-pressure. This^

OSMOTIC PRESSURE 106

like so many other brilliant discoveries, was reserved for Van't Hoff.
In his epoch-making paper, he points out a number of surprisingly-
simple relations, and some of these will now be taken up.

Boyle*8 Law for Osmotio Pressnre. — The law of Boyle for gases
states that the pressure of a gas varies directly as the concentration
of the gas. From Pfeffer's results, it has been shown that the
osmotic pressure of a solution varies directly with the concentration.
This relation for the osmotic pressure of solutions certainly suggests
the relation for gases expressed by the law of Boyle.

(hty-Lussao's Law for Osmotic Pressure. — According to the law
of Gay-Lussac the pressure of a gas increases with the temperature,
at the rate of -^^ for every rise of V C. Pfeffer's results show
that the osmotic pressure of a solution increases with rise in tem-
perature, and the rate of increase is very nearly ^fy for every
degree. Pfeffer did not make an extensive study of the tempera-
ture coefficient of osmotic pressure, but as far as his results go they
led to the conclusion stated above.

If the law of Gay-Lussac applies to the osmotic pressure of solu-
tions, then, solutions which are isosmotic, or have the same osmotic
pressure at one temperature must remain isosmotic at other tem-
peratures, since they would have the same temperature coefficient of
osmotic pressure. This has been tested by the methods for deter-
mining relative osmotic pressures. Hamburger found that solutions
of potassium nitrate, sodium chloride, and cane, sugar, which were
isosmotic at 0®, were also isosmotic at 34®.

There is, however, a still more striking experimental verification
of the applicability of the law of Gay-Lussac to solutions. If a tube
is filled with a gas and all parts of the tube kept at the same temper-
ature, the concentration of the gas will be the same in every part
of the tube. If, on the other hand, one portion of the tube is kept
warmer than the others, the gas will so distribute itself throughout
the tube that the pressure will remain the same in all parts of the
tube. Since the pressure of gas increases with the temperature,
each particle will exert a greater pressure in the warmer region,
and, consequently, there will be fewer particles required in the
warmer portion of the tube to exert the same pressure as exists in
the colder portion. In a word, the gas would tend to become more
concentrated in the colder portion, and more dilute in the warmer
portion of the tube.^

1 It shonid, of course, be remembered that the condition described for a gas is
somewhat ideal. The gas particles, due to their rapid movement, would mix, but
the principle which it is desired to iUustrate holds good.

. - ?> fi: rSORt.AXIC CHEMISTRY

--^^^ * .-. <ifi]nr.jfliif obeys the laws of gas-press-

;. :. :\n uhc>xe should be observed with

- ■• '«.'•: ■ : :.i)f iwo i>arts of a perfectly homo-

: * ^ ..■.f?"rf^T:* :<^nii>eratures for any consider-

^ .: .! Sf»,'*.-*:wes more concentrated in the

*•" -s :.-^ ,vi::u' to be known from its discov-

v^. -,. r J-;< principle is of the very greatest

. ^ •. ,-- » 'L \iiy-Lussac for osmotic jiressure.

V ■•- .. :\ -viUjtfr jx>rtion of the solution should

...*.i^-. \> .^., :or every difference of one degree

• > • ..^ X' ^tjtsily tested by experiment. The

.> .\* A-. y-- <cz^t by placing the solutions in

i..i:!vr -iiAS the upper portions of the tubes

.. v^^wM. t :ti; »^'i^:«n\ and the lower portions cooled

. .. ..V. \^\%s ^^\wT experiments of Soret gave

. .. .. . .JL ^ -^v^ >(nu? not quite as great as that

,, X>-J1 ::;ssAi\ His later experiments, in

. . V . . .^ -.\:. vo stand at constant temperatures

. ■ VjL-'itvv?^ nV.ioh, while a little too low, yet

tS. Kvok, A slight difference between

^v • x*\v?t cw^Uos no surprise when we con-

. .. .. M >a;?.^s\ fvV months at the constant tem-

. . .^. .':\.;t^ r,;Ay Ih> reached, and some mixing

^.....vo ,s ;jirriug is, therefore, unavoidable.

^ ML^ ,.\t«o that it is now quite certain that

. _ v>x vx ■ V N\<t proof of the applicability of

. . , N .\v«N>^v pn^ssure of solutions.

^ v^^^ ^ ^ Omotio Pressure of Solutions. —

\ . » X . ; tVy'o rtud (lay-Lussac to the osmotic

^ vN • V -vo tJxis quantity is analogous to gas-

,.,■ ^vv T V quostiou as to the relative magni-

. v^ ^ -v .-:-:\\x wnanswored. The one might be

\ -* ,i.\ sinAU, and still the two laws which we

. . ^. . .;> ro N'^h. Wo now come to the question,

,. < N»(^ivn tho magnitudes of the two press-

.. . :v»:hMo osWilions?

..v.,::s\ appliotl to gases, states that in equal
..N M \W same tomix^rature and i)ressure, there
s . ,^! Ultimate ]>arts. If the law of Avogadro
., : w.^uld N^ stat^Hi thus; In equal volumes of
' ,• sa?iio tom|>orature have the same osmotic press-
. ' od t he same numlwr of dissolved particles. The

OSMOTIC PRESSURE

107

simplest way in which this law can be tested for solutions is to see
what relation exists between the gas-pressure of a gas-particle and
the osmotic pressure of a dissolved particle under the same conditions
of temperature and concentration. Let us compare the gas-pressure
of hydrogen gas and the osmotic pressure of cane sugar in water.
Given a one per cent solution of cane sugar ; such a solution would
contain one gram of sugar in 100.6 cc. of water, and the osmotic
pressure of such a solution can be calculated from Pfeffer's results.
Hydrogen gas, having the same number of parts in a given volume,
would have the following pressure : The molecular weight of cane
sugar is 342, that of hydrogen 2. The hydrogen gas must, therefore,
contain -^^ grams in 100.6 cm., which is the same as 0.0581 grams
per litre. Hydrogen gas at 0®, and at a pressure of one atmosphere,
weighs i)er litre 0.08995 gram ; the above concentration of hydrogen

0 0541
gas will, therefore, exert a gas-pressure of ' = 0.646 atmos-

phere at 0**.

It is now only necessary to compare the osmotic pressure exerted
by the cane sugar with the gas-pressure, to see if any simple rela-
tions exist between the two. The following table of results is taken
from the paper by Van't Hoff : —

TUfrSKATUBE

Osmotic Pbkmubb of
Cakb Scoab

Qas-pbbbspbb or
IIydboobk Gas

6^8
13°.7

36°.0

0.664
0.691
0.684
0.746

0.665
0.681
0.686
0.735

The remarkable fact is established by these results that the
osmotic pressure of a solution of cane sugar is exactly equal to the
gas^essure of a gas having the same number of parts in a given
volume, temperature being the same in both cases. Under the same
conditions, then, a dissolved particle exerts the same osmotic press-
ure that a gas particle exerts gas-pressure.

Causes of Oas-pressure and of Osmotio Pressure. — That there
should be an equality between these two pressures is very surpris-
ing, if we consider the great difference between the phenomena with
which we are dealing. Gas-pressure is explained in terms of the
kinetic theory of gases, as due to the particles of gas bombarding
against the walls of the confining vessel. It should be stated that
we do not know what is the cause of osmotic pressure. A great

108

PRLXCrPLES OF INORGANIC CHEMISTRY

number of explanations and theories have been offered to account
for osmotic pressure, but in tbe opinion of the writer no one of them
is at all satisfactory. Some ha%'e attempted to account for osmotic
pressure by the attraction of water by the dissolved Hubstance, but
this is only a renaming of the phenomenon, and in no senae an
explanation of it. Others have suggested that water passes through
the semi-permeable membrane from the more dilute to the more con-
centrated solution, because of the screening action of the dissolved
particles. These cannot pass through the membrane, and, therefore,
screen it from the blows of the solvent. Since the greater screen-
ing influence is exerted on the side containing the larger number of
dissolved particles, we have the flow of the solvent from the more
dilute to the more concentrated solution. A careful analysis of this
explanation shows that it is not sufficient The screening influence
of the dissolved particles would be just as great below as it is above,
keeping the water which has passed through the membrane from
rising, since the membrane is quite permeable to water. It is, there-
fore, fairest to say that w^e have at present no satisfactory theory to
account for that phenomenon known as osmotic pressure-

EzoeptioEi to the Applicability of the (Jat Laws to Osmotic Pret»-
ure. — We have just seen that the three best known laws of gas-
pressure apply to the osmotic pressure of solutions of substances
like cane sugar. We might conclude from this that the laws of gas-
pressure always apply to the osmotic pressure of solutions of all
substances. Such is not the case. Van^t Hofif pointed out that there
are not only exceptions to this generalization, but a great many
exceptions. Indeed, the substances wluch present exceptions are
quite as numerous as those which conform to rule. The osmotic
pressure of most salts, of all the strong acids, and all the strong
bases, is much greater for all concentrations than would be expected
from the osmotic pressure of solutions of substances like cane sugar
for the same coneent rations. The osmotic pressures of these three
classes of substances are always greater than would be expected from
the laws of gas-pressure applied to the osmotic pressure of solutiona

The general expression for the laws of Boyle and Gay-Lussac is,

we have seen, —

This applies directly to the osmotic pressure of Kdutioiii of mb*
stances like c^ine sugar. But in order that it may apply to solutions
of salts, acids, and bases, a coefficient must be introduced, which,
for these substances, is always greater than nnity. This coefficient

OSMOTIC PRESSURE 109

was called by Van't Hoff *, and it has come to be known as the Van't
Hoff !.

The above expression when applied to acids, bases, and salts,
becomes— pv = iRT.

While these exceptions were clearly refcognized by Van't Hoff, he
was unable to explain them, or to offer any satisfactory theory to
account for them.

In this case, as in so many others, the exceptions are as interesting
and important as the cases which conform to rule. We shall see
that these exceptions led to a theory which is one of the most im-
portant in modern chemical science.

ORIGIN OF THE THEORY OF ELECTROLYTIC DISSOCIATION

Work of Arrhenins. — Arrhenius was impressed by the generali-
zations reached by Van't Hoff connecting gas-pressure and osmotic
pressure, and especially by the large number of exceptions to these
generalizations. Referring to the equality of gas-pressure and osmotic
pressure under the same conditions, Arrhenius found a difficulty in
that the generalizations reached by Van't Hoff, connecting gas-press-
ure and osmotic pressure, held only for a large number of substances
but by no means for all. The aqueous solutions of a great number
of substances exerted a larger osmotic pressure than they should do
if the genei-alization of Van't Hoff applied.

When a gas shows a deviation from the law of Avogadro we
assume that it is dissociated, and verify the assumption experimen-
tally. The same assumption may be made in the cases of substances
which present exceptions to the laws of Van't Hoff.

Arrhenius then puts forward the assumption of the dissociation
of. certain substances dissolved in water to explain the exceptions
to Van't Hoff's generalization. Osmotic pressure is, as we have
seen, proportional to the concentration of the solution. This is
the same as to say that osmotic pressure is proportional to the
number of dissolved particles. If a substance exerts an abnormally
great osmotic pressure, there must be more parts present in the
solution than we would expect from the concentration. But acids,
bases, and salts, represented by hydrochloric acid, potassium hydrox-
ide, and potassium chloride, are the substances which show the
abnormally great osmotic pressure. How is it possible to conceive
of substances such as these breaking down into any larger number of
parts than would correspond to their molecules ?

110

PRINCIPLES OF IJf ORGANIC CHEMISTRr

This 19 the problem which must be solved, and Arrheniiis has
solved it, as we believe, satisfactorily. He went back to the theory
proposed by Clausius to account for the facts which were known in
eonnection with the phenouieuon of electrolysis. It was found that
an intiiiitely weak current will decompose water to which a Httle
acid is addedj Hberatiug hydrogen at one pole and oxygen at the
other. If the aqueous solution of the acid contained only molecolesj
in order that we might have electrolysis the current must be capable
of dec oni posing the molecules. The fact is that a current far too
weak to decompose a molecule of water will effect electrolysis,
Tberefore, some of the molecules present iu the solution^ either
those of the w^ater or of tbe acid, must be already broken down be-
fore the current is passed, Clausius did not claim that the mole*
cules are broken down into their constituent atoms. Such a theory
would be absurd* His theoi-y ivas that the molecules are broken
down into parts, wliich he called ionSj and each ion is charged with
electricity, either positively or negatively. An ion may be a charged
atom or a charged group of atoms.

The theory that molecules are broken down into ions by a solvent
like water was proposed, tljen, by Clausius in 1S56,

A similar theory was advanced by the chemist Willianison in
1851, as the result of his work on the synthesis of ordinary ether
from alcoLol and sulphuric acid. The theory of Clausius differed
from that of Williamson, in that the former assumed that there are
only a few molecules broken down into ions, while "Williamson
thought that most of the molecules present are in a state of decom-
position. It should be observed that lM>th of these theories are
purely qualitative suggestions. The one thougljt that only a few
molecules in .solution are broken down into ions, the other, that we
have to do mainly ivith ions ; but neither suggested any method by
which we could determine the actual amount of the dissocdation in
any case.

The new featum which w^tis introduced by Arrhenins w^as to
point out a method for determining just what per cent of the mole-
cules is broken down into ions. He thus con%*erted a purely qualita-
tive suggestion into a quantitative theory^ which could be tested
experimentally.

The Theory of Electrolytic Diiseciation. — The theory of ele<3-
trolytic di?>soriation, as we have it tonlay, states that when acids,
bases, and salts are dissolved in water, they break down or
dissociate into iona. Examples of the three classes are the fol-
lowing:—

OSMOTIC PRESSURE 111

»HC1 = H, Ci.

KOH = K, OH.

KC1 = K, Ci.

Each compound dissociates into a positively charged part called
a cation, and a negatively charged part, an anion. These. ions may
be charged atoms as the above cations, or groups of atoms as the
anion OH. The cations are usually simple atoms charged with posi-

tive electricity. The cation of all acids is hydrogen, H ; the nature
of the anion varies with the nature of the acid. It may be chlorine,
bromine, the NOj group, SO^, etc. The anion of bases is the group

(OH); the cation varies with the nature of the base. It may be
potassium, barium, ammonium, etc. The anions and cations of
salts both vary with the nature of the salt. They depend upon
the nature of the acid and the base which have combined to form
the salt

Heasniement of Electrolytio Dissociation. — Although electrolytes,
which we remember include acids, bases, and salts, are dissociated
by water into ions, it is not true that all electrolytes are completely
dissociated under all conditions. Indeed, no electrolyte is com-
pletely dissociated by water, and still less by other solvents, unless
the dilution of the solution is very great. The strongest acids,
bases, and salts are completely dissociated only when the dilution is
so great that a gram-molecular weight of the substance in question
is dissolved in from five hundred to one thousand litres of water.

Electrolytes are dissociated, however, to a greater or less extent
at all dilutions, and it is always a matter of interest and frequently
a matter of importance to know the degree of the dissociation under
the conditions in question. '

Several methods are available for measuring the amount of disso-
ciation of any electrolyte in a solvent like water. All electrolytes
give greater lowering of the freezing-point and produce greater rise in
tJie boiling-point of water than non-electrolytes. When we were dis-
cussing the determination of the freezing-point and boiling-point
constants it was stated that non-electrolytes must be used. The
reason is now apparent. Electrolytes being partly dissociated, con-
tain a larger number of parts in solution than would correspond to
their molecules. Since lowering of freezing-point and rise in boil-

^ The comma between the two ions in this and all subsequent ionic equations
means that the ions were combined as a molecule, or can combine and form a mole-
cule.

112

PRINCIPLES OF INORGANIC CHEM18TET

ing-point are properties which depend only upon nmnbers, the larger
the number of parts present, the greater the value of these quantities.

Knowing the lowering of the freejsing-point aud the rise of the
hoiling'point of water which would be produced if there were no dis-
sociation, and knowing the values actually found, we calculate the
amount of dissociation by simple proportion. If the depression of
the freezing-point or rise of tlie boiling-point is twice as great as if
there were no dissociation, the comi>ound is completely dissociated,
since each molecule yields two ion3 if the electrolyte is bimtrjf like
those already considered* If the electrolyte breaks down into three
ions, — ^ia terna}% — ^ these values are three times as large as if there
is no dissociation.

If the values are one and one-half times as large as if there is no
dissociation, it mea.ns that a binary electrolyte is dissociated fifty per
rent, a ternary electrolyte twenty-five per cent, and so on. These
examples will make the principle clear.

The Condnotmty Kethod. — Another method is frequently used
for measuring electrolytic dissociation. We have seen that solutions
of electrolytes conduct the current, and indeed it is this property
which characterizes a given substance as an electrolyte or a non-
electrolyte. Solutions of electrolytes, however, conduct very differ^
ently even when the concentrations are the same. The amount of
the conductivity has, however, been showu to depend upon the degree
of the dissociation, and for any given electrolyte to be proportional
to the number of ions present. In comparing conductivities, how-
ever, we must of course take into account tlie concentration, A
normal solution in physical chemistry means one that contains a gram-
molecular weight of the electrolyte in a litre of solution. Such a
solution conducts better than a tenth-normal solution, but not ten
times as welh To compare the conductivities of these two sola-
tions we must divide that of the former by ten, or multiply that of
the latter by ten. We adopt the second mode of procedure, and
compare the conductivities of all solutions with those of the normal
solution. Such are known as moiecuiar conducUvitieSf since they
always refer to molecular quantities.

In the Kohlrausch method of measuriug conductivity an alternat-
ing current is passed between platinnm electrcxles, thi-ough the soln*
tiou whose conductivity it is desired to study. Tlie resistance of
the solution is balanced against a rheostat on a Wheatst4>ijc bridge,
the point of equilibrium being determined by means of a telephone*

The apparatus used in the method of Kohlrausch is sketched in
Fig* 23» iris a rheostat or set of resistance coils. The meti*e stick

OSMOTIC PRESSURE

118

AB is divided into millimetres, and over this is stretched a manga-
nine wire (manganine being an alloy of German silver and manga-
nese). c7is a small induction coil which furnishes the alternating
current, i? is a glass cup which contains the solution whose resist-
ance is to be measured. The electrodes are cut from thick sheet
platinum, and a piece of platinum wire is welded into the centre of
each plate. This wire is then sealed into a glass tube, which is filled

with mercury to make electrical contact with a copper wire intro-
duced into the mercury. The telephone is connected between the
rheostat and resistance vessel, and also with the bridge wire, by
means of a slider. The point of equilibrium is ascertained by mov-
ing the slider along the wire until the sound of the coil is no longer
audible in the telephone. Let this be a point C. Let us call the
distance AC, a, BC, b, the resistance in the box r, and the resistance
in the vessel ri. From the principle of the Wheatstone bridge we
would have —

rb = Via ;

rb
ri= —
a

Since conductivity c is the reciprocal of the resistance n —

a

c =

rb

This expression does not take into account the concentration of
the solution. In practice it is best to express concentrations in terms
of gram-molecular weights of the electrolytes in a litre (gram-molec-

114

PRINCIPLES OF INORGANIC CHEMISTRY

xilar normal)* As we liave seeiij tlie number of litres of the solution
containing a gram -molecular weight of the electrolyte may be repre-
sented ty Vf when the above expression becomes —

By intrctflucing v into the above expression, we pass from specific
to molecular conduct! vitiesj and we express the molecular condnc-
tirity by the letter ^, In order to indicate the concentration v to
which ^ aj>pliesj we write for the molecolai conductivity /t^ —

''' = Tb

This ejcpression takes into account all of the factors except the
cell constant A% which depends upon the size of the electrodes which
we are using, and their distance apart. Introducing the constant, we
have —

Calculatioti of the Blssociation from Conductivity ITeasiirementi. —

The molecular conductivity of an electrolyte increases with the dilu"
tiou of the solution up to a certain pointy where it acquires a maxi-
mum, coustaut value. This corresponds to the condition of complete
dissociation, and is represented by the symbol /jt^.

When there is no dissociation there is no conductivity. When
there is partial dissociation the value of the molecidar conductivity
is between zero and ^^, These intermediate values of the molecu*
lar conductivity are I'c presented as values of ^^ ; v, representing the
dilution of the solution or volume, is the number of litres of the
solution which contains a gram-molecular weight of the electrolyte.

If we wish to know the percentage of dissociation n, at any dilu-
tion K, it is tmly necessary to divide the value of il^ at that dilution
by the value of ^^*

For details in applying the freezing-point, boiling-i>ointj and con-
ductivity methods to the measurement of electrolytic dissociation^
some of the physical chemical manuals must be consulted.

CHAPTER IX

CHLORINB (At. Wt. = 35.45)

Chlorine an Element or a Compound. — Although chlorine does
not occur in the free state, it was discovered as early as 1774 by the
great Swedish chemist Scheele, who, however, did not recognize its
elementary nature. The question of chlorine being an element or a
compound is closely connected with an interesting chapter in the
history of chemistry. It was thought at one period that oxygen is
essential to acidity. In order that a compound should be an acid
it must contain oxygen. Chlorine combines with hydrogen, forming
one of the strongest acids known to man. The question arose where
does the oxygen come from in the compound of chlorine with hydro-
gen, known at that time as muriatic acid ? It was obvious that it
could not come from the hydrogen, whose elementary nature was
recognized at that time. It must, therefore, come from the chlorine.
Chlorine was then regarded as an oxide of some element which was
unknown, and which they could not isolate. However, it was termed
murium, and chlorine was regarded as the oxide of murium. Hydro-
chloric acid, since it contained this oxide of murium, was known as
" muriatic acid," a name which it bears even to-day. These views
were held about 1785-1790.

The French chemist, Gay-Lussac, however, made it probable by
his investigations that chlorine is an element, and this same conclu-
sion was reached by the Englishman, Humphry Davy, in the early
years of the nineteenth century.

During the last century a number of attempts were made to
decompose chlorine into simpler substances, but all of these have
failed; the evidence all pointing unmistakably to the elementary
nature of chlorine.

Ooonrrence and Preparation of Chlorine. — Chlorine does not
occur in the free condition in nature. This is due in part to its
great chemical activity. If once set free it would quickly combine
again with other substances. It occurs in combination with many
other elements, such as magnesium, potassium, silver, lead, but
especially in combination with the element sodium, as sodium chlo-

115

116 PRDfClPLES OF INORGANIC CHEMISTRY

ride. The clilorides of all the above elements are readily soluble ia
water, except the chlorides of lead and silver. The soluble cddoridea
cannot exist on the surface of the earthy where they are subjected
to the influence of water, but pass luto solution and are swept down
to the sea.' This accounts for the large amounts of chlorine in sear
water, mainly in the form of potassium, magnesium j and sodium
chlorides.

In certain protected localities, however, which are not readily
accessible to water* bs in the great salt beds of the earthy the chlorides
may remain in solid form. As examples, take the great deposits at
Stassfnrt in Germany, Salzburgj and the like. The deposits were
made by the evaj^oration of the seas which once covered these reg-ions,
and wliich contained the various salts in solution.

The more common minerals containing a large amount of chlorine
are ccimaUlte^ srrflvmej rock sait^ and the like.

A nnml>er 4Df methods have been devised for preparing chlorine,
but most of these ai-e now only of historical interest*

The process devised by Deacon consists in oxidizing hydro-
chloric acid by means of the oxygen of the air* Hydrochloric acid
and air are passed througli heated tubes containing balls of clay
saturated with copper sulphate. Under these conditions the oxygen
of the air unites with the hydrogen of the hydrochloric acid, form-
ing water and liberating chlorine*

Another method for obtaining chlorine, based upon the oxidation
of hydrocldoric acid, is the following: ^ —

When a compound rich in oxygen, like manganese dioxide,
MnOj, ia heated with hydrochloric acid, the latter is oxidized to
water and chlorine.

MnO, + 4 HCI ^ MnCl, + 2 HiO + CV

A part of the chlorine combines wnth the manganese, forming man-
ganese chloride, and is lost as far as free chlorine is concerned,
A method ( Wehlon-s) has been devised for obtaining this part of the
chlorine by converting the chloride of manganese into an oxygen
compound, but this is of little importance at present

A very convenient method for obtaining chlorine in the laborar
tory consists in treating yencJiing-powder with hydrochloric acid.
The bleaching-powder is introduced into an ordinary Kipp'S appara-
tus, and the acid allowed to come in contact with it It will be
rememliered from the preparation of hydrogen that this apparatus
ivorks automatically. When no more gas is desired a stoi>cook is
closed^ and the pressure of the gas liberated forces the acid away

CHLORINE 117

from the bleaching-powder. The reaction which takes place here
will be discussed under the element calcium.

All of these methods have practically given place to the electro-
lytic. Most of the chlorine is now prepared by the electrolysis of
aqueous potassium or sodium chloride. Wlien a solution of potassium
chloride in water is electrolyzed, hydrogen separates at the cathode
and chlorine at the anode. The potassium remains in solution
around the cathode as potassium hydroxide. The following equa-
tion represents the reaction which takes place : -^

2KCl-f-2H,0 = 2KOH-f-Hs-f-Cl^

In an analogous manner chlorine is prepared by the electrolysis of
camcUlUe, a double chloride of potassium and magnesium having
the composition KMgCla.

As already indicated, the electrolytic method has practically
replaced all others for preparing chlorine on the large scale.

Chemical Properties of Chlorine. — The yellowish-green gas chlo-
rine is, chemically, one of the most active substances known. It
combines with nearly all the elements and with many compounds by
simple contact, often with evolution of much heat and even light,
and in some cases almost with explosive violence. The best method
of collecting chlorine for experimental purposes is by displacement
of air. Being heavier than air the chlorine gas is conducted to
the bottom of the vessel containing air, and the latter is displaced
upward by the heavier chlorine.

When copper foil is brought in contact with chlorine gas it com-
bines with the chlorine, shown by the fact that it glows and forms
chloride of copper.

2Cu-f-Cla = 2CuCl,

or, Cu-f-Cls = CuCl2.

When finely divided antimony is allowed to fall into a vessel
containing chlorine gas, we have literally a rain of fire — each anti-
mony particle becoming incandescent as it combines with the
chlorine.

2Sb4-3Cl8 = 2SbCl8.

Phosphorus, boron, silicon, and other elements readily bum in
chlorine, forming the corresponding chlorides. Other substances,
like brass, burn in chlorine only when they have been heated to an
elevated temperature.

Combustion in Chlorine. — We have here examples of combination
taking place between substances and chlorine, which are analogous

118

PRINCIPLES OF INORGANIC CHEMISTRY

to combustion in oxygen. In the former case, as in tlie latter, the
combination takea place with evolution of light and heat^ and tbe com-
bustion in cbloiuie is even more energetic than in oxygen, in that it
starts at ordinary temperatures. We have, then^ combustion in chlorine
just as truly as in oxygen. The tenn coujbiistion, however, as ordi-
narily used, always refers to conibinatiun with oxygen, since we never
know chlorine in the free condition unless it is specially prepared.

Action of Gblorine on Hydrogen. — Hydrogeu unites with chlorine
at ordinary temperatures if exposed to diffuse light, and with explo-
sive violence if exposed to direct sunlight A jet of hydrogen can,
however, be burned in chlorine just as it can be burned in oxygen.
When hydrogen was bumed in oxygen, the two gases combined,
fonning water. When hydrogen is burned in chlorine, the two gases
c^uubine, forming the important compound hydrochloric acid, which
we shall study a little later. Introduce a jot of burning hydrogen
into a vessel filled with chlorine, and notice the pale green color of tbe
flame* alno the fumes of hydrochloric acid formed. A jet of chlorine
al^o burns rt*iulily when plunged into a vessel filled with hydrogen gas.

Action of Chlorine on Water, — Chlorine is readily soluble in
WAter, and tb« rtmultiug solntion is known as chlorine water. Chlo-
rine vmtrr, if kept in the dark, is a stable substance, but if exposed
to the lights a deep-seated change takes place. The chlorine acta
fhenjically np<m the water, combining with the hydrogen and liber-
atiu^c o?cygen* The resulting solution contains the hydrochloric
noUl f^irmed, while the oxygen gas is liberdted. Such chemical
rt^ai^tion^ which are brought al>out by the action of light are known
an ithiditchfrnimi tf actions. We shall encounter a number of tbem as
our «ubjtM*t dini*l(>p»* Since oxygen is liberated, chlorine is known
Iki a dnmij itA'iditiH(j arjenL Its oxidizing power renders chlorine one
of %lw vt*ry bent hlnKhing agents which is at our disposal. The
uvyt^'U winch \n si^t frt*e when chlorine acts on moisture oxidizes
i*rtjnnui eoloiiug-mutlt'r, iukI leaves iKduml the colorless substance.
Tbis K^\y b«» ilUi!itnittH.l by bringing into the presence of chlorine gas
liHiK^ luvH^t tiowi^ra oi- a moist piece of calico, when tbe color will dis-
ap|H*ar iu a wsxs short tinu\

Chloiino i» al»o an excellent disinfectant ^ having an unusual
l^iwer U* di^stioy kuitf^ria and other forms of life. This is due in
piiit Ui iU MXiduiuk* lU^iiJii, and in part to direct combination of
i^likii'iui^ wjih the organic nmtter of snch forms of life. All things
^WttAUlurod, chluvine \s one of tbe most ]iowerfijl disinfectants known*

4ot4ati ol Chloriiit on Certain Organic Compoimdi, — Chlorine not
only mU ou t^lenieuUry substances and simple compounds, but also

CHLORINE 119

on complex organic substances. When brought in contact with the
elements it combines with them in the proportions represented by
the above equations. When brought in contact with organic sub-
stances which contain hydrogen, the chlorine first replaces the
hydrogen, taking its place in the molecule, and then more chlorine
combines with the replaced hydrogen, forming hydrochloric acid.
Such a reaction, known as substitution, can be illustrated by bring-
ing a piece of filter-paper saturated with oil of turpentine into a
vessel filled with chlorine gas. A violent reaction takes place,
resulting in the liberation of a large amount of finely divided carbon.

Oil of turpentine consists of carbon and hydrogen. The chlorine
drives out the hydrogen in part at least, combining with it and form-
ing hydrochloric acid, and leaves the carbon behind in the finely
divided condition.

A better example of the substituting action of chlorine is its
action on the compound benzene. Benzene also contains carbon and
hydrogen, having the composition expressed by the formula C^H^
Chlorine displaces the hydrogen atoms, combining with them and
also taking their place in the molecule : —

CeHe-f-Cl2 = CeH,Cl-f-HCl.

This process can be continued until all the hydrogen has been
replaced by chlorine, the resulting compound being CeCle.

Chlorine Hydrate. — When chlorine gas is conducted into a mix-
ture of water and ice, a crystalline compound separates, having a
greenish color, and the composition is represented by the formula
Cl,.8 HjO or Clj.lO H^O. At ordinary temperatures it decomposes
into chlorine and water, while at somewhat elevated temperatures
the decomposition is quite rapid, resulting in a copious evolution of
chlorine gas.

This compound is of
special historical interest
in connection with the
liquefaction of chlorine,
and also in connection
with the liquefaction of
gases in general. The
earlier work on the lique- ^^' *

faction of gases was carried out almost exclusively by the great
English physicist, Faraday. He succeeded in liquefying chlorine by
means of chlorine hydrate. Some of this compound was placed in
one end of a thick-walled glass tube, and the other end closed, as

ISO

PRINCIPLES OF INORGANIC CHEMISTRY

shown in Fig, 24. The end of the tube containing the chlorine
hydrate was gently warmed, while the other end waa surrounded by
a freezing-mixture of ice and salt Under these conditions the
chlorine was liberated, and produced a pressure in the tube which
was sufficient to liquefy it This is one of the classical experi-
menta of Faraday on the liquefaction of gases.

PHYSICAL PROPERTIES OF CHLORi:o:

Certain Physical Properties of Chlorine. — The yellowish-green
gas, chlorine, u alxjut two and one-half times as heavy as the air, a
litre weighing 3.22 grams. It hns a most ctisagreeable odor, ami an
injurious effect when inhaled. It a<;^ts upon the mucous membrane
of the nose and throat, and disintegrates these tissues if inhaled in
sufficient quantity and for sufficient time. It is therefore necessary
in working with chlorine to take every precaution to be protected
from the gas. Such work should always be don t? under a good hood,
with a strong draft to remove the gas as rapidly as it eocapes into
the air. Even when all ordinary precautions are taken, enough of
the gas escapes into the atmosphere to be very unpleasant to the
experimenter^ and to produce uncomfortable results if inhaled for
a sufficient time.

Chlorine does not obey the laws of Boyle or Gay-Lnssac with the
same exactness as oxygen and hydrogen. This is probably connected
with the fact that chlorine at onlinary temperatures is below its
critical temperature, and not far above its liquefaction temperature
under atmospheric pressure. It is a general rule that gases near
their point of liquefaction do not obey the gasdawa.

Liquefaction of ChloriiLe. — That chlorine can be liquefied with
comparative readiness has been shown by Faraday's experiment, by
means of wliich liquid chlorine was obtained from chlorine hydrate.
At zero degrees a pressure of six atmospheres are ret^uired to liquefy
chlorine, while under a pressure of one atmosphere it passes over
into a liquid at — 33^G. This is the boiling-iwint of liquid chlorine.
Its critical tcmperaturej m shown by the experiment of Faraday, is
above the temperature of a mixture of salt and ice, and, indeed, is
quite high. It is 146'', and the pressure required to liquefy chlorine
at this temperature is 94 atmospheres, this being the critical pressure
of chlorine. Liquid chlorine has a specific gravity of 1.66 at — S0%
but the coefficient of expansion is very large, and at higher teinperar
tures the specific gravity is much less.

Liquid chlorine has a yellow color, showing little or none of the

CHLORINE 121

green which is characteristic of the gas. It freezes at — 102®, form-
ing a greenish-yellow solid.

Comparative Inactiyity of Dry Chlorine. — While moist chlorine
is one of the most active substances chemically, dry chlorine is com-
paratively inactive. A lecture-table experiment, which is frequently
shown, is to pass chlorine through a glass tube containing a piece of
metallic sodium which is heated by means of a Bunsen burner. If
the chlorine has been carefully dried, as is frequently done, the sodium
will melt and remain with untarnished surface in contact with the
chlorine gas. If, on the other hand, the water-vapor has not been
removed from the chlorine, vigorous chemical action will take place,
accompanied with intense heat and a bright light, and the compound
sodium chloride will be formed.

The comparative inactivity of dry chlorine is further shown by
the fact that liquid chlorine can be kept and transported in strong
steel cylinders. Indeed, it can be obtained on the market in this
form, and is the most convenient means of obtaining chlorine. The
steel cylinders are provided with a stop-cock, so that when the gas is
desired it is only necessary to open the stop-cock and obtain it.

We have already studied one reaction which would take place
only when water was present — the union of hydrogen and oxygen.
We shall meet later with a number of similar examples.

HYDROCHLORIC ACID

Hydrochloric Acid, HCL — We have already examined a number
of reactions in which hydrochloric acid was formed. We shall study
more closely some of these, and other reactions in which hydrochloric
acid is produced. We have seen that when hydrogen is burned in
chlorine the product is hydrochloric acid. In order that this re-
action should take place it is not necessary that an ignited jet of
hydrogen should be introduced into the chlorine. We have seen that
when the mixture is exposed to diffuse light a gradual combination
takes place, and when exposed to direct sunlight the gases combine
with explosive violence.

When hydrogen and chlorine gases are mixed in equal volumes,
and an electric spark passed through the mixture, the gases combine
with explosive violence. Such a mixture is known as chhr-electrolytic
gaSf or chlorine detonating gas, and is readily obtained by electrolizing
a concentrated aqueous solution of hydrochloric acid.

Volume Belations in which Hydrogen and Chlorine Combine. —
We have studied the relations by volume in which hydrogen and

122

PBIXCIPLES OF mORGAXIC CHEMISTltr

oxygen combine, aod the ratio between t!ie volumes of the gases
which enter into conibination and the volume of the prcxiuct formed.
It will be remembered that one volume of oxygen combines with
two volumes of hydrogen, and forms two volumes of water-vapor,
The relations which obtain in the case of hydrogen and chlorine are
even simpler. ^Tien one volume of hydrogen is mixed with one
volume of chlorine and combination takes place^ all of both gases are
used up, and just two volumes of hydrochloric acid are formed. The
law of the simple volume relations in which gases combine holds
here even more strikingly than in the case of oxygen and hydrogen,
there being no contraction in volume when the gases hydrogen and
chlorine combine; and further, these gases combine in the simplest
ratio -by volume, viz. equality.

Preparation of Hydrochloric Acid* — Hydrochloric acid gas is
prepared most conveniently on a large scale by a methoil entirely
different fi-ora any of the above. When a salt of hydrochloric acid
is treated with a n on- volatile acid such as sulphuric, the hydrochloric
acid gas is set f i-ee. The best known salt of hydrochloric acid is, as
we have seen, sodium chloride. When this is treated with sulphuric
acid, a reaction takes place in the sense of the following equation ; —

2 NaCl + H,S04 = Na^SO^ + 2 HCL

The hydrochloric acid gas thus formed is conducted into water,
which has the power of absorbing large quantities of it. This is the
form in which it is used in the laboratory and in the arts. When it
is desired to obtain the gas again from its concentrated solution in
water, it is only necessary to add something to the solution which
has greater attraction for water than the hydrochloric acid. Such a
substance is ordinary sulphuric acdd. ^Yhen concentrated sulphuric
acid is di*opped slowly into concentrated, aqueous hydrochloric aeid>
the latter escapes from the solution as a continuous stream of gas.

Another method of preparing hydrochloric acid in quantity is by
heating the chlorides of certain metals with water-vapor. The chlo-
ride of mugnestum is frequently used, the reaction taking place in
the sense of the following equation ; —

MgCl, + H/) = MgO 4- 2 HCL

Chemical Properties of HydrocMoric Acid- — Hydrochloric acid,
as the name implies, is an acid, and since this is the first stibstance
which we have thus far encountered with acid properties, a few words
should be added in reference to acids in general. Hydrochloric acid
has the composition represented by the formula HCL Its molecule

CHLORINE 128

therefore contains one atom of hydrogen and one of chlorine. The
question arises to which constituent are the acid properties due ? It
may be due to either or to both. When we come to study other acids
we shall learn that many substances are acids which do not contain
any chlorine, and many compounds containing chlorine are not acids.
Therefore chlorine is not essential to acidity. We shall also learn
that all substances which are acid contain hydrogen, and no other
element in common. Hydrogen is therefore essential to acidity.

There are, however, many compounds which contain hydrogen
and which are not acids. The question which arises is how does
the hydrogen in the latter class of compounds differ from the hydro-
gen in the former? The answer is furnished by the theory of
electrolytic dissociation. When the compound hydrochloric acid is
dissolved in water, its solution conducts the electric current. It is,
therefore, an electrolyte, and its molecules are dissociated to a greater
or less extent into ions. An aqueous solution of hydrochloric acid
is, then, a solution of hydrogen and of chlorine ions.

As we shall have to deal frequently with ions, we adopt some
method of distinguishing between atoms and ions. Since ions are
charged atoms or groups of atoms, we shall use the positive sign
over the symbol of an atom to mean that it. is charged positively,
or is a cation. The negative sign over an atom or group of atoms
means that it is carrying a negative charge and is an anion. Hydro-
chloric acid is dissociated by water in the sense of the following
equation : —

HCI = H,C1.

When a solution of hydrochloric acid is brought in contact with
a metal like zinc, the latter takes the positive charge from the
hydrogen ion, becoming itself an ion and passing into solution, while
the hydrogen ion having lost its electrical charge becomes an atom.
We have seen, however, that an atom of hydrogen cannot exist by
itself, two atoms combining and forming a molecule of hydrogen.
The reaction between zinc and hydrochloric acid is represented by
the following equation : —

Zn 4- H, Ci 4- H, Ci = Zn, CI, CI + H,.

Hydrochloric acid acts upon metals in general in the sense of the
above equation — the metal taking the charge from the hydrogen ion,
becoming itself an ion, converting the hydrogen into the atomic con*
dition. A few examples will make this clear.

124 PEINCirLES OF INOHOANIC CHEMISTKY

K+H, Cl = Kpd + H|

Ca + M, Cl + H, (Jl = Ca, (Ji, Cl + H,;

Fe + H, CI + H, Ci + H, ci='Fe, CI, CI, C1 + 3H.

Hydrochlovie acid has also the jiower of acting eliemically upon
HubataiJi^es oth^r tkiti the metals. Indeed^ the action of hydrochlo-
ric afid u(>oii metals is not simply a chemical act, since it consists
i'hiefly in the tratisfer of an electrical charge from the hydrogen ion
of the acid to the metal.

Take a aubstance like calcinm hydroxide ^ ordinary lime-water —
having tlie composition Ca(OH)^ When this is treated with hydro-
cldoric acid a reaction takes plaee which we at present represent by
the following equation, and will study it later in more detail; —

Ca(OH)s 4- H. CI + H, CI = C^ Cl, CI + 2 11,0.

('alcinm hydroxide is a type of substances known as bases^ with
whioh we shall become familiar a little later. Hydrochloric acid
atHs in general itpou bases.

BefinitiOD of an Acid- — Having studied hydrochloric acid as
ihii ty[H* of a large class of chemical c (mi pounds known as acids,
wit are prt^pared to consider these a little more closely. The old
iHUJceptiun was that acid properties depend for their existence upon
thti presence of oxygen. Indeed, the terra oxij^en means acid-
former. This had to be almndoned after it was shown that many of
our atrojigeat acids contain no oxygen whatsoever.

As iatlicated above, all acids hare certain properties in cam-
inon* They all taiite sour; they have the property of coloring
eertsdn vegetable dyes red. They have the power of dissolving cer-
tain metala in the sense of the above equations, and they all
eontaiu hydrogen wliieh can give up its electrical charge to certain
metals, itself escai*ing as hydrogen gas.

Hydrogen in this form is known as ionic hydrogen, and, as
has been utated, wlierever we have ionic hydrogen we have acid
propertii's, and wherever we have acid proiK?rtie3 we have ionic
hydrogen. To say tlxat a compound ha,** acid properties n jeans, then,
that ivhen it tit dtmolvt^d in water or mma other dissociating ml vent it
yit*ldii htjdrtitft'n ion»,

Thia dt^tinition says that a compound is not an acid unless it is
bnmght into the i*resenco of a dissociating solvmit. This is the
same as to say that no pure» homogeneous substance is au acid.
This seems on the face of it like going too far. Can we think
of pure, dry hydrochloric acid, for example, as nut having acid

CHLORINE 125

properties ? The definition goes still farther, and says that in order
that a compound should have acid properties it must be dissolved in
a dissociating solvent. If the definition is true, when a substance
like hydrochloric acid gas is dissolved in a non-dissociating solvent,
it should have no acid properties. To any one who is familiar with
the strongly acid properties of hydrochloric acid when dissolved in
water, this definition seems to lead to pretty serious consequences.
What are the facts ? Pure, dry, liquid hydrochloric acid has no
acid properties. When pure, dry hydrochloric acid gas is dissolved
in pure, dry, non-dissociating solvents like chloroform or benzene,
the solutions have not the least trace of acid properties. A solution
of dry hydrochloric acid gas in dry benzene will not even color
blue litmus red. It should be added that such solutions do not
conduct the electric current, showing that there are no ions, and,
therefore, no hydrogen ions present. They have not the slightest
power to dissolve metals, showing again that there are no hydrogen
ions present to give up their charges to the metal atoms.

Such solutions would, however, taste sour, since as quickly as the
molecules of the acid in the chloroform or benzene are brought in
contact with the tongue, they are also in contact with moisture, and
would be dissociated at once into hydrogen ions and chlorine ions.
The hydrogen ions produce the characteristic sour taste of acids.

W^e shall learn that the above relations hold for all acids. No
pure, dry, homogeneous substance is an acid. It becomes an acid
only when dissociated by a solvent or some other means into
hydrogen cations, and into anions whose nature depends upon the
acid in question and varies with every acid.

Detection of Hydrochloric Acid. — There is one reaction which
serves to detect hydrochloric acid under all ordinary conditions.
Hydrochloric acid is, as we have seen, dissociated into hydrogen
ions and chlorine ions. Any reaction which would detect hydrogen
ions would not be a characteristic reaction of hydrochloric acid,
since all acids when dissolved in water yield hydrogen ions.

To detect hydrochloric acid, then, we must make use of some
reaction which is characteristic of the chlorine ion since all acids
yield hydrogen ions. Such a reaction takes place whenever a silver
ion is brought in contact with a chlorine ion.

H, CI + A'g, NOs = AgCl -f- H, NO3

expresses the reaction between hydrochloric acid and silver nitrate.
The silver chloride formed is a white solid, readily soluble in ammonia.
Since soluble chlorides in general yield in solution chlorine ions,
this is a means of detecting also the presence of such chlorides.

126

PRIXCrPLES OF INORGANIC CHEMrSTRr

Fliysical Properties of HydrocMorio Acid. — Hydrochloric acid
is a colorless gas, wtiU a sharp, pungent odor, and produces marked
irritation of the tn neons niembrane of the nose and throat when
inhaled even in small quantity. Its critical temperature is 52°, so
that it can be readily liquefied. Its critical pressure is 82 atmos-
pheres. At lower temperatures it can be liquefied at much lower
pressures. At zero degrees it can be liquefied by a pressure of about
thirty atmospheres.

Liquid hydrochloric acid is colorless, and as already indicated, is
much less active chemically than the solution in water. Indeed, it is
a comparatively inactive substance. When carefully freed from water
it does not act on metals, and does not even color blue litmus red.

The liquid boils under atmospheric pressure at — SO^'.S, and solidi-
fies at — 112^5, The gas shows unusual solubility in water, one
volume of water at zero degi-ees dissolving al)out BOS volumes of the
gas. The solubility dimitushes as the temperature rises, whicdi is in
keeping with tlie general rule for the solubility of a gas in a liquid.
The gas has such great attraction for water that if the breath is
blown across the open mouth of a bottle of concentrated hydrochloric
acid, the particles of water are condensed around the escaping hydro-
chloric acid, and a mist is produced which can be readily seen* The
same eifect is observed when the breath is blown into a stream of
hydrochloric acid gas escaping from a generator.

Aqueous Solution of Hydrochloric Acid. — Hydrochloric acid gas
dissolved in wati/r is not a trui? solution of a gas in a liquid. That
this is the ease is shawn in several ways. When hytlrixdiloric acid
gBA is dissolved in water there is a marked evolution of heat, which
does not take place when a gas is simply dissolved in a liquid. This
would indicate that there is chemical union between the acid and
water.

Further, when a solution of the gas in water is boiled even under
diminished pressure, it is not possible to remove all the gas from the
water, hut a considerable portion remains dissolved in the water.
When boiled under a pressure of 760 mm. gas escapes until the
remaining liquid has a composition corresjwnding approximately
to one molecule of hydrochloric acid and eight molecules of water.
This mixture lx>ila at a fairly constant temperature (110°)» and does
not change in composition, the distilled portion having the same com-
position as the undistilled liquid which remains behind. Further,
if the aqueous solution of hydrochloric acid is more dilute than
would correspond to this composition^ water distils over until this
composition is reached* AH of these facts would indicate that this

CHLORINE 127

particular mixture of hydrochloric acid and water is a definite chemical
compound. It is well known that a chemical compound has a definite
boiling-point, which is a characteristic constant of the substance.

There is one fact, however, which shows that this substance with
a specific gravity of 1.102, and containing 20.3 per cent of hydro-
chloric acid is not a chemical compound. Its composition changes
as we change the pressure under which it is boiled. If the press-
ure is greater than 7G0 mm., the distillate, or portion which distils
over, contains more water in proportion to acid ; while if the press-
ure is lower than the normal, the distillate is richer in acid than
would correspond to one of acid to eight of water. This alone shows
that the substance is not a chemical compound. Its constant com-
position when boiled under constant pressure is satisfactorily ex-
plained by physical chemistry, but it would lead us too far to dis-
cuss the subject in full in this connection. Suffice it to say that
there are many such constant boiling mixtures known, none of
which, however, are definite chemical compounds.

There is, however, a definite compound of hydrochloric acid and
water which is well known. When hydrochloric acid gas is con-
ducted into a concentrated aqueous solution of hydrochloric acid
which has been cooled to —22®, well-defined crystals separate,
having the composition HCI.2H2O. These melt when heated to
— 18®, and decompose at higher temperatures.

When hydrochloric acid gas is dissolving in water, it is advisable
to take one special precaution. The gas is so very soluble in water
that if the tube through which the gas is escaping is plunged far
beneath the surface of the water, the liquid is liable to rise rapidly
in the tube and fiow back into the generating fiask.

An arrangement by means of which this can be avoided is the
following : —

The generating flask is provided with an exit tube through which
the gas escapes when the sulphuric acid acts on the sodium chlo-
ride. The end of this tube is provided with a funnel which dips just
beneath the water in the receiver. If the water should tend to rise
in the tube, due to the rapid absorption of the gas and the pro-
duction of a partial vacuum in the tube, it will rise in the funnel
until air can enter and restore the pressure to the normaL In this
way the water is prevented from flowing back into the flask and break-
ing it while hot, or causing an explosion by contact with the sulphuric
acid, which has become hot as the result of the chemical action.

This precaution is always taken when a very soluble gas is dis-
solved in a liquid.

128 PRINX'IPLES OF IXORGANIC CHEMISTRY

COMPOUNDS OF CHLORINE WITH OXYGEN AND HYDROGEN

Compoiinds of Chlorine with Oxygen. — Although chlorine and
oxygen cannot be made to combine directly, several compounds of
these two elements have been made by indirect methods. These
compounds are chlorine monoxide (Cy3), chlorine dioxide (010,),
and chlorine septoxide (Cl^tO?). These compounds are all character-
ized by instability. They are prepared by methods with which we
shall become familiar in our study of the compounds of chlorine
with oxygen and hydrogen, and we shall therefore turn to this class
of substances.

Compounds of Chlorine with Oxygen and Hydrogen. — Thus far
we have studied compounds between only two elements. We might
suspect, therefore, that only two chemical elements can combine with
one another, forming a definite molecule, or at least that this is by far
the most common form of chemical union. Such is by no means the
case. We shall now study briefly a class of compounds between
the three elements, chlorine, oxygen, and hydrogen, which, if not
very common substances, have considerable chemical interest. These
compounds are : —

Hypochloroiis acid HCIO.

Chlorous acid HCIO,.

Chloric acid HClOji.

Perchloric acid HCIO4.

Hypochlorout Acid, HOCL — Hypochlorous acid is formed in very
small quantity when chlorine acts upon water, and in the sense of
the following equation : —

H,0 4-Cl,= HCl-t-H0Cl.

The free hydrochloric kcid formed also as the result of the reaction
acts upon the hypochlorous acid and decomposes it. If there is some
substance present to combine with the hydrochloric acid, such as
freshly i)recipitated mercuric oxide, it is removed from the field of
action and does not decompose the hypochlorous acid. The reaction
between chlorine and water, then, takes place in the sense of the
following equation: —

U,0 4- 2 CI, -f HgO = HgCl, -f 2 HCIO.

Another method for preparing hypochlorous acid, which on the
whole is the best, consists in preparing first the potassium or sodium
salt of the acid. When chlorine is conducted into a cold, dilute

CHLORINE 129

solution of potassium or sodium hydroxide, the reaction takes place
in the sense of the following equations : —

2 KOH + CI, = KCl -f KOCl -f H,0 ;
2 NaOH -f Cla = NaCl + NaOCl + H,0.

The salts of hypochlorous acid are termed hypochlorites. These
salts, therefore, are potassium hypochlorite and sodium hypochlorite.
When either of these salts is treated with a cold, dilute solution
of hydrochloric acid, the hypochlorous acid is set free : —

KOCl 4- HCl = KCl + HOCl.

The hypochlorous acid is then distilled off and collected.

Properties of Hypochlorous Acid. — Hypochlorous acid is a weak
acid, and is an unstable compound. It readily gives up oxygen,
passing over into hydrochloric acid : —

2HC10 = 2HCl4-Oj^

It is, therefore, a powerful oxidizing agent, and if brought in contact
with substances which can take up oxygen, readily gives it up to
them. Its value as a bleaching ageiU depends upon this fact

Calcium Hypochlorite, Ca(OCl)«. — The calcium salt of hypochlo-
rous acid is used extensively as a bleaching agent, on account of the
ease with which it gives up chlorine. When chlorine is passed into
slaked lime, the following reaction apparently takes place : —

2 Ca(0H)2 + 2 Clj = CaCl, + Ca(OCl), + 2 H,0.

This apparent mixture of calcium chloride and calcium hypochlorite
is known as " bleaching-powder," and is largely used as a disinfectant.
Chlorine Monoxide, ClgO. — Hypochlorous acid is not known in
the anhydrous condition. When an attempt is made to free it from
water, it loses water itself and passes over into chlorine monoxide.

2 HOCl = HjO + ClsO.

The compound CljO, since it is formed from hypochlorous acid
by the removal of water, is also known as hypochlorous anhydride ;
the term anhydride of a substance being a generic term for a com-
pound derived from another by loss of water. The most convenient
method of preparing chlorine monoxide is by passing dry chlorine
over dry, yellow, mercuric oxide.

HgO 4- 2 CU = HgCl, + Cl^O.
The chlorine monoxide, being a gas at ordinary temperatures, can be
readily collected. It is readily converted into a dark-yellow liquid,
which boils at — 19®.

ISO

PRTXCrPLES OF IKOHGANTC CnEMrSTRY

Gafleous chlorine monoxide is somewhat explosive, but the liquid
is very explosive. By wanning or jarring, it explodes easily, yield-
ing clilorine and oxygen.

2Cl,0 = 2Cl, + 0^

Chlorine monoxide dissolves readily in water, combining with it
and forming hypochloroiis acid.

Cl,0-hHjO = 2HC10. 1

Chlaric Acid, HCIO,. — The most convenient method of preparing
chloric acid is first to prepare the porassium or barium salt, and from
the salt to obtain the free acid. When cldorine gas i& conducted
into a hot, concentrated solution of potassium hydroxide, the follow-
ing reaction takes place : —

6 KOH H- 3 CI, = 5 KCl + KClOj + 3 HA

The solution contains, after the reaction is over, two salts, potas-
sium chloride and potassium chlorate. These can, however, be
readily separated by their diflferent solubilities in water; potassium
chloride being quite soluble, while .potassium cldorate is veiy much
less soluble.

From the solution potassium chlorate readily crystallizes, espe-
cially on evaporation, leaving 1>ehind in solution the potassium chlo-
ride. With potassium chlorate we have already become somewhat
familiar when we were studying methods of preparing oxygen. It
will be remembered that this compound gives off all of Its oxygen
when heated to an elevated temj>erature.

When potassium chlorate is treated with a dilute solution of sul-
phur ic acid, the following reaction takes place i —

2 KCIO3 + H,80, = KjSO* -h 2 HClOa.

Care must be taken not to treat potassium chlorate with concen-
trated sulphuric acid, since violent explosions almost always result
from such a reaction.

The solution contains the chloric acid, but since the latter cannot
be distilled without undergoing decomposition, this method does not
yield pure chloric acid.

To obtain pure chloric acid a salt of this acid must be used which
will form an insoluble precipitate with the sulphuric acid. The
barium salt is the most convenient. When barium chlorate is
treated with a dilute solution of sulphuric acid in equivalent quan-

CHLORINE 131

tity, insoluble barium sulphate is precipitated, and pure chloric acid
remains in solution.

(Insoluble)

BaCClOs), + H,S04 = BaS04 +' 2 HClQs.

The barium sulphate is then filtered off, or the clear, supernatant
liquid decanted from the precipitate and concentrated in a vacuum
or over sulphuric acid.

Chloric acid may also be prepared by the action of hydrochloric
acid on silver chlorate, the silver chloride formed being insoluble.

Propertiet of Chloric Acid. — Chloric acid is a colorless liquid with
very strongly acid properties and with great oxidizing power. It
contains a large amount of oxygen, which it readily gives up. When
chloric acid is warmed or exposed to the light, it passes over into
perchloric acid, which we shall study a little later. When a piece of
paper is saturated with a concentrated solution of chloric acid, it is
oxidized so energetically that it bursts into flame.

Chlorates. — The chlorates, as the salts of chloric acid are called,
are very energetic oxidizing agents. This is due in part to the large
amount of oxygen which they contain, and in part to the ease with
which they give it up. To illustrate the unusually strong oxidizing
power of the chlorates and their decomposition products, perform
the following experiment: Mix some finely powdered potassium
chlorate with some finely powdered cane sugar, and place a small
amount of the mixture on a stone slab under the hood. Add cau-
tiously from a pipette a few drops of concentrated sulphuric acid.
In a few moments the entire mass will burst into violent flame.

The Chlorine Ion and the Ion of Chlorates. — We have studied one of
the characteristic properties of the chlorine ion, viz. its power to com-
bine with the silver ion and form insoluble silver chloride. We would
naturally ask whether the chlorine in potassium chlorate has this
same power. When a solution of potassium chlorate is electrolyzed,

the potassium ion moves to the cathode and the chloric ion ClOg to
the anode. Potassium chlorate, therefore, dissociates as follows : —

KC10, = k, cfOs.

Chlorine in this case, instead of forming the anion, forms only a part
of the anion. It is in combination with three oxygen atoms, and the
chlorine and oxygen form the anion. If a solution of potassium
chlorate is treated with a solution of silver nitrate, no precipitate is
formed, showing that chlorine has very different properties when
alone in the ionic state, than when combined with another element
forming part of a complex ion.

132

PRIKGIPLES OF IKORGANIC CHEMISTRY

Perchloric Acid, HCIO^. — When potassium chlorate is heated
vigorously it gives off all of its oxygen, as we saw when we were
studying methods for the preparation of oxygen. If, however,
potassium chlorate is heated moderately, it gives off only a part of
its oxygen. The decomposition of potassium chlorate by heat takes
place^ then, in two stages. In the first place tlie potassium chlorate
melts and gives off oxygen. If the temperature is now kept eon-
stantj OKjgen will cease to come off after a time and the melted sub-
stance will solidify. This solid mass is a mixtvire of potassium
chloride and potassium perchlorate, and this reaction is in general
formulated thus : —

2 KCIO, = KCl + KCIO, + 0^

It is, however, questionable whether this expresses the whole truth.
The perchlorate, KCIO4, can be obtained from the mixture by dissolv-
ing out the potassium chloride by means of cold water ia which
potassium chlorate is very slightly soluble*

AVhen potassium perchlorate is treated with sulphuric acid, per-
chloric acid is set free.

2 KCIO, + H^O^ = K^S04 -h 2 HCIO4.

When the solution is distilled under diminished pressure, perchloric
acid fmsses over and is condensed. By fractional distillation under
diuunished pressure it can be obtained in pure condition,

Fropertiei of Perchloric Acid. — Perchloric acid is more stable
tlmil any of the other oxygen and hydrogen compounds of chlorine.
The pure acid is a very vigorous oxidizing agents and explodes
^^i\y sfhmx brought in contact with substances which can be 0x1-
dizv<t T\w acid containing from thirty to forty per cent of water
i», howi*vet| quite stable* Perchloric acid forms a definite, crystal*
Uu(? ixuuiiivuud with water, having the composition HC10|.HaO.

IVfcvhlmv acid is a very strong acid, which is the same as to
i«\v »hAl it IS vi*ry mucli dissociated by water* It readily replaces
hvdn»i^hK»rio at* id from it-s salts, but this is partly due to the com-
|HU'utiVt» iu.^vdubility vt tht^ pereblor*ates ; it Ijeing a general law in
oh^Hii-Hirv that u^^u an inmlMe comjmnnd can be formed it isforrtiecL
If .| r.*, = u . nnciHilmt^'d solution of potassium chloride is treated
Hith u^ aeki uiH^dU^s of potassium perchlorate are precipi-

UkH\i 1 ' I i» ouo oC the few substances which form

tlitlk^uUl} - , .. h^huhU with potassium and similar elements,
iittd thit ii»4i<?tiu4i ^Wh thcjwtam \m used to detect the presence of
IMii^klukiio mid.

CHLORINE 133

Chlorine Septozide, Clfij. — When perchloric acid is dried with
the powerful dehydrating agent, phosphorus pentoxide (PjO,), water
is removed and chlorine septoxide is formed.

2 HCIO4 + P2O5 = P A.H,0 4- CIA.

Chlorine septoxide is a colorless oil, boiling without decomposition
at 82^

Chlorine Dioxide, ClOj, and Chlorous Acid, HClOg. — Chlorine
combines with hydrogen and oxygen, forming an acid with more oxy-
gen than hypochlorous, and with less oxygen than chloric acid. This
acid was not taken up until we had studied chloric acid and the
chlorates, since, in its preparation, a chlorate is used. When potas-
sium chlorate is carefully heated with oxalic acid, at a temperature of
about 70°, and the decomposition products passed through a freezing
mixture of salt and ice, a reddish oil which is explosive condenses.
This is chlorine dioxide, having the composition CIO,. The same
gas is formed when a chlorate is decomposed with sulphuric acid.
The liquid boils at 10**, and passes over into a yellowish-red
crystalline solid at —79**.

The gas and liquid are explosive. Indeed, it is this gas which
explodes when potassium chlorate is treated with sulphuric acid.
This gas dissolves in water without rendering the solution acid.
When conducted into a solution of a strong alkali, such as potassium
or sodium hydroxide, the following reaction takes place : —

2 CIO2 + 2 KOH = H,0 + KClOs -t- KCIO^

The new compound, KCIO2, which we have met with for the first
time, is known as potassium chlorite, and is a salt of the acid,
HClOj, chlorous acidj which has thus far not been isolated.

Power of Chlorine to combine with Oxygen. — We have thus
seen that there are a number of compounds of chlorine with oxygen,
containing for one atom of chlorine very different amounts of oxygen.
This shows that one element may combine with very different
amounts of another element, and form definite chemical compounds.
This raises the question. What is the true combining weight of
chlorine ? Are we to take the largest or the smallest amount of
chlorine which enters into a molecule? We always take the
smallest, and this is for us the atom. If we take the amount of
chlorine which enters into the molecule, which we write CljO, we find
that it is just double the amount which combines with the element
hydrogen to form hydrochloric acid. The amount of chlorine in a

134

PRIKCrpLES OF IKORGASTC CHEMISTRY

molecule of hydrochloric acid is the smallest quantity of chlorine
known ; chlorine never eomhining with any element in a quantity
which is smaller than that which enters into a molecule of this acid.
This is for us the atom of chlorine, and all other quantities of chlo-
rine are referred to this as the unit. The weight of this amount of
chlorine is i35.45, on the has is of oxygen = IG.

Yalence. — One atom of chlorine sometimes combines with a
half atom of oxygen, i.e. two atoms of chlorine are required to
combine with one atom of oxygen, a^ in the compound CljO. In other
cases, one atom of chlorine combines with two atoms of oxygen, as tu
ClOj, while in the compound CltO^, one atom of chlorine combiDCs
with three and a half atoms of oxygen. The power of an atom to
hold other atoms in combination is known as its valence. This is
not to be taken as a definition, but simply as a description of the
action of chemical valence.

We have already seen examples of chemical action taking plae«
between ions, which are atoms or groups of atoms charged with
electricity. There is abundant evidence furnished by physical chem*
istry, and some of this will be discussed later, that nearly all, if
not all, chemical action is between ions or charged parts. The
establishment of this fact has given us a definite physical basis for
the conception of valence.

An ion is unh^kntf or can combine with one hydrogen atom^
which is really the unit of valence, if it carries one electrical charge.
An ion is bivcUent^ or can combine with two hydrogen atoms, which
are equivalent under all ordinary conditions to one oxygen atom^
if it carries two electrical charges; trlvalent, if it carries three
charges; (imidrivaletitf if it carries four charges; quinquivalent, if
there are five charges upon it; seximdentj if it carries six charges;
aeptivijUejitf if there are seven charges connected \vith it ; and octtv-
cUent, if it carries eight charges* There are no ions known which
have a greater valence than eight, fe« which have the power of
combining with more than eight hydrogen atoms, or four oxygen
atoms.

Faraday's Law the Basis of Chemical Talenca. — It is obviousi
from the alK>ve, that the law of Faraday lies at the basis of chemical
valence, Faraday passed an electric current through a solution of
an electrolyte, and observed that the amount of the electrolyte
decomposed was proiK>rtional to the amount of current which had
passed through the solution. On the basis of this experimentally
e»tahltHhed fact he enunciated the first part of his well-known
law : —

CHLORINE 135

The amount of cliemiml decomposition effected by the passage of the
current, is proportional to the amount of electricity which flows through
the conductor.

Since electricity can flow through a solution of an electrolyte
only by being carried by the ions in the solution, the above part of
Faraday's law shows that each ion of the same substance carries
exactly the same amount of electrical energy.

Faraday determined also the amounts of different elements which
are separated from their compounds, by passing the same current
through solutions of these compounds. A generalization of very
wide significance was reached, which is the second part of the law
of Faraday : —

The amounts of the different elements which are separated by the
sams quantity of electricity bear the same relation to one another as the
equivalents of these elements.

This is saying in other words that all univalent elements carry
exactly the same quantity of electricity, all bivalent elements carry
exactly twice this quantity, all trivalent elements three times the
quantity, and so on. In a word, all univalent ions carry the same
amount of electricity, and all polyvalent ions a simple, rational mul-
tiple of the amount carried by univalent ions — the multiple being
the valence of the ion.

By referring all chemical valence to the law of Faraday, which is
one of the few laws of nature to which no exception is known, we
place what has hitherto been only a name for a large number of em-
pirical facts, upon an exact physical basis.

We have studied thus far three elements which are more or less
typical, — oxygen, hydrogen, and chlorine. There still remain more
than seventy elements, and we might continue our study by taking
these up more or less as suited our convenience. Such a treatment
of chemical phenomena would, to say the least, not be scientific.
This is especially true since cei*tain deep-seated connections between
various elements have been unmistakably established. Certain ele-
ments are very closely allied in all of their properties, while between
others the relationship is more remote, and others still have veiy
little in common. We shall now take a bird's-eye view of the field
yet before us, and see what elements are related most closely to one
another. These we shall naturally treat in close sequence in order
that the resemblances and differences may be pointed out.

CHAPTER X

THE PEHIODIC SYSTEM

Hypothesis of Pront. — Attempts were early made to discover
relationships between the elements ; also relationships between the
chemical properties of the elements aptl certain of their physical
properties, especially their atomic weights. One of the first of
these was pointed ont by Proat as early as IS 15. He observed that
the atomic weights of the elements as then determined were nearly
all whole numbeis when referred to hydrogen as unity. This sug-
gested to him the hypothesis which bears his name, viz. that all
the elements are simply condensations of hydrogen. The atoms of
the different elements are composed of hydrogen atoms, the number
being expressed by the atomic weight of the element. More accurate
determinations of atomie weights have, however, shown that they are
not all simple, whole numbers when referred to hydrogen as the unit,
but in some cases differ markedly from whole numbers*

The hypothesis of Front, while not true as stated by its author,
undoubtedly contains the germ of a great chemical truth.

The Triads of Bobereiner^ — ^Tlie next to point out any important
relation between the chemical properties of the elements and their
atomic weights was Dobereiner, in !82o. He called attention to
relations such as the following. If we add the atomic weight of
caleium, 40,1, to that of barium, 137.4, and divide the sum, 177.5,
by 2j the product is 88.7, which is very close to the atomic weight
of strontium, 87,6. These three elements are obviously very closely
related ohemically.

Again, if we add the atomic weight of sulphur, 32.06^ to the atomic
weight of tellurium, 127.6, and divide the sum, 159.66, by 2, the prod-
uct is 79.83, whieh is very close to the atomic weight of selcniuni, 79.2*
The chemical relationship between these three eleujenta is also very
close, as we shall learn. Similar relationships between the atomic
weights of other groups of three closely allied elements were pointed
out by DcVbereiner, These are known as the TnadH of JMberemer.

The Octaves of Newltndj. — The question of relations between
the atomic weights was taken up by Newlands* In his earlier

13^

THE PERIODIC SYSTEM

137

papers he pointed out connections between atomic weights and
chemical properties, but it was not until 1864 that he announced any
important discovery. In a brief note to the Chemical News, "On
Relations among the Equivalents," he arranged the elements in the
order of their equivalents, and stated that " it will be observed that
elements having consecutive numbers frequently either belong to the
same group or occupy similar positions in different groups. . , . The
difference between the number of the lowest member of a group and
that immediately above it is 7 ; in other words, the eighth element
starting from a given one is a kind of repetition of the first, like the
eighth note of an octave in music." In the following year Newlands
announced his " Law of Octaves " in a very brief note : " If the ele-
ments are arranged in the order of their equivalents with a few slight
transpositions, it will be observed that elements belonging to the
same group usually appear on the same horizontal line. It will be
seen that the members of analogous elements generally differ either
by 7, or by some multiple of 7. In other words, members of the
same group stand to each other in the same relation as the extremi-
ties of one or more octaves in music." The table given by New-
lands brings out the relation to which he refers. It is of such
historical interest that it should be given in this connection.

Newlands's Table

H 1

F 8

CI 16

Co&Ni22

Br 29

Pd 36

I

42

Pt&Ir

60

Li 2

Na 0

K 16

Cu

23

Rb 30

Ag 37

Cs

44

Tl

63

G 8

MglO

Ca 17

Zn

25

Sr 31

Cd 38

Ba&y45

Pb

64

Bo 4

Al 11

Cr 19

Y

24

Ce&La33

U 40

Ta

46

Th

66

C 6

Si 12

Ti 18

In

26

Zr 32

Sn 39

W

47

Hg

62

N 6

P 13

Mn20

As

27

Di&Mo34

Sb 41

Nb

48

Bi

66

0 7

S 14

Fe 21

Se

28

Ro&Ru35

Te 43

Au

49

08

61

A comparison of this table with the periodic system proper will
show that it contains more than the germ of this important general-
ization.

The Periodic System of Hendeleeff and Lothar Heyer. — The
periodic system of the elements, as we now have it, was undoubtedly
discovered independently, and very nearly simultaneously, by the.
Russian, Mendeldeff, and the German, Lothar Meyer. The latter
published in 1864 a table containing most of the then known ele-
ments, arranged in the order of their atomic weights. In this
arrangement elements which are closely allied in their chemical

138

PRINCIPLES OF INORGANIC CHEMISTRY

properties appear in the same columns, but the system is so incom-
plete that it is scarcely an advance on that of Newlands.

The first to point out the most important features in the arrange-
ment of the elements according to their atomic weights was
undoubtedly Mendel^eff. In 1869 he arranged the elements in a
table in the order of their atomic weights, and showed clearly that
there is a periodic recurrence of properties as the atomic weights
iincrease. This will be seen best in the following table: —

Mexdelbeff*s Obigixal Table

{

OROur I

Group II

Gsotr III

Gsour IV

GBOur V

Grol'p VI

Group VII

Group VIII

_

_

_

KIf«

KU,

KII,

KH

K/)

KO

RfO.

KO,

K«<>.

KO,

i;,(),

KO4

1

H=l

•

2

Ll=7

Be=9.4

B=ll

C=12

N=14

0=16

F=19

3

Na=23

Mg=34

Al=27.3

8i=28

P=31

8=32

CI = 35.5

4

K=39

Ca=40

-=44

Ti=4S

V=51

Cr=52

Mu=56

Fe=56, Co=
59, Nl = 59,
Cu=63

5

(Cii=63)

Zn=G5

— =68

-=72

A8=75

8e=78

Br=80

6

Rb=86

8r=87

Y=88

Zr=90

Nb=94

Mo=96

-=100

Rii=10*, Bh

= 104, Pd=

106, Ag=108

7

(Ag=108)

Cd-=112

In=113

8n=118

Sb=122

Te=125

1=127

8
9
10

C8=133

Ba=137

Di=138

Ce=140

—

—

—

— —

Er=178

La=180

Ta=182

w=m

08 = 195, Ir

= 197, Pt =

198, Au=199

U

(An=199)

Hg=200

T1=2M

Pb=207

Bi=208

—

— —

12-

—

—

Th=231

—

U=240

—

— —

m& tible contains all the elements known at that time, and the
Utt&k spM^ indicate that the elements which would naturally fall
iuu *ii«^ ruk"'*s were unknown. Tlie general plan of the Mendel^ff
iBiuf 1* <iTi7c*. All the elements are arranged in succession in the
mw: 't "Ui**:: lacwtiudnjr atomic weights. If we start with the ele-
wnr T-ra -:n« ^mijll^t Atomic weight next to hydrogen, ».e. lithium,
«u. irraacr ^t" siu.'vwJia^ elements in the order of their atomic
»viS3t55t ••• x» hiuriw. w^ 6ml that the next element, sodium, has
ty^-vrn^. "t"v >iuuiur ow ch^HW of lithium. If we place sodium in
-^ ^^wn rrrrviu oiiuttii With lithium, and then arrange the next
. r rtv 'n(t«r .0.' iU«>ir Hifioittic weights, we find that magne-

THE PERIODIC SYSTEM 189

slum falls in the same column with beryllium or glucinum, aluminium
with boron, silicon with carbon, phosphorus with nitrogen, sulphur
with oxygen, and chlorine with fluorine. This is, of course, a re-
markable relation, since in every case those elements which fall in
the same vertical column resemble each other very closely. The
first seven elements, starting (not with hydrogen, since it does not
fit into this scheme) with lithium, and ending with fluorine, agree
very closely in properties with the second set of seven elements
arranged as in the above table. We come now to the first member
of the next series of seven elements, — potassium ; it falls right into
the group with lithium and sodium, calcium with glucinum and
magnesium, titanium with carbon and silicon, vanadium with nitro-
gen and phosphorus, chromium with oxygen and sulphur, and man-
ganese with fluorine and chlorine. Here again striking analogies
appear between the different members in the same groups. The
blank space between calcium and titanium contained no known ele-
ment when this table was prepared. The element has since been
discovered, and has peculiar interest in connection with this whole
system ; to this reference will again be made. After we leave
manganese we encounter one of the weakest points of the Periodic
Law. The next elements in order of atomic weights are iron, cobalt,
and nickel ; but it is obvious that neither of these can be placed in
the same group with the alkali metals. They must, therefore, be set
aside and left out of the system. Then we come to copper, which, is
very questionably placed with the members of group I. Then irregu-
larities appear again. At the end of the sixth series we find three
or four more elements which do not fit into the scheme, but after
leaving these, regularities again begin to manifest themselves.

A more detailed account of the relations between properties and
atomic weights will be taken up a little later. The above suffices to
show the general relation, and also the periodic recurrence of proper-
ties with increase in the atomic weights.

The same general relations as those pointed out by Mendel^fE
were undoubtedly discovered independently by Lothar Meyer, and
published the following year (1870). His table is almost exactly
the same as that of Mendel^eff, and he recognized clearly the periodic
recurrence of properties. To quote his own words, " We see from
the table that the properties of the elements are, for the most part,
periodic functions of the atomic weights."

Meyer has since changed the form of this table, arranging it as a
spiral. "If we regard this table as wrapped around an upright
cylinder so that the right and left sides touch ; therefore, nickel next

140

PRINCIPLES OF INORGANIC CHEMISTRY

to copper, palladium to silver, and platinum to gold, we obtain, as is
easily seen, a continuous series of all the elements in the order of
their atomic weights, arranged in the form of a spiral. The elements
which, in this arrangement, fall into the same vertical column, form
a natural family, the members of which, however, bear a very unequal
resemblance to one another." This spiral arrangement of the ele-
ments is shown in the following table : —

Meyer's Table (using the present atomic weights)

III

IV

VI

VII

VIII

Li
7.03

Na
23.05

K
3W.16

Cu
63.6

Rb
85.5

Ag

i07.as

Cs
132.0

Au
197.2

Be
9.1

Mg
24.36

Ca
40.1

Zn

65.4

Sr
87.6

Cd
112.4

Ba

137.4

Hg
200.0

B

11.0

Al
27.1

Sc
44.1

Ga
70.0

Y

8D.0

In
116.0

La
138.9

Yb
173.0

Tl
204.1

C
12.0

Si
28.4

Ti
48.1

Ge
72.5

Zr

90.6

Sn
119.0

Ce
140.26

Pb
206.9

Th
232.5

N
14.04

P
31.0

V

51.2

As
75.0

Nb
94

Sb
120.2

Ta
183

Bi
208.5

O
16.0

S
32.00

Cr
52.1

Se
79.2

Mo
96.0

Te
127.6

W

184.0

U

2.38.5

F
19.1

CI
35.46

Mn
55.0

Fe
55.9

Co
59.0

Ni
58.7

Br
79.90

Ru
101.7

Rh
103.0

Pd
106.5

I
126.97

Os
191.

Ir
193.0

Pt
194.8

This table brings out more clearly than that of MendeWeff the
idea of a continuous arrangement of all the elements in the order of
their atomic weights. And it is equally successful in showing the
periodic nature of the properties of the elements. The blank spaces
are for unknown elements. Meyer calculated the probable atomic
weights of these elements, but these values being for the most part
unverified, are omitted.

TUE PERIODIC SYSTEM

141

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142 PRINCIPLES OF INORGANIC CllEMISTHlT

One of the iieweBt forms of the Periodic System, and in some

respects the best, is the following, which was recently proposed by
Brauner. This mcludes in group 0 the rare elements discovered ia
the atjnosphere by Eiimsay j however, this atidition to the system
was made some time ago by Ramsay himself. The distinctive feat-
ures of this arrangement are; the grouping of a number of the closely
related, rate elements in group IV, series 8, These elements have
atomic weights ranging from 140 to 173. The ninth series in the
Mendeleeff table, which contains no elements, is entirely abandoned;
the tenth series is made an extension of the eighth, while the eleventh
and twtjlfth series in the Mendel^ff table are made the ninth aad
tenth series in the new talde.

This system lias marketl advantages over the earlier forms. It
includes all the known elements, and what is more important, it
omits the uinth series in the Mendel^eff table, which never had
any real existence, since not a member of this series has ever been
discovered. It also simplifies the system by reducing the number of
series from twelve to ten; and what is perhaps most important, it
brings together those elements which differ from one another in
properties less than any other known elements.

CHEMICAL PROPERTIES AND ATOMIC WEIGHTS
COMBINING POWER

If we start with lithium in ^fendeleeff s table and proceed to the
right along the second series, this striking fact is observed; the
elements increase in their power to combine with oxygen regularly
from left to right* Take first the power of the elements to combine
with oxygeiL Lithium forms the compound Li^O, beiy Ilium BeO,
aluminium AltOs, carbon €0^ nitrogen XjO^j oxygen and fluorine
may be disregarded for the moment. Take the third series* SotUum
forms the compound KaO, which is a superoxide^ magnesium MgO,
alnminium AljOai silicon SiOj, phosphorus PjjOfft sulphur J5*>a, and
chlorine Clp-. The fourth and fifth series show the same regnlari-
ticH, and similar relations are obiierved throughout the table. The
best example of an element octivalent towards oxygen is osmium,
which forms the comjiound OsO^* We havej then, NaO, MgO, AJ,Oaj
SiOj. PjO,, 30.^ Cl,Oj, OsO^

We may saj' in general that the power of the elements to combine
with oxygen is smallest in group I, and increases regularly by unity
in each succeeding group j reaching a maximum in group VIII, w^hete,
at least in the case of osmium^ it is eight

THE PERIODIC SYSTEM 143

Results of a similar character are obtained if we study the power
of the elements to combine with chlorine. Sodium combines with one
chlorine atom, magnesium with two, aluminium with three, silicon
with four, phosphorus with five. Sulphur does not combine directly
with six chlorine atoms, but combines with both oxygen and chlo-
rine, forming the compound SOaClj, in which the sulphur has a
valence of four towards the oxygen, and of two towards the chlorine,
or of six in all. But there is a member of group VI which combines
directly with six chlorine atoms. This is tungsten, in the tenth
series. We. would express the combining power of the elements
towards chlorine as follows; —

NaCl, MgClj, AlCla, SiCl4, PCI,, SO^Cl^

(WCle)

Exactly the same regularity which was observed in the case of
oxygen exists here. The elements in group I have the smallest power
to combine with chlorine, and this increases by unity from group
to group as we pass from left to right; reaching a maximum of six
in the sixth group. We know of no element that has the power of
combining directly with more than six atoms of chlorine.

When we examine the power of the elements to combine with
hydrogen, a regularity is observed, but of a different kind from those
already considered. The elements in groups I, II, and III in general
do not combine directly with hydrogen to form stable compounds,
although hydrides of some of these elements are known. When we
come to group IV, we find in carbon a remarkable power to combine
with hydrogen. The highest valence of the elements towards hydro-
gen is manifested in this group, where one atom of the element com<»
bines directly with four atoms of hydrogen. As we pass to the
right, the power of the elements to combine with hydrogen decreases,
and decreases regularly. Nitrogen combines with three atoms of
hydrogen, oxygen with two, and fluorine with one. Starting with
group IV, we have : —

CH^ NHs, OH2, FH.

The valence towards hydrogen manifests itself to a maximum
degree in group IV, and diminishes regularly as the valence towards
oxygen increases.

The relations pointed out between the combining power of the
elements are general, extending throughout the entire table of the
elements. It should, however, be stated here that there are many
breaks in the system, irregularities appearing on every hand. Some
of these defects will be pointed out in a later paragraph.

lU

PRINCIPLES OF DfORGAXlC CHEMISTRY

Mtfations mitkhi the Groups

In the tikble of MeadeleefE the members of the even Beries ara
|itii€«4 al>ov« one another, and, similarly, the members of the odd
ammi. Eiu.*h group is thus divided into two columns^ whose mean-
ing il first sight is not so appai-ent* If the members of these two
Qolumtis in any group be compared, it will be found that those ele-
mfiiits whioh fall in the same column are more closely allied in their
gsnond properties than the elements in different columns in the
same gnnip. Thus, lithium, pota^sinnij rubidiumj and caesium re-
ieinble eat'h other chenucally more closely than tliey resemble
sodium, copper, silver, and gold. This is more strikingly shomi by
the awoud group, where glucinuuij calcium, strontium, and barium
fall in one culumu, and magnesiumj zinc, cadmium, aud njercury in
the other. Tlie chemical relation between the individuals in a
given column is very close in this group, while it is not so striking
between the members of the different columns. Thus, calcium is
much more closely related to strontium and barium than it is to
zinc or mercury; and, similarly, cadmium is much more closely
allied to zinc and mercury than it is to the calcium group.

Passing to the last group, chlorine, bromine, and iodine fall in the
mme column, and are very similar in thetr ehemieal behavior, while
IbdT relation to manganese is at first sight not very close. These
bot% while purely empirical, are of profound interest, and give to
Hm Fmodie Law a deep significance. It is certainty true that the
t of even series are more closely related to one another than
! to members of o<ld series, and the same obtains for the rela^
ts«ili the odd series. We seem to have here not only a
|y«taa of the elements, but one such system i^ithin

tl^tl^

Bask and Acid Properties

m teltttiQii between the ehemieal properties of the

atomio weights must be pointed out. In any

fmm% with the lowest atomic weight has the

^ with oxygen, as has already been stated.

bttsic character* Thus^ lithium is more

^v^lM^ ii& turn, is far more basic than boron.

^k^ SifDesium, while aluminium begins to

HtlqpilwiAi* Potassium is far more basic

«ji« ^iMi illWitiiim, ciesium than barium. The

uDIi^ Mut iiu^* and silver and cadmium, is

^ ^^ tiWt lki» infill baste elements in the first

THE PERIODIC SYSTEM 145

group, we would expect to find the most acid in the last, and such is
the case. Through the middle groups we find elements which show,
now more, now less basic or acid properties, depending upon condi-
tions ; but in the last column of the last well-defined group we have
elements which manifest only acid-forming properties. The hydro-
gen and hydroxy 1 compounds of the halogens are always acids, and
always react as such with all other substances. These facts are very
surprising. As we pass upward in the table of atomic weights, say
from oxygen, the first element we encounter is fluorine, with very
pronounced acid-forming properties. The element with the next
higher atomic weight is sodium, which is one of the strongest base-
forming elements. Similarly, next to sulphur comes chlorine, which
has much stronger acid-forming properties than sulphur, but next to
chlorine comes potassium, which is one of the most strongly basic
elements. In the same way bromine is followed by rubidium, and
iodine by caesium, where the contrast in properties is quite as great
as in the cases referred to above.

Many other relations between chemical properties and atomic
weights have been pointed out, but those already considered are
among the most important.

Physical Properties and Atomic Weights. — The relations between
many of the physical properties of the elements and their atomic
weights are striking. A number of these have been pointed out by
Lothar Meyer.

Atomic Volnmes. — The atomic volume of an element is the atomic
weight divided by the specific gravity or density of the element in
the solid form. In this connection the atomic weight of hydrogen is
taken as the unit, and the specific gravity of water as the unit of
density. Take the first element in the Periodic System which exists
normjdly in the solid state, — lithium. Its atomic weight is 7, its

density 0.059. The atomic volume of lithium = ** = 11.9.

Meyer plotted the curve showing the change in the atomic
volume with increase in atomic weight, and found that it had re-
markable properties. The curve is shown in Fig. 25. The abscissas
are atomic weights, and the ordinates atomic volumes.

In some cases the specific gravity of the element in the solid form
could not be determined ; as with hydrogen, oxygen, nitrogen, fluo-
rine, etc. In the places corresponding to these elements the curve
is a dotted line.

We see at once from the curve that the atomic volume is a peri-
odic function of the atomic weight. As the atomic weight increases,

140 PRIXC

RIXCIPLES OF INORGANIC CHEMISTRY

a ? 8

S3HmOA 0IH01»

THE PERIODIC SYSTEM 147

the atomic volume decreases and increases regularly. The curve
presents five maxima, at which we find the five alkali metals, —
lithium, sodium, potassium, rubidium, and caesium. At the minima
fall those elements whose atomic weights are approximately the mean
between the atomic weights of the element at the preceding and suc-
ceeding maxima. In fact, at the third, fourth, and fifth minima we
find the elements which do not fit into Mendel^efPs table, and are
placed by themselves in group VIII. We see also in this curve the
distinction between the shoi-t and long periods of Mendel^efPs table.
The first loop of the curve contains the first short period, or the ele-
ments from lithium to fluorine ; the double loop from sodium to nickel
the first long period, and so on. It sometimes occurs that elements
with similar chemical properties have very nearly the same atomic
volumes, as with chlorine, bromine, and iodine.

It is quite remarkable that for elements with very nearly the
same atomic volumes, the properties are markedly different, depend-
ing upon whether the element is on an ascending or a descending
arm of the curve ; and, therefore, upon whether the element with the
next higher atomic weight has a larger or smaller atomic volume
than its own ; e.g, phosphorus and magnesium, chlorine and calcium.
If we follow the curve from its origin, we find the most strongly
base-forming elements at the maxima, and the remainder on the
descending arms of the curve. The acid-forming elements are on
the ascending arms of the curve. Relations between a number of
physical properties and atomic volumes have been pointed out.
These include refraction of light, specific heat, power to conduct
heat and electricity, magnetic properties, etc.

Old Atomic Weights corrected and Hew Elements predicted by
Heans of the Periodic System. — A scientific theory to be of the high-
est value must not simply be able to account for all the facts known,
but must suggest new possibilities which were not realized when the
theory was first announced. The Periodic Law has fulfilled the lat-
ter condition in a beautiful way. By means of it a number of
erroneous atomic weights were corrected. The atomic weight of
indium was supposed to be 75.6, and the composition of the oxide,
InO. This would place it in the Periodic System between arsenic
and selenium. The chemical properties and atomic volume showed
that it belonged rather between cadmium and tin. Meyer gave it
the atomic weight 113.4 (75.6 x 1|), and regarded the oxide as having
the composition lujOg. This was confirmed by Bunsen from
specific heat determinations. The atomic weight of beryllium was
thought to be 4.54, or 4.54x2 = 9.08, or 4.54x3 = 13.62. The

148

PRINXIPLES OF INORGANIC CHEMISTRY

cliemical and physical nature of the element showed that it must
come between lithium and boron, and, indeed, be the head of the
magneaiura -calcium group. The true atomic weight was subse-
quently shown to be 9,1. Similarly, uranium was supposed to have
the atomic weight GO, 120, or 180, and it was difficult to decide between
these values. But it was probably close to 210 in terms of the Peri-
odic System j and this conjecture has also been verified. It should be
observed that in these cases the vapor-density method of determining
the number of atoms in the molecide could not be employed,

The Periodic System has been used not simply to decide between
an atomic weight and a multiple of this quantity, but to actually
correct atomic weights imiK^rfectly determined. Bunseu found the
atomic weight of caesium to be 123.4* This value was smaller than
would be expected from the Periodic System. The correct atomic
weight of caisium was found later to l>e 132,9, which is iu perfect
accord with the system, 31ore I'ccent work in connection with
osmium, iridium, platinum, and gold make it very probable that
the order for these four elements suggested by the system is the
correct one, and that the earlier determinations of atomic weights
contain considerable error.

The prediction of the existence of unknown elements and the
Dature of their properties has been so beautifully verified in a num-
ber of eases that this has become the most striking application of
the Periodic Law. Mendeleeff recognized that the atomic weight
and other properties of an element can be determined from the
properties of the two neighboring elements in the same series and
the two neighboring elements in the same half of the same group.
The properties are as a rule the mean of those of the four elements.
These four elements were termed by Mendeleeif the Atomic AHalogues
of the element in question. This will be clear from the following
example : —

Ca

40

•

Bb

Sr

Yt

85

88

89

Ba

137

THE PERIODIC SYSTEM 149

The atomic weight of strontium is the mean of the atomic
weights of its four analogues, and the same holds in general for the
other properties.

On the basis of this fact Mendel^ff predicted the existence and
properties of a number of elements which had not been discovered
when the Periodic Law was announced. The element predicted was
named from the element in the same group which immediately pre-
cedes it, adding the prefix "eka." In the third group the element
immediately following boron was unknown, and was termed eka-
boron. Since it followed calcium with an atomic weight of 40, and
preceded titanium whose atomic weight is 48, its atomic weight
must be 44. The oxide must have the composition EbjOj, and bear
the same relation to aluminium oxide that calcium oxide does to
magnesium oxide. The sulphate must be less soluble than alumin-
ium sulphate, just as calcium sulphate is less soluble than magnesium
sulphate. The carbonate would be insoluble in water. The salts
would be colorless and form gelatinous precipitates with potassium
hydroxide and carbonate, and disodium phosphate. The sulphate
would yield a double salt with potassium sulphate. Few of the salts
would be well crystallized. The chloride would probably be less
volatile than aluminium chloride, since titanium chloride boils higher
than silicon chloride, and calcium chloride is less volatile than
magnesium chloride. The chloride would be a solid, its volume
about 78, and its density about 2. The specific gravity of the oxide
would be about 3.5, and its volume about 39. Ekaboron would be a
light, non-volatile, difficultly fusible metal, which would decompose
water only on warming ; would dissolve in acids with evolution of
hydrogen, and would have a specific gravity of about 3.

In a similar manner Mendel^eif predicted the existence and prop-
erties of an element between aluminium and indium, terming it ekor
aluminium. The atomic weight would be approximately 68.

Again, an element should exist between silicon and tin, and this
was termed ekasilicium, with an atomic weight of 72.

The properties of the last two elements and their compounds are
described in considerable detail from the properties of their atomic
analogues, but for these the original paper must be consulted.

These elements have now all been discovered. The element
described by Nilson as scandium, proved to be ekaboron, having an
atomic weight of 44.1. Gallium, discovered by Lecoq de Boisbau-
dran, was the predicted ekaaluminium, with an atomic weight of 70.
And germanium, discovered by Winkler, proved to be the ekasilicon,
having an atomic weight of 72.5. The properties of these elements

150

PRtNCirLES OF INORGANIC CHEMISTRY

and their compounds corresponded about as closely with the proper-
ties predicted fur them as the atomic weights.

Imperfeotioiifl in the Periodic System. — WTiile admiring the many
deep-seated rehitions which are brought out by the Periodic ♦System,
we must not fail to observe that it is far from complete. At the
very outset there is evidence of tins incompleteness — hydi-ogeu
does not fit at all into the scheme, and yet it is one of the most im-
portant elements. In the very first group of the elements, again,
there is apparent inconsistency. Along with Hthiunij potassium j
rubidium, and caesium, we find copper, silver, and gold. There is
evidently no very close connection between the last three elements
and the first four. Further, sodium does not fall into the same
division of the group with the other strongly alkaline metals, but
with copper, silver^ and gold. It is at once apparent that sodium is
not as closely allied to these elements as to the alkali metals which
constitute the other division of group L

Passing over the intermediate groups, which contain a number of
more or less serious inconsistencies, we find in group YII manganese
placed with the halogens and not falling into the same gronp either
with chromium or with iron. The relations of manganese to the
halogens are not more striking than the diiferences, and we do not
find manganese falling into the same division of the group with
chlorine, bromine, and iodine, bnt with fluorine, to which it bears a
much less close resemblance than to the remaining halogens.

When M*e come to group YII I, we find nothing but discrepancies.
These elements do not fit into the system at all, and are placed by
themselves m a separate group. It is questionable whether it is
desirable to call this group VlII, since it is in no chemical or physi-
cal sense a true extension of the system one step beyond group VII,
Take as an example the power of the elements to combine with
oxygen. There is a regular increase in this power from unity in
group I, through the several gtoups up to group VII,— where we
find the compounds CIgO-, BrjO„ IjO^, — fluorine not combining at
all with oxygen. Of all the elements in the so-called group VIII,
there is only one, osmiutu, which has a valence of eight towards oxy-
gen. The remainder all show a lower valence towartls this element.

It seems better to recognize these elements as distinct exeeptions^
which do not fit into the Periodic System at all satisfactorily; yet
even here we must recogniEe a certain periodicity in the recurrence
of these exceptions^ and that they occur in every case in groups of
three. The PericKlic System seemed to be hard pressed for a time to
find a plaee for some of the elements described by Ramsay as oeeur-

THE PERIODIC SYSTEM 151

ring in the atmospheric air. Quite recently, however, Ramsay has
shown that these elements have a place in the Periodic System.
These apparent discrepancies in the Periodic System have not been
pointed out with the desire to undervalue the merits of this impor-
tant generalization, but simply to arrest attention to the fact that
the system is still far from complete. What has already been ac-
complished is of tremendous importance, as is shown by the single
fact that we can correct atomic weights and predict the properties
of elements entirely unknown. Indeed, we can do more ; we can pre-
dict with what elements the unknown element in question would
form compounds, the composition of these compounds, and even the
color and other physical properties possessed by them.

We shall probably never have a complete and perfect Periodic
System of the elements until our knowledge of these substances and
their compounds is far deeper than at present If the system was
perfect and complete, it is more than probable that it would lose
some of the interest which it now possesses; since it would then
offer far less incentive to investigation, which is one of the best tests
of the scientific value of any theory or generalization.

General Scheme to be Followed. — In dealing with the remaining
elements we shall be guided largely by the Periodic System. This
system, however, is, as we have seen, defective, and we shall, there-
fore, not follow it blindly, but depart from it whenever the relations
can be more clearly established by doing so.

We shall begin with the members of group VII in Mendel^fTs
table, omitting, however, one member, manganese, which will not be
taken up until much later.

We shall then take up some of the members of group VI — sul-
phur, selenium, and tellurium ; while chromium, molybdenum, tung-
sten, and uranium will not be studied until very much later. The
nitrogen group (V) will then be studied, and following this the car-
bon group (IV).

The metallic elements will then be taken up. Groups I and II
will be studied very nearly in the order indicated in the Periodic
System, while the remaining metallic elements will be studied more
or less independent of the system.

CHAPTER XI

BROBflNB, lODINi:, FLX70RINB

BROMINE (At. Wt.= 79.96)

Ooonrrenoe and Preparation. — The element bromine, which was
discovered by Balard in 1826, and named from its bad odor, closely
resembles the element chlorine. Like the latter it does not occur in
the free condition in nature, and occurs in very much smaller quan-
tity than chlorine. The compounds of bromine, like those of chlorine,
being in general very soluble, the chief occurrence of bromine is in
the waters of the sea. Where sea-water has evaporated and depos-
ited the great salt beds of the earth we find the bromides mixed with
a large number of other salts. Among these should be mentioned
especially the great deposits at Stassfui-t in Germany, from which
much of the bromine of commerce is obtained. Bromine also occurs
in combination with metals as bromides, in many of the mineral
springs of central Europe and Ohio.

Bromine is prepared by three methods : The electrolysis of bro-
mides, which is strictly analogous to the electrolysis of chlorides,
the bromine ion passing to the anode and separating in the free con^
dition, while the metal passes to the cathode.

A second method which was much used formerly, but is now
seldom employed on a large scale, consists in the oxidation of
hrdtobitMuic acid by manganese dioxide ; the hydrobromic acid being
9«l ft^ from the bromide by means of sulphuric acid.

2 NaBr + H^O, = Na,SO, + 2 HBr ;

S HBr + MnO, + HjSO^ = MnSO^ 4- 2 H^O + Br^

vV*«vHttiit^ ihvtM^ in one equation, we have : —

i \.%l^ s. Miu\ 4- 2 UfiO, = MnSO^ + Na^O^ + 2 H,0 -hBr^

t'^*! '^iKt ^tt^Nvxl <\^n»ij»t8 in the replacement of bromine from

SKI^-hCl,-2KCH-Br,.

BROMINE, IODINE, FLUORINE 153

Chemical Properties of Bromine. — Bromine in its chemical prop-
perties strikingly resembles chlorine. Like the latter it unites
directly with most of the elements. The compounds formed — the
bromides — are not as stable as the chlorides. This is shown by the
fact that chlorine replaces bromine from the bromides. The bromides,
like the chlorides, are in general soluble in water, the bromides being
more soluble than the chlorides. In solvents other than water, such
as the alcohols, etc., the bromides are almost always more soluble
than the chlorides.

Bromine has a remarkable power to disintegrate organic sub-
stances. Like chlorine it replaces the hydrogen in such compounds,
and effects even deeper transformations. Its action upon the mucous
membrane of the throat and nose is much more vigorous than that
even of chlorine, and great precaution must, therefore, be taken in
working with this substance, to be protected from its disintegrating
fumes. A bottle containing bromine should never be opened except
under a hood with good ventilation.

Detection of Bromine. — The chemical properties of bromine are
so closely allied to those of chlorine that it might at first sight seem
diflBcult to determine with which we are dealing, especially if it is
combined with other substances. This difficulty is, however, only
apparent. If a bromide is treated with chlorine, the bromine, as we
have seen, is set free and can be recognized by its odor and color.
If a solution of a bromide is treated with a little chlorine water, a
little carbon disulphide being added to the tube and shaken vigor-
ously, the bromine which has been set free by the chlorine is dis-
solved by the carbon disulphide and imparts its characteristic
reddish-brown color to the solution.

Again, if we add a solution of silver nitrate to a solution of a bro-
mide, the bromine ion combines with the silver ion, forming silver
bromide, winch is practically insoluble in water. Silver bromide,
however, is white like silver chloride, and the eye could not distin-
guish between them. It might, therefore, seem difficult to deter-
mine whether we were dealing with a chloride or a bromide. Such,
however, is not the case, since silver chloride readily dissolves in
ammonia, while silver bromide is much less soluble. When we have
a precipitate formed which we know is either silver xxhloride or sil-
ver bromide, it is only necessary to add a few drops of ammonia.
If the precipitate dissolves it is the chloride, if it does not disappear
in solution it is the bromide. In making this test care must be
taken not to add much ammonia water, since silver bromide is solu-
ble also in a large amount of ammonia.

154

raixcirLEs of inorganic chemistry

Bromine Atoms and Bromine lozu. — Bromine in the atomic con-
dition behaves very differently from bromine in the ionic condition.
We have just seen how the iiromiii© ions behave when brought in
0Oiita4?t with the silver ions. They combine with them at once, form-
ing the characteristic precipitate, silver bromide.

K, Br + A^g, m\ ^ K, m^ + AgBn

When bromine in the unionized condition is brought into the pres-
enee of silver ions, it does not combine to the slightest extent with
the silver ions. Thus, bromine in such compounds as ethyl bromide,
CtH^Br, is not in the ionized condition. It is presumably in the
atomic condition, forming a part of the undisaociated molecide of
the non-electrolyte, ethyl hroniide* Bromine in this condition doea
not rt^act with a jiolution of gjlver nitrate — with silver ions.

Fhyiical Fropertiea of Bromina — The reddish-brown liquid, bro-
mine^ has a density of 3.1. It boila at 03^, yielding a reddish-yellow
va|>or. The density of the vapor is 160 in terms of hydrogen as two,
provided the bromine vapor lias not been heated above 300^ The
atomic weight Vieing 80, the mdecale of bromine therefore contains
two atoms, or has the composition represented by the formula Br,.
If the vapor of bromine is heated above 3<>0°p the density becomes
less as tlte teini>erature is raised^ showing that the molecules of Bfj
are gradually breaking down into molecules of Br,

Bromine solidities at — 7"*, forming an orange-red solid. A beauti-
ful experiment consists in introducing some brtimine into a flask
with a narrow neckj and drawing tlie neck out to a fine tube. The
tobe is se^^led, and is then full of bromine vajKjr, Upon the bottom
ef the flask place some solid cartmn dioxide and ether, or, still Ijetter,
mbltk liquid air. At these extremely low temperatures the bromine
I St once from tlie sfiot on the glass where the solid carbon
mr the liquid air was i>!aeei

^BUt dissolves in water, forming what is known as brmnine

MA m loalogous to chlorino water. The saturaterl solution

111 iioiitliBtr i!^"'t 3 per cent of bromine. When a solution of

»iiCSP09ed to the light, hydrobi-omic acid is formed

^ lust as when a solution of chlorine water is

t lirdtocWoric acid is formed and oxygen ja lil^r-

' a kfdnte with cold water analogouxs to ehlo-

1^^ liydrftte is more stable than chlorine

0( bromine hydrate is apparently

, te change with the conditions. Bro-

[ 4fc M^^etion with photography and with

Ti^

BROMINE, IODINE, FLUORINE 155

Hydrobromic Acid, HBr. — When hydrogen and bromine are
mixed and the mixture subjected to an electric spark or exposed to
the light, the elements do not combine with explosive violence as
was the case with chlorine and hydrogen. They do combine, how-
ever, to a cei-tain extent, forming hydrobromic acid. The amount of
combination can be greatly increased by passing the mixed gases
over finely divided, hot platinum, which acts catalyticaJly upon the
mixture.

Hydrobromic acid can be prepared by the action of bromine upon
compounds containing carbon and hydrogen. The bromine displaces
a part of the hydrogen, combining with it and forming hydrobromic
acid. When bromine acts upon benzene the following reaction takes

place: —

CeH« + 2 Br, = QH^Br, + 2 HBr.

The best method, however, of preparing hydrobromic acid is by
allowing water to act upon the bromides of certain acid-forming ele-
ments such as phosphorus. The reaction that takes place in this
case is —

PBr5 + H^O = POBr, -|- 2 HBr,

resulting in the formation of phosphorus oxybromide and hydrobromic
acid.

A method which was formerly used, but which is far less satis-
factory than that just described, consists in treating phosphorus with
bromine in the presence of water. The phosphorus and bromine
combine, forming the tribromide or pentabromide of phosphorus,
depending upon the amount of bromine used. If the tribromide is
used, the reaction which takes place is the following : —

PBra -h 3 H,0 = HaPO, + 3 HBr.

[Phosphorous acid.]

If the pentabromide is formed, the decomposition with water takes
place as represented by the above equation, or if more water is added
the oxybromide is decomposed into hydrobromic acid and phosphoric
acid, thus : —

POBr, + 3 H,0 = H8PO4 H- 3 HBr.

The question which would naturally be asked is why not prepare
hydrobromic acid by the action of sulphuric acid on bromides?
This would be analogous to the preparation of hydrochloric acid by
the action of sulphuric acid on chlorides.

This method is theoretically possible, but precaution must be
taken or a secondary reaction between the hydrobromic acid formed
and the sulphuric acid takes place, which interferes with the value
of the method.

156 PRDTCIFLES OF INOBGANIC CHEiUSTRr

tf the salptiiuic Beii is dilute^ the hydrobTOmic aeld formed di^
solTes in the aqueous sulphuric acid od account of its great solubility
in water. If, on the other hand, the sulphuric acid used is concen-
trated, the following reaction takes place: —

MSO, + 2HBr = 2H^ + SO* + Br^

The hydrobTOmic acid is oxidized to bromine by the sulphnric
acidf which is reduced to sulphur dioxide.

Properties of Hydrobromic Acid. — Hydrobromic acid resembles
hydrcjchlonc a^^id very closely iu its chemical and physiesU pro^ierties.
It is a colorless gas with penetrating odor, very soluble in water.
An aqueous solution saturated at zero contains about 80 per cent of
the acid. Such solutions give off dense fumes when exposed to the
air, or when the breath is blown across the mouth of the flasks which
contain them.

Hydrobromic acid is a reducing agent^ as we have seen in its
action on sulphuric acid* It is also a very strong acid, A dilute,
aqueous solution of hydrobromic acid is completely dissociated into
its ions : — + -

HBr = H, Br.

The bromine inns manifest their presence by combining with
the silver ioiis when brought in contact with them, forming tlie
white precipitate, silver bromide, which is soluble with difficulty
in ammonia.

When chlorine is brought into the presence of a bromide, as we
have seen* or of hydrobromic acid, the bromine separates and chlorine
passes into sotution. If we examine this reaction more closely, we
find that what has taken place is a transfer of the electrical charge
from the bromine ion to the chlorine atom, converting the latter into
an ion, while the bromine ion having lost its charge is converted into
an atom. The atoms of bromine then combine and fonn the molecules
of bromine. The reaction which takes place would be represented
thus : —

H, Br + CI = H, CI + Br,

This action of chlorine on hydrobromic acid is analogous to the
action of metals on acids in general. In the former case we have a
transfer of the negative charge from the anion bromine to the chlorine,
ttonverting the bromine into an atom, and the chlorine into an anion.
In the latter case the positive cljarge is transferred from the cation
hydrogen to the metal^ converting the hydrogen into an atom and
the metal into a cation. The main difference between the two re-

BROMINE, IODINE, FLUORINE 157

actions is that in one case we have a transfer of the negative charge,
in the other case a transfer of the positive charge.

When a concentrated solution of hydrobromic acid in water is
cooled suflBciently, a crystalline hydrate separates having the compo-
sition HBr.2HsO.

Hydrobromic acid gas can be liquefied. It boils at — 73** and
solidifies at — 87**.

Compoundfl of Bromine with Oxygen and Hydrogen. — Bromine
forms two well-characterized acids with oxygen and hydrogen.
These are hypobromous and bromic acids. While hydrobromic acid
is less stable than hydrochloric, as we have seen, oxygen acids of
bromine are more stable than the corresponding acids of chlorine.

When bromine is conducted into water in the presence of mer-
curic oxide, hypobromous acid is formed : —

HgO H- 2 Br, -h HjO = HgBr, + 2 HBrO.

The acid resembles very strikingly hypochlorous acid. Like the
latter it gives up its oxygen readily, being, therefore, a good oxidiz-
ing agent.

The sodium salt of hypobromous acid is prepared by the action
of bromine on sodium hydroxide : —

2NaOH-|-2Br = H,0-|-NaBr4-NaOBr.

This method of preparing sodium hypobromite is strictly analo-
gous to the method of preparing sodium hypochlorite.

Bromic Acid, HBrOj, is prepared by oxidizing bromine by means
of chlorine monoxide, also by fusing bromides with chlorates ; the
bromide being converted into the bromate and the chlorate into the
chloride.

The best method, however, of preparing bromic acid, is by the
action of bromine on caustic potash. A mixture of potassium bro-
mide and bromate is formed : —

6K0H 4- 3 Br, = 5 KBr + KBrOj + 3 H,0.

From this mixture the potassium bromate can be readily separ
rated by fractional crystallizationy — the same method which was em-
ployed to separate potassium chlorate from potassium chloride.
The bromate is much less soluble in water than the bromide, and
readily crystallizes from a not too dilute solution of the two salts in
water.

Bromic acid is obtained from potassium bromate by methods
strictly analogous to those employed for obtaining chloric acid from
potassium chlorate, by treating barium bromate with dilute sul-

158

PRINCIPLES OF INORGANIC CHEMISTRY

phuric acid, or silver broiuate with dilute hydrochloric acid. The
barium aulphate, or silver cliloride formed, ia insoluble and can be
readily filtered off from the solution of bromic acid.

The bromine analogue of perchloric acid — perbromic acid —
does not exist

No compounds of oxygen and bromine have been made. If they
are formed they are too unstable to be isolated.

CampQEnd of Iromiae with Chlorme^ BrCl. — Bromine combines
with chlorine, forming one compound, brtjmine chloride, having the
composition BrCL It is formed when cold chlorine gas is conducted
into liquid bromine. It is a reddish-brown liquid, which decom-
poses at 10^

It ia rather surprising that this compound should exist, when we
consider how closely allied chlorine and bromine are in their chemi-
cal properties ; it being a general rule that the more closely related
elements are the least liable to enter into chemical combination.

IODINE (At Wt= 120.85)

Occurrence and Preparation. — Iodine, w^hich was discovered by
Courtois in 1812, and named from the violet-bhie color of its vapor,
occurs very rarely in the free condition. It is widely distributed in
nature, occurring, however, usually only in small qnautities. It
occurs along with chlorides and bromides, but in very much smaller
quantities than either of these substances. It also occurs in sea-
water, as we would expect, on account of the solubility of its com-
pounds. It is taken up from the waters of the sea by certain sponges
and plants, and exists in considerable quantity in the ashes of such
plants. It occurs also in certain ores, especially those of silver, and
in the deposits of soda saltpetre in Chili and Peru. It also occurs
in small quantity in deposits of rock-salt.

Iodine is obtaineii to-day mainly from CfiiU mUpetrej in which
it occurs in the form of sodium iodate. The iodine is obtained from
this salt by reduction with aulphnrous acid : —

2 NalO, + 5 H^O, ^ 4 H^SO, + Na^SO* + H,0 + 1^

The sulphurous acid is oxidized to sulphuric acid, the iodate being
reduced and iodine set free.

There are other methods of obtaining iodine, such as the oxida-
tion of hydriodic acid by one or another oxidizing agents,

4PII + Ot = 2H,0-»-2It,

BROMINE, IODINE, FLUORINE 159

The agent raost frequently used is manganese dioxide. When an
iodide is treated with manganese dioxide and sulphuric acid the fol-
lowing reaction takes place : —

2 KI -f MnO, H- 2 HsS04 = MnS04 H- K8SO4 + 2 HjO + 1^

Iodine can also be displaced from iodides by means of chlorine.
When a solution of potassium iodide is treated with chlorine water
the following reaction takes place : —

K,i4-Cl = K,Ci4-L

This is analogous to the action of chlorine upon bromides, which
we have seen consists in a transf errence of the electrical charge from
the bromine ion to the chlorine. Here we have the charge trans-
ferred from the iodine ion to the chlorine, which becomes an ion, the
iodine having lost its charge becoming an atom.

Chemical Properties of Iodine. — The purplish-black solid, iodine,
resembles strongly in its chemical properties the elements chlorine
and bromine. It combines with many other elements, but is not
quite as active chemically as bromine and chlorine. Bring together
a few flakes of iodine and a small piece of dry phosphorus in a por-
celain dish. Combination between the two will take place at once,
resulting in the formation of an iodide of phosphorus. Iodine, like
bromine and chlorine is a strong oxidizing agent. When brought in
contact with substances which can take up oxygen in the presence
of water, it takes the hydrogen from the water forming the com-
pound hydriodic acid, with which we shall soon become familiar;
and the oxygen is taken up by the oxidizable substance. Thus,
when iodine is brought in contact with the easily oxidizable com-
pound sulphurous acid, its color disappears rapidly and the iodine
passes into solution as hydriodic acid.

H,SOa -f- H,0 H- 1, = H,S04 + 2 HI.

The oxidizing power of iodine is frequently made use of to determine
the quantity of this substance present.

Detection of Iodine. — Iodine forms a characteristic blue color
with starch paste, which enables its presence to be easily detected.
The starch paste must be carefully prepared in order that the reac-
tion may be sensitive. Some granules of starch should be placed in
a porcelain mortar and a little cold water added. The starch should
be ground with the water to a fine paste. The paste should be
poured into a beaker, and hot (not boiling) water added in consider-
able quantity with vigorous stirring. Under these conditions some

160

raWCrPLES of inorganic CFIEJnSTRT

of tlie starch dissolves in the water. After the solution has stood
for a time the supernatant liquid i^ poured off, and is then ready to
be used in detecting iodine. This reaction is so sensitive that it can
be used to detect a mere trace of free iodine.

If the iodine is combined as in the iodide of a raetal^ it can he
detected by simply adding concentvated sulphuric acid. The hydri-
odic acid set free reacts with the sulphuric acid reducing it, and the
iodine is liberated as such. It can be recognised by the characteristic
dark-brown color which it imparts to the solution.

Detection of Iodine in the Presence of Bromine and Chlorine, —
We have seen that iodine resembles closely iu its chemical behavior
both bromine and chlorine. The question arises^ How^ can iodine be
detected in the presence of one or both of these closely related
elements ? Tliere are several methods by which this can be effected,

If to a solution containing an iodide, bromide, and chloride, a
little chlorine water is added, the iodine will separate first. If a
little carbon disulphide is added and vigorously shaken with the
solution, it will dissolve the free iodine and acquire the characteristic
porplish-red color of this substance. In detecting iodine by this
method care must be taken to add the chlorine water drop by drop
and not in excess* If an excess of chlorine is introduced, it will
cause the free iodine to be oxidized to iodic acid, which we shall
study a little later, and which is colorless.

If, after the iodine has separated, more chlorine water is added
to the solution, the iodine color will disappear for the reason indi-
cated above, and free bromine will then begin to separate, which
will give its characteristic reddish-brown color to the carbon disul-
phide. It is thus possible to detect iodine in the presence of bromine
and chlorine, and also bromine in the presence of the other two sub-
stances.

Another method for detecting these three cJosely related ele-
ments is based upon the different solubilities of the three silver salts
in ammonia. We have seen that silver chloride is very easily solu-
ble in ainmonia, and that silver bromide is soluble with difficulty-
Silver iodide is insoluble in ammonia, or soluble to only such a
slight extent that it is practically insoluble. Further, while silver
chloride and bromide are wiute, silver iodide is yellow.

If to a solution containing the three substanceSj chloride, bromide,
and iodide, silver nitrate is added in excess, we will have the three
salts of silver — chloride, bromide, and iodide — precipitated. The
precipitates are now filtered off and treated with a few drops of dilute
ammonia. If any appreciable quantity of the precipitate dissolves,

BROMINE, IODINE, FLUORINE 161

this is silver chloride. Enough ammonia is now added to the pre-
cipitates to dissolve all of the silver chloride. If on further addition
of a considerable volume of ammonia more of the precipitate dis-
solves, this would mean that we had bromine present. If a yellowish
precipitate remains behind which is insoluble in ammonia, this is
silver iodide.

In determining whether, under any given conditions, a part of
the precipitate has dissolved in the ammonia, a little nitric acid is
added to the ammoniacal solution. Any precipitate held in solution
by the ammonia would be again precipitated.

Physical Properties of Iodine. — Although iodine is a solid at
ordinary temperatures, it can readily be converted into vapor at a
slightly elevated temperature. When iodine is heated exposed to
the air, i.e, under ordinary conditions, it does not melt, but passes
at once into vapor. It can, however, be melted at 114**, and boils at
184^

The vapor-density of iodine between 200** and 600** is such as to
show that the molecular weight is 256. Since the atomic weight is
126.85, the molecule of iodine at these temperatures consists of two
atoms. As the temperature rises, the German chemist, Victor Meyer,
has shown that the vapor-density decreases, and that above 1400®
the density is only about one-half the value at the lower temperature.
Above 1600° it is quite certain that the vapor-density of iodine would
remain constant, since at this temperature the atom and molecule
would be identical, and no further dissociation of the molecules could
take place.

The vapor of iodine can be readily condensed to a solid when the
temperature of the vapor is again lowered. This process of convert-
ing a solid into a vapor, and recondensing the vapor to a solid, is
known as mbUmation. The best method of purifying iodine is to
sublime it The impurities, being for the most part non-volatile at
the temperature at which iodine volatilizes, remain behind in the
solid form.

Iodine dissolves in water to only a slight extent. If to the water
potassium iodide or hydriodic acid is added, the solution dissolves
iodine in considerable quantities. Iodine dissolves readily in carbon
disulphide, chloroform, alcohol, and ether. Solutions of iodine in the
last two solvents are known as tincture of iodine.

Hydriodic Acid, HI. — When a mixture of hydrogen and iodine
is heated, there is partial combination between the two forming hydri-
odic acid, HI. Only a portion of the mixture, however, combines.
The velocity of the reaction, i.e. the amount of hydriodic acid which

162

PRINCIPLES OF mOHGAKTC CHEMISTRY

is formed in a given time can be greatlj increased by heating the
mixture in the presence of finely dividal platinum, which acts by
contact, or catalyficallij, as we say. Even under these conditions a
complete combination of the two gases cannot be effected*

A far more convenient method of preparing hydriodic acid is by
the action of water on phosphorus triiodide, Vl^\ —

PI, + 3 H,0 = II.PO, + 3 III,

or by the action of iodine on phosphorus in the presence of water; —

3I + P+3H,0 = H»POa + 3HL

Hydriodic acid cannot be prepared by the action of sulphnrie
acid on iodides, since, as we have seen, hydriodic acid reduces sul-
phuric acid. The extent to which this reduction takes place depends
largely upon the temperature and concentration of the solutions.
If the solutions are cold and dilute, the rednetion of the sulphuric
acid may only proceed as far as sulphur dioxide, —

H3SO, -h 2 HI = Is + 2 HjO + SO,.

If th© solutions are more concentrated and warmer^ we may
have free sulphur formed: —

HgSO,4'6HI=?l2 + 4H30 + S,

If the solutions are still more concentrated and hot, the reduction
may proceed still farther and give us, from the sulphuric acid,
hydrogen sulphidcj — the lowest reduction product of sulphuric acid.

H(SO, + 8 HI =4 1, + 4 H^O -h H,9.

Hydriodic acid is in general a strong reducing agent, readily

giving up its hydrogen to substances which can take it, or taking
oxygen from substances which can lose it. This is due to the ease
with which hydriodic acid is broken down, or dissociated by heat, into
hydrogen and iodine.

The gas, hydriodic acid, is readily liquefied. At 0* a pressure of
four atmospheres is suHicient to convert the gas into a liquid. The
liquid is readily solidified, the solid melting at — ol^

The gas dissolves very readily in water, one volume of water
dissolving, at 0**, about 500 volumes of the acid. There is a solution
of hydriodic acid and water which has a constant boiling-point.
This contains 57 per cent of acid and boils at 126*. This is analo-
gous to the constant boiling mixture of hydrochloric acid and water.
Like the latter, also, it is not a definite chemical compound^ since its

BROMINE, IODINE, FLUORINE 163

composition can be changed by varying the pressure under which it
is boiled.

An aqueous solution of hydriodic acid quickly turns brown when
allowed to stand exposed to the air. This is due to the oxidation of
the acid by the oxygen of the air, and the consequent liberation of
iodine.

We have seen that iodine dissolves far more readily in a solution
of an iodide, or hydriodic acid, than in pure water, — in a word, in
a solution already containing a large number of iodine ions. We
can now see a reason for this rather remarkable fact. The iodine
combines with the iodine ion, already in the solution, forming the

complex ion, I3, which is brown, while the iodine ion, I, is colorless.

This complex ion, Ij„ gives its characteristic brown color to the
solution.

Compounds of Iodine with Oxygen and Hydrogen. — Iodine
forms two well-characterized compounds with oxygen and hydrogen.
These are iodic acid and periodic acid. When iodine is dissolved
in caustic potash the reaction takes place thus : —

6KOHH-3l2 = 5Kl4-KI03 4-3H,0.

It is probable that during this reaction potassium hypoiodite is
formed. If so, this is so unstable that it breaks down into iodide
and iodate.

The iodates are more readily prepared by the action of iodine on
chlorates in the presence of water.

5 KClOs + 3 1, -h 3 HjjO = 5 KID, 4- HIO3 4- 5 HCl-

When no water is present we have : —

KC103 + I = KI03+C1.

When potassium iodate is treated with barium chloride, barium
iodate, which is difficultly soluble, is formed, and potassium chloride,
in accordance again with the general principle, that whenever an
insoluble, or difficultly soluble compound can be formed, it is formed.

2 KIO3 H- BaCl = Ba(I03), + 2 KCl.

When barium iodate is treated with a dilute solution of sulphuric
acid the following reaction takes place : —

Ba(I03), H- H2SO4 = BaS04 + 2 HIO3,
barium sulphate being much more insoluble than barium iodate.

164 PEINCIPLES OF INORGANIC CHEMISTRY

This is the best method of preparing icxiic acid.
Iodic acid is a well-crjstalUzed compound^ soluble in water, and
a very strong acid.

It may dissociate thus : —

HIO, = H,IO^

or thus : 2 HIO^ ^ H, H(IO,)^

or thus : 3 H 10, = H, Hs(I O^^p

ghoTOH by the fact that we have salts of tlie foUowitig compositions: —

MIO» MHCIOs)^, and MH,(IO,V

When iodic acid is cajefiilly heated above 100', it loses water
and passes over into a white powder, iodine pentoa^id^f —

2HI03=HaO-hrA^

Iodine pentoxide dissolves in water, combining with it and form-
ing again iodic acid. Iodic acid and iodine peotoxide readOy give
up their oxygen, and are^ therefore, good oxidizing agents.

When the salts of iodic acid are heated or subjected to the action
of strong oxidizing agents, they pass over into the salts of periodic
acid.

Periodic acid does not have a composition strictly analogous to
perchloric or perbromic acid, but this plus two molecules of water : —

HI04 + 2H/>=H,IO«.

The acid is best prepared by trausfomnng the acid sodium salt,
KajKjIOfl, into the acid silver salt, AgjH^IOu, the acid sodium salt
being prepared by conducting chlorine gas into a boiling solution of
sodium iodate to which sodium carbonate ha*H been added.

When the silver salt, Ag^Hj^IO^ is decomposed with water it
yields periodic acid, HJOq, and the silver salt of this acid, AgjIO^: —

5 Ag.HJO, ^ 2 Ag,IO, + 3 H,IO^

Periodic acid, H^IOe, is easily soluble in water, and when gently
heated loses water and passes over into iodijie sepioxide^ laO,, When
heated to 140°, periodic acid loses oxygen as well as water, and
passes over into iodide pentoxide, IjOj,

Some of the periodates indicate that the composition of periodic
acid is HIO|, others that the acid baa the comtKisition, HglO^ and
others still that the acid ia H^IOj. These facts, however, are not

BROMINE, IODINE, FLUORINE 165

inconsistent. One of these acids is easily derived from the others
by the addition or subtraction of water : —

HI04 + H,0 = H3l05,
H3l05-hH,0 = H,IOe.

Periodic acid readily loses oxygen, and is therefore a good oxidiz-
ing agent.

Compounds of Iodine with Chlorine, ICl, ICl^. — Iodine forms two
compounds with chlorine, ICl and IClj. When chlorine is brought
in contact with iodine in the absence of moisture, both of these com-
pounds are formed. Iodine monochloride is a reddish-brown liquid,
boiling with partial decomposition at 101°, and forming crystals
which melt at 14° or 27°, depending upon the conditions of their
formation. The one with the higher melting-point is the more stable
form.

When an excess of chlorine is conducted over iodine, the trichlo-
ride is formed. It is a reddish-yellow solid, breaking down easily
into chlorine and the monochloride. Both of these compounds are
decomposed by water into iodine, hydrochloric acid, and iodic acid.

Compound of Iodine with Bromine. — Iodine forms only one com-
pound with bromine — iodine bromide — having the composition IBr.
It is an unstable solid, easily decomposed by water or by rise in
temperature.

The relation pointed out earlier in the case of bromine obtains
here. Iodine is more closely allied to bromine, chemically, than it
is to chlorine. With the latter it forms two compoimds, and with
the former only one, which is very unstable. The more closely
allied elements are least likely to enter into chemical combination.

FLUORINE (At. Wt. = 19.0)

Ocourrence and Preparation. — Fluorine, on account of its unusual
chemical activity, does not occur in nature in the free condition.
It occurs mainly in combination with the element calcium as fluor
spar, from which it derives its name. Fluor spar is so called be-
cause it readily melts and flows, serving as a flux for other sub-
stances. Fluorine also occurs in another mineral, cryolite^ in
considerable quantity. Cryolite is a double fluoride of sodium and
aluminium, occurring mainly in Greenland, and having the com-
position NasAlFg.

The problem of isolating fluorine remained for a long time un-
solved. On account of the great chemical activity of this substance^

166

PRINCIPLES OF INORGANIC CHEMISTRY

m soon as it was set free from its compounds it would combine again
with whatever it came in contact The problem of isolating fluorine
was solved by the French chemistj Moissan. He electrolyzed hydro-
fluoric acid and obtained hydrogen at the cathode and fluorme at
the anode, H there is any water present, the fluorine as rapidly as
formed would act upon it and decompose it, yielding oxygen and
hydrofluoric acid. An aqueous solution of hydrofluoric acid, there-
fore, cannot be used. Liquid hydrofluoric acid or the anhydrous
gas cannot be used, since they do not conduct the electric cuiTcnt-

l-S

Fio. *iiiv

Moissan found that when potassium fluoride is dissolved in an-
hydrous hydrofluoric acid the solution cotuluets the current^ hydrogen
being libemted at the cathode and fluorine at tlie anode. At first
the attempt was made to use vessels lined with calcium fluoride
where the fluorine eseaives, but vessels of platinum were subsequently
employed and found to work very satisfactorily. Indeed, it has
subsequently been shown that fluorine does not act very vigorously
upon copper, and copper vessels have been used in which to liberate
the element fluorine.

BROMINE, IODINE, FLUORINE 167

The apparatus used by Moissan for preparing fluorine is shown
in Fig. 26. The apparatus and electrodes, made of platinum-iridium,
have the form shown in the figure. The electrodes are insulated by
means of stoppers (S.S.) of fluor spar. The apparatus is kept at a
temperature of — 23° by means of methyl chloride. The fluorine,
liberated on the anode, passes out through a spiral platinum tube
cooled to — 50° by means of methyl chloride, and then through two
tubes of sodium fluoride to remove all traces of hydrofluoric acid.
The apparatus holds about 160 cc. The solution which is to be elec-
trolyzed contains 20 grams of potassium fluoride dissolved in 100
grams of anhydrous hydrofluoric acid.

Chemical Properties of Flnorine. — Fluorine is one of the most
active chemically of all the elements. It replaces chlorine from
chlorides and from hydrochloric acid. At ordinary temperatures it
combines with most of the metals, converting them into fluorides.
Platinum and gold are about the only metals which resist its action,
and these are transformed into fluorides at elevated temperatures.

What is more surprising, fluorine combines also with most of the
metalloids, and at ordinary temperatures. The only elements which
resist the action of fluorine are nitrogen, chlorine, oxygen, and argon
and its associates.

When fluorine decomposes water, the oxygen which is liberated
contains a large amount of ozone.

Moissan describes a number of very beautiful experiments, where
the fluorine as it escapes from his apparatus is allowed to come in
contact with various metallic and non-metallic elements. Many of
these take fire and burn readily in a stream of fluorine, at ordinary
temperatures, evolving a large amount of light and heat.

Physical Properties of Flnorine. — Fluorine is a gas at all ordinary
temperatures, having a light, greenish-yellow color. It is much
lighter in color than chlorine, and much more active upon the mu-
cous membrane of the nose and throat. Its vapor-density indicates
that at ordinary temperatures it is broken down in part into mole-
cules which are identical with the atom. Its vapor seems to be
composed in part of molecules of Fj and in part of molecules of Fj.

Similar relations were observed in the cases of bromine and
iodine, but only at much higher temperatures.

Fluorine has been liquefied by the combined efforts of Moissan in
France, and Dewar in England. The fluorine was cooled to — 190**
by means of liquid air, when it liquefied and was received in a glass
bulb with a vacuum-jacket. Fluorine boils at — 187°, and at this
low temperature has lost much of its chemical activity, as is obvious

108

PRINCIPLES OF INORGAOTC CHEMISTRY

from i\m fact that it can be received in a glass vessel Liquid fluor-
ine doefi not act uijon iron, and does not even replace iotline from its
Ooiii|Hjmirlr*. All attempts to solidify fluorine, except the most reeentj
Vern unautu^essfuJ, It has recently been converted into a solid.

Hydrofluoric Acid, HF. — Hydrogen fluoride, or hydrofluoric acid,
in pti^iared mo^it couveniently by the action of sulphuric acid on
calcium fluoride; —

CaFjH- H^SO, = CaSO, + 2 HF.

Th« most characteristic chemical property of hydrofluoric acid is its
]jower to act upon glass, etching it, as we say* It is extensively u&ed
for this purpose in preparing measuring apparatus especially for
chemical work. Hydrofluoric acid does not act upon paraffine and
similar organic substances. The glass vessel upon which it is
desired to make a pernjanent line is covered with pjtraftine by dip-
ping it into the molten material. A fine line is then dra^^Ti through
the parafi^Qe at the place where it is to apiiear ou the glass. The
glass is thus exi>03ed at this place. It is now subjected to the action
of the fumes of hydrofluoric acid. Where the glass is protected by
tlie paraffine, it is not acted npon by the fumes of the acid. Where
the |>iuaffiiie has been removed, however, the glass is etched.

This can be readily tried in the laboratory by covering a watch-
crystal with parafline and scratching a line, letter, or number upon
it* Then exiHJse the glass to the action of the fumes of hydrofluoric
acid* After a time remove the parafline by dissolving it in oil of
turpentine, and the crystal will be etched wherever the ].>araffine has
been removed.

It ID obvious from the above that hydrofluoric a*ml cannot he pre-
pared or kept in glass bottles. It does not attack platJtmm vessels,
and acts only slightly upon vessels of lead. It is presen^-ed in ves-
sels of gutta iwicha, upon wliich it acts only slightly. Hydrofluoric
acid is much weaker than either hydrochloric, hydrobromic, or hydri-
odic ftC5id. It is readily soluble in water, and its aqueous solution
is tint fiirm in which it is nearly always employed. Like the ai^ids
m**ijtiuni*d uliove, it fanus a constant boiling mixture with water,
but thiis is not a definite compound*

The vapor*density of the pure acid would show a molectde much
morrt complex than wovdd l>e indicated by the formula HF, The
oom position of cryolit<% Na^AlF^, points to the same conclusion.
There are a uuiuIrt of examples of fluorine and the other halogens
tending to form aggregates of m% atoms^ as in cryolite.

Fluorine differs from chlorine, bromine, and iodine in that its siU

BROMINE, IODINE, FLUORINE 169

ver salt is readily soluble and its calcium salt insoluble. Anhydrous
hydrofluoric acid boils at 19°.4, and solidities at 92°.5.

On account of the great chemical activity of fluorine and hydro-
fluoric acid, they are very dangerous substances to work with, acting
upon organic matter with great vigor. Special precautions must
therefore be taken in dealing with them.

Compound of Flnorine with Iodine, IF5. — One compound of fluor-
ine and iodine is said to have been prepared. This has the compo-
sition IF5, and is obtained by allowing iodine to act on dry silver
fluoride : —

3I, + 5AgF = 5AgI-hIF5.

The compound IF5 is known as iodine pentafluoride.

Comparison of the Several Acids formed by the Halogens. — We

have seen that chlorine, bromine, iodine, and fluorine all form com-
pounds with hydrogen which are acid. Taking these in the order
of the increasing atomic weight of the halogen, we have seen that
hydrofluoric acid is the most stable of all the compounds of the halo-
gens with hydrogen. Hydrochloric acid is next, and this is followed
in the order of decreasing stability by hydrobromic and hydriodic
acids. Indeed, the last named substance is quite unstable.

If we turn to the compounds with oxygen and hydrogen, we find
the order exactly reversed. Fluorine forms no known compound
with oxygen. Chlorine forms very unstable compounds with hydro-
gen and oxygen, bromine still more stable compounds, while iodine
forms fairly stable substances when combined with oxygen and
hydrogen.

If we compare the strengths of the acids formed by the union of
hydrogen with the several halogens, we would find that they were
all acids, but to a very different extent. The method of determin-
ing the relative strengths of acids is to determine the relative num-
ber of hydrogen ions in their solutions, since acidity is due entirely
to hydrogen ions. To determine the relative number of hydrogen
ions is the same as to determine the degree of dissociation of the
several substances. The dissociation of a compound is most readily
determined, as we have seen, by measuring the conductivity of its
solutions, and also the conductivity of its completely dissociated
solution. The molecular conductivity at any dilution is known as
ft^, the molecular conductivity at complete dissociation, yi.^. The
dissociation a is the ratio between these two quantities. The molec-
ular conductivities of a number of solutions of the compounds of the
halogens with hydrogen are given below. The dilutions in the four

170

PRINCIPLES OF INORGANIC CHEMISTRY

cases are the same, being the number of litres of the solution which
contain a gram-molecular weight of the electrolyte : —

Dilution

HCl

HBr

HI

IIF

¥

MK(t»°)

ILy{<afi)

Mv(.>50)

Mk

4

343

354

353

27.8

82

360

373

372

55.8

128

376

380

380

08.2

250

378

380

380

129

M»

378

880

381

380

The dissociations in each case are calculated by dividing the
molecular conductivities at any given dilution by the maximum
molecular conductivity of the substance.

Dilution

na

HBr

HI

HF

«

«

«

«

4

90.8%

93.2%

92.7%

7.3%

32

97.6

98.2

97.6

14.7

128

99.5

100.0

99.7

26.0

256

100.0

100.0

99.7

84.0

While the strengths of hydrochloric, hydrobromic, and hydriodic
acids are almost exactly the same at all dilutions, that of hydroflu-
oric acid is very much less. Indeed, at the greater concentrations
hydrofluoric acid is less than one-tenth as strong as the other halo-
gen acids. This is one of many examples of physical chemistry cor-
recting erroneous conceptions in inorganic chemistry.

CHAPTER XII

SULPHUR (At. Wt. = 32.06)

Occurrence and Purification. — Sulphur occurs in great abundance
in nature in the free condition. This is especially true in volcanic
regions such as those of Italy, Sicily, Iceland, China, etc. A vol-
canic region in which the deposition of sulphur is still going on is
known as a solfatara. Sulphur also occurs in combination with a
number of other elements. In combination with oxygen as sulphur
dioxide, SOg, it escapes in certain volcanic regions. Combined with
hydrogen as hydrogen sulphide, H^S, it also issues from the earth
in the neighborhood of certain volcanoes. It also exists in combina-
tion with a number of metals as sulphides. We have zinc sulphide
or zinc blendCy ZnS, lead sulphide or galena, PbS, iron sulphide
or pyrites, FeSj, mercury sulphide or cinnabar, HgS, antimony sul-
phide or stibnite, SbjSs, and copper iron sulphide or copper pyriteSy
CujFejS^.

The sulphur which occurs in the free condition in nature does
not all escape from the interior of the earth in the form of sulphur,
but is deposited as the result of the action of one sulphur compound
on another, or of oxygen on hydrogen sulphide : —

2H,S-h02=2H,0 + 2S,
SO, + 2H,S = 2H,0-|-3S.

In this last reaction sulphur dioxide, which generally takes up oxy-
gen and is, therefore, a reducing agent, gives up oxygen and is an
oxidizing agent.

Sulphur as it occurs in nature contains more or less impurities
which are generally non-volatile. The sulphur is freed from these
by fusion. The first more or less crude product of molten sulphur
is known as crude brimstone. The crude brimstone is then redis-
tilled and either condensed in a state of fine division as flowers of
sulphur, or the molten mass poured into moulds as roll or stick
sulphur.

Chemical Properties of Sulphur. — Sulphur at ordinary tempera-
tures is comparatively inert. Indeed, at elevated temperatures it is

171

PRINCIPLES OF IKORGAFIC CHEMISTRY

less active chemically than members of the halogen group at ordi-
nary tetiiperatuiea* When heated in the presence of the air sulphur
combines with oxygen, forming sulphur dioxide i —

Sulphur, however, combines with many of the elementSj forming
well-detined comjjounds. The compounds with the acid- forming ele-
ments, such as those wliich we have already stndied, are, in general^
much ]e9S stable thau the compounds of sulphur with the metals.
The latter class of compounds^ known as the sitlphtdesj have charac-
teristic properties whieh are very useful, as we shall learn, in qualit4i*
live analysis.

Physical Properties of Sulphur, — Sulphur is a yellowish solid at
all ordinary temperatures, inciting at US'* and boiling at 448M,
Tlie solid sulphur is known in two forms. If allowed to crystallize
from a solution of carbon di sulphide, and as found in nature, it cry**
tallizes in the orthorhombic system; the characteristic of this system
being that the three crystallographic axes are all at ri^'ht angles and
all of unet|iial length.

If, on the other hand, ordiuary flowers of sulphur, roll sulphur,
or orthorhombic sulphur is melted and allowed to cool slowly in a
hessiau crucible, we obtain the sulphur in the form of needles which
do not belong at all to the orthorhombic system, but to a crystallo-
graphic system having a much lower order of symmetry — the mono-
clinic system. The characteristic of this system is that the three
crystallographic axes are all of unequal lengths^ and one of them
makes an obliqne angle with the other two.

Substances which can crystal] ize in several systems are known as
poiifmorphoii.^ ; when they crystallize in two systems, as dimorphoits.
Sulphur i s, therefore, dimQrphons, Below 9o^*(j orthorhombic sulphur
is the stable phase, while from 9/J^C to 131** nionoclinic sulphur is
the stable phase. This point 95^6^ at which the transformation from
one solid phase to the other solid phase takes place, is known as the
tmnsUfon [mnL Substances which like sulphur exist in two phases
of the same stiite of aggregation, and the two phases can be recipro-
cally transformed into one another by changing the temperature, are
knowTi as enantmiropk. Where only one form is st^ible undc? condi-
tions which can be realized and, consequently, the unstable form can
be transformed into the stable but not mce vermis we have what is
known as monotropisni, ^

In addition to the above two solid modifications of svdphur, we
have solid, amorphous sulphur, exempUfled by flowers of sulphur.

SULPHUR • 173

milk of sulphur, etc., which are insoluble in carbon disulphide and,
therefore, differ from crystallized sulphur. The differences between
the various solid phases of sulphur are far more deep-seated, as we
would expect, than mere external form. Take the two crystalline
modifications. They are obviously analogous to oxygen and ozone,
the former being, however, enantiotropic, the latter monotropic, but
this difference is not fundamental, simply depending upon whether
the transformation point is below the melting-point.

In the case of oxygen and ozone we have seen that the f imda-
mental and important difference is in the amount of intrinsic energy
present in the molecules of the two substances. We would naturally
ask whether any such difference exists between orthorhombic and
monoclinic sulphur. This can be answered very simply in the case
of sulphur, indeed more simply and directly than with oxygen and
ozone. It is only necessary to burn equal quantities of the two
modifications of sulphur in oxygen and measure the amounts of heat
liberated. The products being the same in both cases — sulphur
dioxide — any difference in the heats of combustion is the expression
in thermal units of the difference between the intrinsic energy in
orthorhombic and monoclinic sulphur.

A considerable difference was found in the amounts of heat liber-
ated in the two cases. Thirty-two grams or a gram-atomic weight
of orthorombic sulphur when burned to sulphur dioxide liberate
71,000 calories of heat. An equal weight of monoclinic sulphur
burned to sulphur dioxide sets free 73,300 calories of heat. The
difference, 2300 calories, is the thermal equivalent of the difference
in the intrinsic energy of the two modifications.

The relations are just what we would expect from our study of
oxygen and ozone. Oxygen is the more stable form under ordinary
conditions, and contains the smaller amount of intrinsic energy. So,
also, orthorhombic sulphur is the more stable form under ordinary
conditions and contains less intrinsic energy than monoclinic.

Sulphur, as has been stated, melts at 118**. It first passes over
into a thin, light-yellow liquid, which, on further rise in temperar
ture, passes through a remarkable series of transformations. When
heated to 160** the yellow liquid passes over into a reddish-brown,
viscous mass, which becomes deeper brown in color and more viscous
as the temperature is raised to 250**. If the temperature is still fur-
ther raised until 400** is reached, the viscous mass becomes a yellow
liquid again, which boils at 448**.4. When the boiling sulphur is
allowed to cool, it passes through these same changes again, but in
reverse order.

174 PRINCIPLES OF INORGANIC CIIEMrSTRY

The vapor of sulphur when formed at its boiltiig temperature
\b reddiah-brown^ but this beoornes much lighter in color m the tem-
l>eratiire is rLiiseiL

Vapor-density of Sulphur.— The vapor-density of sulphur has
attracted much attention, since it vaiies so greatly with the tem-
perature. If sulphur is boiled under diminished pressure, so that
the teniperature is quite low, its vapor-density corresponds to the
molecular weight 2o6. Since the atomic weight of sulphur is 32,
the molecule of the vapor under these conditions consists of eight
atoms — Sf^ If sulphur is boiled under ordinary atmospheric press-
ure, the vapor-density corresponds to a molecular weight of 102,
which means that the molecules are composed of six atoms each —
Sj, If the vapor of sulphur is heated to S00° its vapor-density
corresponds to a molecular weight of 70, while if the vapor is
heated to lll>0% its density shows a molecular w^eight of 64* This
corresponds to the molecule %

As the teniiJerature rises the complex molecules of sulphur break
down ujto simpler molecules, and when a temperature of 1100'' is
reached, practically all of the more complex molecules have broken
down into molecules containing two atoms each.

The further question arises, Do the molecules Sg break down in
stages or do thej decompose at once into molecules of S^? This has
recently l>een answeretl satisfactorily by methods which it ivould
lead lis too far to discuss here. The molecules of Sg do not first
break down into molecules of S^, S^, etc.j but decompose at once into
molecules of Sj, in tlie following sense ; —

Sjj ^ 4 Sj.

The opposite opinion was held for a time, but is undoubtedly
erroneous.

The Temperature-prestnre Diagram of Sulphur. — If we plot the
tern peratmt*- press lire diagram of sulphur as we did that of water, it
would have the following form (Fig. 2T): —

The diagram is considerably more complex than the diagram for
water, whtjru only three phases were present; yet the principles
involved are exactly the samej and if we understood the diagram
for water, this shotjld offer no serious difficulty*

Hoginning with the conditions of equilibrium between or thorhom-
bic sulphur and sulphur vapor, these are represented by the curve
PJh Tlie curve PP^ is the vajx^r- pressure curve of monoclinic sul-
phur, while /\C7 is the vapor-pressure curve of liquid sulphur. The
point P it the transition point of orthorhombic and monoclinic sul-

SULPHUR

175

phur. The curve PPn represents the conditions of equilibrium
between orthorhombic and monoclinic sulphur, and any point on
this curve is therefore a transition point. The curve PiPu repre-
sents equilibrium between monoclinic and liquid sulphur, and is
therefore the curve of the melting-point of monoclinic sulphur.
Just as the curve (PPn) of the transition point of orthorhombic and
monoclinic sulphur slopes to the right as it rises, showing an in-
crease in temperature with increase in pressure^ so the curve of the

TEMPERATURE

Fio. 27.

melting-point of monoclinic sulphur (PiPu) slopes to the right as it
rises. This is but one of many analogies between transition points
and melting-points. These two curves, however, meet at the point
Pu, which corresponds to a temperature of 131®. The curve P^E
is the curve of equilibrium between orthorhombic and liquid sulphur,
I.e. the curve of the melting-point of orthorhombic sulphur with
increase in pressure, monoclinic sulphur being incapable of exist-
ence beyond 131**, no matter how high the pressure.

Let us turn now to the dotted curves. PA represents the vapor-
pressure of metastable monoclinic sulphur. This is greater below

176

PRINCIPLES OF IKOHGANIC CHEMISTRY

tlie transition point, as we would expect, than the vapor-pressure of
the stable ortliorhombic phase. Alxjve the transitiun point ortbo
rhombic sulphur is the metastable phase, and it lias in this region
a higher vapor -pressure than the stable monoelinic phase. This is
represented by the curve PPmt the prolongation of PB. If now we
prolong the curve, I\0 representing equilibrium between liquid sul-
phur and its vapor until it meets the prolongation of PB^ it will do
so at Pu,. If now we join Pn, and P^y the curve will represent the
equilibrium between orthorhombie sulphur and liquid suljihur, i.e.
the melting-point of orthorhombie sulphurj and the elf ect of pressure
as increasing the temperature at which this phase will laelt

We have now examined all the curves in the diagrjun. Let us
see what kinds of systems they represent. The point P repre-
sents equilibrium between the three phases orthorhomhiCj mono-
clinic, and vapor, and is, therefore, a triple point. Similarly, Pj
represents equilibrium between monoclinic, vapor, and liquid j Pn,
between orthorhombie, monoclinic^ and liquid j and Pju (in the meta-
stable region), betweeu orthorhombie, liquid^ and vajxjr; and these
are all triple points. We have^ then^ four triple pointsj and since
there is one component and three phases the systems are nonvariant

Take the curves. PB represents equilibrium between orthorhom*
bic and vapor, PP^ between monoclinic and vapor, P^C between liquid
and vapor, PiP,i between monoclinic and liquid, PnP between ortho-
rhombie and monoclinic.

Take the dotted curves representing equilibria in metastable
regions. P-:l is the curve of equilibrium between monoclinic and
vapor, PPiji between orthorhombie and vai>or, PjPju between liquid
and vapor, and PuPm between orthorhombie and liquid.

These systems i-epreseut conditions of equilibria between two
phases, and since the number of components is one they are mono-
variant systems.

Take finally the areas. Within BPP^C sulphur is stable only in
the form of vapor, within CPiP^E the liquid is the st-able form,
within EPiiPB the orthorhombie is the stable phase, and within
PPiPn the monoclinic is the stable form. These areas eatdi repre-
sent one stable phase of the substance, and since there is only one
component these systems are di variant.

So much for the conditions of equilibria where there is one com-
ponent and four phases.

Compounds of Sulphur with Hydrogen. — Bidphur forms one very
important compound with Ijydmgen — hydrogen sulphide, HjjS^ and
on# which is far less important — hydrogen persulphide, llfS^

SULPHUR 177

Hydrogen Sulphide^ H^S. — This compound occurs in nature in the
free condition, as we have seen. It escapes from fissures in certain
localities and is dissolved in certain waters, producing sulphur water.
Hydrogen sulphide is formed by the direct union of hydrogen and
sulphur. When hydrogen and sulphur vapor are heated together,
especially in the presence of porous porcelain, they combine in part,
forming hydrogen sulphide. If nascent hydrogen is brought into
contact with sulphur, there is a certain amount of hydrogen sulphide
formed.

By far the best method of preparing hydrogen sulphide, and the
one which is always employed, is to treat certain sulphides with an
acid. The sulphide which it is most convenient and economical to
use is ferrous sulphide, FeS. When this is treated with sulphuric
or hydrochloric acid the following reaction takes place : —

FeS + HjSO^ = FeSO^ + H,S,
FeS + 2 HCl = FeCl, + HjS.

Chemical Properties of Hydrogen Sulphide. — Hydrogen sulphide
is acted upon by ceiiain metals at ordinary temperatures. Thus,
silver and lead decompose it, combining with the sulphur and
liberating the hydrogen.

The oxides of certain metals also react with hydrogen sulphide,
forming the sulphide of the metal and water.

By far the most important chemical application of hydrogen
sulphide is in connection with its action on the soluble salts of the
heavy metals. Take as an example its action on solutions of silver
nitrate : —

2 AgNO, + H,S = Ag,S + 2 HNO,.

In such cases we have the sulphide of the metal precipitated with its
characteristic properties.

It is generally true that when hydrogen sulphide is passed into
solutions of salts of the heavy metals, the sulphide of the metal is
precipitated. If this does not take place otherwise, it is effected by
making the solution slightly alkaline.

The sulphides of certain metals resemble one another with respect
to a given property. This enables us to separate the metals into
groups by means of hydrogen sulphide. The sulphides of certain
metals are soluble in water. Thus, the sulphides of sodium, potas-
sium, ammonium, caesium, lithium, rubidium, calcium, barium, stron-
tium, magnesium dissolve readily in water, and these elements as a

178

PRINCIPLES OF INORGANIC CHEMISTBr

group can be separated from all other elements. Other means musir
be employed to separate the individual membera of this group from
one another*

There is another group of elements whose sulphides dissolve in
dilute acids. These include among the more common elements zinCj
manganese, uranium, iroo^ cobalt, niekeL

These can be precipitated as a group, not by hydrogen sulphidej
since this would necessitate the formation of free acid and the con-
sequent solution of the sulpbidej but by a soluble sulphide. The
sulphide employed is ammonium sulphide, and this group of elements
is known as the ammonium sulphide group* If we add ammonium
sulphide to a solution containing salts of all the above elements,
they would all he precipitated together.

In order to separate the several members of the group from one
another, individual differences in one property or another of the
sulphides must be utilized. The reaction expressing the precipita-
tioa of a sulphide by means of ammonium sulphide is —

ZnCl, + (NH,)^ = ZnS + 2 NH^CL

It is obvious that there is no acid set free in this reaction.

There remains a group of elements whose sulphides are not solu-
ble in cold, dilute acids. These include arsenic, antimony, tin, plati-
num, gold, mereuiy, cadmium, copper, silver,- lead, and bismuth.
Salts of these metals are readily precipitated by hydrogen sulphide,
since the acid set free does not dissolve the sulphide when formed
Thus: —

CdCl, + HtB = CdS + 2HCl,

PtCl, + 2 HjS ^ PtS^ + 4 HCl,

CUSO4 -h HjS =t CuS + HjSO^.

This gi*oup of elements is known as the hydrogen sulphide group.
Here again individual differences between the sulphides are utilized
to separate the several members* Thus, the sulphides of the first
five elements are soluble in yellow ammonium sulphide, and they
are thus separated from the remaining sulphides.

Hydrogen sulphide is easily oxidized in the sense of the follow-
ing equation : —

2H«S + 0, = 2H,0 + 2S.

\Yhen hydrogen sulphide in water is exposed to the air, the above
reaction takes place and the sulphur is precipitated as a white

SULPHUR 179

powder. Because of the ease with which hydrogen sulphide reacts
with oxygen it is a good reducing agent.

Where hydrogen sulphide is passed through a tube heated to red-
ness it decomposes in part into hydrogen and sulphur. This is very
surprising when we consider that hydrogen sulphide is formed by
the direct union of hydrogen and sulphur when the two are heated
together. These statements seem directly contradictory. This
brings us to consider a new phase of chemical reactions which we
have thus far not taken up at any length.

Seversible Chemical Eeactioiui. — We have regarded chemical reac-
tions thus far as taking place only in one direction. Two substances,
A and J5, unite and form the compound AB^ and we have written the
equation expressing the reaction in the following manner : —

A^-B^AB.

This regards the compound AB as the static or unchanging con-
dition into which the elements A and B have passed when they
combine.

This is frequently not the whole truth. The compound AB
often undergoes decomposition at the same time that it is being
formed, giving again the elements A and B. This reaction would be
represented as follows : —

AB^A^-B,

The reaction which took place originally between the elements
A and B is exactly reversed in the second stage. Such reactions, of
which there is an unlimited number, are known as reversible reactions.
Indeed, some are of the opinion that all chemical reactions are
reversible, the original reaction in some cases proceeding, however,
very rapidly, while the reverse reaction proceeds very slowly. This
gives us the key to the formation of a certain amount of the sub-
stance AB from the elements A and B when the reaction is revers-
ible. At first we have only the elements A and B, These begin
to combine and fonn the compound AB with a certain velocity which
is at first very considerable, but which becomes gradually slower and
slower as the amounts of A and B become less and less. At first
there is none of the compound AB present, only the uncombined
elements. When AB begins to form, the reverse reaction resulting
in its decomposition into A and B begins, but at first has very small
velocity. As the amount of the compound AB increases, the
velocity with which it is decomposed also increases. The result is
that the velocity of the original reaction is becoming less and lesSi

180

PRINCIPLES OF INORGANIC CHEMISTRY

while the yeloeitj of the reversed reaction is hecoming greater and
greater* After a time the two velocities become equal and we have
then the condition described as equilibrium^

When eqnilibriuia is reached, it does not mean that the
original reaction has ceased or that the reversed reaction has
ceased, but that the two are taking place with the same velocity,
just as much of the compomid AB decomposing in a given time
as is formed in a given time. This is the same as to my that
the condition of equilibrium is not a static condition as was for
a long time supposed, but is a dpiamic condttion^ The impor-
tance of the recognition of this fact is very great indeed. It
underlies the entire chapter of chemical dynamics and equilib-
rium, which is one of the most important in modern physical
chemistry.

Let OS apply this conception to the reaction under consideration.
Hydrogen and sulphur combine forming hydrogen sulphide, with a
velocity which becomes less as the quantity of the elements present
decreases. Hydrogen sulphide decomposes into hydrogen and sul-
phur, with a velocity which becomes greater as the amount of
hydrogen sulphide present increjises j after a time just as much
hydrogen sulphide heing decomposed in a given unit of time as is
formed in the same time. Equilibrium between the two reactions
is then established.

When equilibrium is established we have the maximum amount
of hydrogen sulphide formed, which, under the conditions ever
could be formed. This is usually referred to as the $klti of f/ie re-
action.

All chemical reactions in which the yield is less than one hun-
dred per cent are reversible, and since this theoretical yield is
probably never quite fully realized^ all reactions aie probably
strictly speaking reversible.

In some cases, however, the combination is so nearly com-
plete that we must regard the velocity in one direction as infi-
nitely great with respect to the velocity in the reverse direction.
In such cases we would have, when equilibrium was reached,
nearly all of A and B combined to form the compound AB^ while
a very slight amount of AB was decomposed into ^1 and B. This
is the condition which obtains in most reactions where a solid is
precipitated. The solid is formed with a velocity which is usually
far too great to measure, and the reaction proceeds nearly to the
end before equilibrium between the two opposite reactions is estab-
lished

SULPHUR 181

Only such reactions can be used in quantitative analysis which
depend upon the reaction in one direction being practically com-
plete.

Acid Sulphides. — Hydrogen sulphide, as we have seen, has the
composition HjS, and forms salts with univalent elements having the
composition M2S, M representing any univalent element. If M repre-
sents a bivalent element, the salt has the composition MS.

Hydrogen sulphide can also form a different class of salts in
which one of the hydrogen atoms is still present. With univalent
elements these would have the general composition MHS, with
bivalent elements M(HS)j. The acid in these salts is monobasic —
one molecule of the acid, as we usually express it, combining with
one ion of a univalent element. We have a number of examples of
these hydrosulphides or acid sulphides as they are called. Ammonia
forms the hydros ulphide very readily when hydrogen sulphide is
conducted into aqueous ammonia : —

NH3 -}- H,S = NH4HS.

Indeed, this is the compound which is formed when aqueous
ammonia is saturated with hydrogen sulphide. In order to form
the normal sulphide, an equal quantity of ammonia must be added
to the hydrosulphide : —

NH4HS + NH8 = (NH,)2S.

Many hydrosulphides are known, such as NaHS and KHS.
Indeed, those metals in general, whose sulphides are soluble in water,
form hydrosulphides.

Dissociation of Hydrogen Sulphide. — The hydrosulphides or acid
sulphides are, as we have seen, salts of a monobasic acid. The
sulphides, on the other hand, are salts of a dibasic acid, i.e, one which
combines with two univalent ions. How can we explain these facts
on the basis of the dissociation theory ?

There are two different ways in which hydrogen sulphide can
dissociate. These are the following : —

H^ = H, HS, (1)

H^ = H, H, 8. (2)

When the compound dissociates as in equation (2), it is obviously
dibasic ; when the dissociation takes place as in equation (1), it is
monobasic. This is the general method by which dibasic acids dis-

t8t MaXCirtES OF IKORGANTC CHEmSTRY

a»ucial9w T^ftke m a general example H,A, in which A is the anloii of
IM ilBMmti MritL This dissociates first into H and HA : —

H,A ^ H, HA.

Wilt MtOfl HA coii tains the anion proper combined with a hydrogen
ilQIlL If the dilution of the solution is still further increased, this
biMyki down into Hj A : —

HA ^ H, A.

When the dissociation has taken place in the sense of (2), the acid is
iKUnpletely dissociated, and we have a dibasic acid.

The amount of the dissociation in terms of (1) or (2) at any given
dilution is determined solely by the nature of the acid. If the acid
ia very strong it begins to dissociate in terms of (2) before any very
great dilution is reached, while if the acid is weak it may re<]uire
very high dilution to effect any appreciable amount of dissociation in
the sense of the second equation.

The salts of acids such as we have been considering dissociate
very much like the acids themselves. The hydrosulphide or acid
sulphide of sodium^ NaHS, dissociates thus : —

mHS = I^a, HS.
The anion HSj being the anion of a we-ak acid, diaBOciates slightly

into lU ^t h^it only slightly, A fpw htidroffen ions are thus formed
in the solution of sodium hydrosulphide or acid sulphide, and these
give the characteristic acid reaction of this salt with indicators. The
mud reat^tion is, however, weak, since there are relatively only a few
hydrogen ions present in the solution. If suhsliances like hydrogen
sulphide dissociate as monobasic acids yielding only one hydrogen
ion, why do they reacrt like dibasic acids when treated with a solu-
tion of a strong base like sodium hydroxide ? The answer t^ this
apparently difficult question is very simple. The sodium hydroxide
reacts with all the dissociated portion of the hydrogen sulphide
thus r —

H^S, H + OH, Na = H,0 + ^a, HS,

All the hydrogen ions originally present are, therefore, combined
with the hydroxy! anion to form water and removed from the field of
action* As sfKrn as this has taken place, we have the remaining
undiasociated portion of the acid under the same conditions as the
original acid — in the presence of water in which there are no hydro-

SULPHUR 188

gen ions. The dissociation will then proceed until all the molecules
have been broken down in the sense of (1). If more sodium hydroxide
is now added, we would have in the solution an excess of hydroxyl ions,

and these will cause the anion HS to dissociate to a much greater
extent than it would do in the presence of water alone ; since as fast
as hydrogen ions are formed they combine with hydroxyl ions from
the sodium hydroxide, form water and are removed from the field of
action. There is thus no accumulation of hydrogen ions as in the
case of water alone, and the dissociation proceeds until all the ions

HS have dissociated into H, S. Hydrogen sulphide therefore acts
as a dibasic acid towards a strong base.

Physical Properties of Hydrogen Sulphide. — Hydrogen sulphide
is a colorless gas with an extremely disagreeable odor, suggesting
decomposing organic matter. A\Tien taken into the lungs in any
great quantity, it is quite poisonous. For this reason and on account
of its disgusting odor it should never be allowed. to escape into the
air of a laboratory. Since it is so frequently used in connection
with qualitative analysis, a separate room is attached to every well-
equipped chemical laboratory in which the gas is generated and
used. This is known as the " sulphuretted hydrogen " room.

Hydrogen sulphide dissolves in water to the extent of from 2^ to
3 volumes of the gas in one volume of water. Even at such concen-
trations the law of Henry holds — the amount of gas dissolved is
proportional to the pressure to which the gas is subjected. The law
of Henry, in general, holds- better the more dilute the solution, i,e.
the less the solubility of the gas.

A solution of hydrogen sulphide in water soon becomes cloudy,
due to the deposition of sulphur, the hydrogen sulphide being
oxidized, as we have seen, by the oxygen of the air, yielding water
and sulphur.

When hydrogen sulphide is subjected to a pressure of fifteen
atmospheres at ordinary temperatures, it passes over into a colorless
liquid having a specific gravity of 0.9. When cooled to — So^, it
solidifies. Liquid hydrogen sulphide is somewhat explosive.

Hydrogen Persnlphides. — There are a number of compounds of
sulphur with potassium which contain much more sulphur than
potassium sulphide — KjS. These are K2S2, KgS^ K2S4, and K2SJ.
These are formed by the union of potassium sulphide with sulphur.
When an ordinary solution of ammonium sulphide is allowed to
stand in contact with the air for a time, a part of it is oxidized with
liberation of sulphur, which then combines with the ammonium sul-

184 PHINCIPLES OF INOriGANIcf CHEMISI

phide forming poljsiilphides of ammonia. This h shown by the
change from tho colorless ammonium sulphide to the deep-yellow
poly sulphides of ammonia*

When any of these polysulphides is treated with an acid, it gives
off hydrogen snlphidej and snlphnr is set free. When^ however^ the
process is reversed and the solution of the poly sulphide added to
the acid, no hydrogen sulphide is formed, but a yellow, oily liquid
separates. The couiposition of this liquid probably depends on the
particular polysulphide which is present It is probably the acid
corresponding to the polysulphide in questiom The fact is that we
almost always have a mixture of a number of these polysulphides,
and when we pour such a mixture into an acid, the resulting oil
probably cont^ius a number of hydrogen persulphides, ranging in
composition from H^Sj to HSi-

This whole question of the composition of the hydrogen persul-
phides is, however, still an open one.

COMPOUNDS OF SULPHUE WITH OXYGEN AND HYDROGEN

Sulphur Dioxide, SOj. — The simplest compound of sulphur and
oxygen is sulphur dicjxide, having the composition SOg, It is formed
by the direct combination of the two elements when sulphur is burned
in oxygen : —

S + O, = SO*.

A more convenient method of preparing sulphur dioxide is by
the action of strong acids on sulphites, or acid sulphites. Xormal
potassium sulphite has the composition KjSOg. When this is treated
with sulphuric acid the following reaction takes place: —

K^SO, + H^O* = K^SO, + H,0 + SO,.

The a^id sulphite has the composition KHSOj. When this is treated
with sulphuric acid we have : —

2 KHSO, + H^0| = K^SO^ + 2 H,0 + 2 SO^

Another very convenient method for preparing sulphur dioxide
is by the action of sidphuric acid on metallic copper. The reactioa
may be regarded as taking place as follows : —

Cn + H3SO4 = CuSO^ + n^

The nascent hydrogen then reduces the sulphuric acid: —
H,S04 + H,^2H,0 + SO>.

SULPHUR 185

Its most characteristic chemical properties are its oxidizing and
reducing actions. Under certain conditions it gives up its oxygen,
serving as an oxidizing agent ; under other conditions it readily takes
up oxygen, passing over into sulphur trioxide, SOg, and is, therefore,
an excellent reducing agent.

The composition of sulphur dioxide is determined by the follow-
ing considerations. When sulphur is burned in oxygen the volume
of the sulphur dioxide formed is just equal to the volume of the
oxygen used up. From Avogadro's law there are, therefore, the
same number of molecules of sulphur dioxide formed as there were
molecules of oxygen used up. Each molecule of oxygen, however,
contains two atoms of oxygen ; therefore, each molecule of sulphur
dioxide must contain two atoms of oxygen.

It is a colorless gas with very penetrating odor, and a taste which
persists for an unusual time. Its odor and taste are characteristic
of a burning sulphur match, with which every one is familiar. It
does not obey the laws of gas-pressure, being too near its point of
liquefaction. Its critical temperature is 157®, and its critical pressure
is 79 atmospheres. It can, therefore, be readily liquefied. At ordi-
nary temperatures it is liquefied if subjected to a pressure of three
atmospheres. If cooled by a mixture of salt and ice it readily lique-*
fies under atmospheric pressure. Liquid sulphur dioxide is perfectly
clear and transparent, boiling at —10®. It can be readily solidified
when allowed to evaporate imder diminished pressure, a temperature
of — 50® to — 60® being produced. Liquid sulphur dioxide is an
excellent solvent for ajarge number of substances, and according to
the recent work of the Russian, Walden, has considerable power to
dissociate electrolytes into their ions. Indeed, solutions of certain
salts in liquid sulphur dioxide frequently show better conductivity
than solutions of the same salts at the same concentrations in water.
It is readily obtained on the market in steel cylinders. Sulphur di-
oxide dissolves readily in water, one volume of water at ordinary
temperatures dissolving about fifty volumes of the dioxide. The
solution of sulphur dioxide in water has an acid reaction and is
known as sulphurous acid.

Sulphurous Acid, H2SO3. — The acid formed when sulphur diox-
ide is dissolved in water yields salts having the composition MjSOj,
where M is a univalent metal. The acid must therefore have the
composition HjSOs, and be a dibasic acid. It can also form acid sul-
phites of the composition MHSO3.

It has been impossible to isolate the acid H2SO3, since it breaks
down so readily into water and sulphur dioxide.
H^03=H,0 + S0,.

:-'_i J 3'3i>.=-L"z:

"" --»■* zaji. xjir -mr cus-

Ire. •:•—=- r "Lr- J.— :j-l i

' >*^ • 7->.v.te v: —

^"^-.stlI 111* »' 'TTrnn

ff I

' ' /■•'•■".'■ . :. i.r ■ : i :- "lie rcw-

' "' *'"';^•■'■'. /•! U'-:tU',\ in thr :r»^s.>nce

"" / ".H.r„n. ,*.,/),Jy, foTunir^ Milphur tri.

' '■ '• • » I.

'' »■ ■ '» '.'... ,., f,. ,„,..^ ,,,.,, Htiil U,tter, finely

SULPHUR 187

divided platinum. If the two gases are passed through a tube con-
taining platinum sponge, and the tube heated, they combine very
readily. Instead of using platinum sponge, it has been found to
be more economical to use asbestos covered with finely divided
platinum — platinum asbestos.

This is distinctively a catalytic reaction^ the platinum in this case
does not enter at all into the reaction, and a very small amount of
platinum is capable of effecting a large amount of the reaction. This
method of preparing sulphur trioxide has been foimd to be so effi-
cient and economical that it is rapidly supplanting all others, and
in the near future will probably be used almost exclusively for pre-
paring this substance. The sulphur dioxide obtained by roasting
various sulphur ores, and especially pyrite, is mixed with air and
passed over heated platinum asbestos. The sulphur trioxide is then
dissolved in concentrated sulphuric acid, forming the so-called solid
sulphuric acid, having the composition HjSjOy.

Properties of Sulphur Trioxide. — Sulphur trioxide is a powerful
oxidizing agent, readily giving up one of its oxygen atoms and passing
over into sulphur dioxide. It has an unusual attraction for water,
combining with it at once on mere contact, and even causing the
hydrogen and oxygen in organic compounds to combine and form
water with which it instantly combines.

Sulphur trioxide is a transparent, mobile liquid, which boils at
46®.2. It can be cooled to zero without solidifying. When further
cooled it forms a white solid, which melts at 14**.8.

When kept at ordinary temperatures it undergoes polymerization,
passing into a white solid composed of fine needles, having the
general appearance of asbestos. This form, which is very probably
a polymer of the other, is the stable modification. When it is heated
it does not melt, but passes at once into vapor. When the vapor is
condensed it forms the liquid first referred to.

One reason for supposing that the solid is a polymer of the
liquid form, is that the latter is the stable modification at higher
temperatures where polymers tend to pass into simpler forms, and is
formed directly by condensing the vapor, in which form the molecule
is generally the simplest possible.

Sulphur trioxide dissolves in water with a crackling sound, and
with the evolution of an enormous amount of heat, forming sulphuric
acid.

Sulphuric Acid, H2SO4. — Sulphuric acid occurs, in the free con-
dition, in small quantity in certain waters on the surface of the
earth, and in abundance in sulphates. It is prepared now very

z^ xz

* " T^

.-rc-i

:.i,-T:[>fg

pa.?:-.-./ :j.. .-;- v-i. ^ ;,. . ... .- 3^^^

t*. i - • .- : .. -. : <- .... . - ^ . .7-^: hi'^

Till' r)i«'frijr;il r«'.i/*...r. - i ... ,. -./-;- - '. i - :- -* _• - ;- - — 7»r .f

Biil|iliiini- Ji'riiJ l,y f.|,i: '^x0,',i: :;.r:!. .•:. ^i fir i? :1.tv j.r»^ k-: -s-t:. .»:«
till* I'ni lowing:

Nitiii' iM'iil (U!t.H iiji*#ri HnlfiKur 'i.ozil^r ::. t;.^ ; rosvr.oe ■. : wA:er as
follows :

SULPHUR 189

When the sesquioxide of nitrogen, NjOg, reacts with water-vapor
it forms nitrous acid : —

When nitrous acid reacts with sulphur dioxide in the presence of
the oxygen of the air, we have —

2 SO, + 2 HNO, + Oj = 2 NO,. SO3IL

The compound NO2.SO3H, is known as nitrosyl-sidphuric acid, or
nitro8ulpho7uc acid. When this is treated with water the following
decomposition takes place : —

NO,. SO3H -}- H,0 = H,S04 + HNO,.

The HNOj, or its anhydride, N2O3, then reacts with more sulphur
dioxide and oxygen and forms again NO,. SO3H, which then decom-
poses with water-vapor in the sense of the last equation, and the
process is continuous, the nitrous acid or sesquioxides of nitrogen
being collected in the Gay-Lussac tower, as we have seen.

The acid obtained from the leaden chambers is known as " cham-
ber acid." It contains about G5 per cent of the compound H2SO4.
In order to further concentrate this acid it is allowed to flow
through hot, shallow, lead pans, and when the acid has become
sufficiently concentrated to act chemically upon the lead, it is trans-
ferred to a platinum vessel and more of the water distilled off. The
acid thus obtained has a specific gravity of 1.82, and is ordinary,
commercial, concentrated sulphuric acid. The acid can be still fur-
ther concentrated in vessels of platinum.

Chemical Properties of Solphuric Acid. — One of the most char-
acteristic properties of sulphuric acid is its power to take up water
and combine with it. For this reason it is an excellent drying
agent, readily taking water from other substances. When we wish
to dry a gas which contains water-vapor, the best method with one
exception is to allow the gas to stream slowly in fine bubbles
through concentrated sulphuric acid. Its power to combine with
water is the key to many of the reactions which concentrated sul-
phuric acid can effect When brought in contact with many organic
substances, it causes the hydrogen and oxygen to combine and form
water- with which it itself combines. This is the explanation of the
charring of wood and similar substances effected by the concentrated
acid. The hydrogen and oxygen in the wood or other organic mat-
ter combine, form water which is taken up by the acid, and free
carbon remains behind as a black substance.

190 PRINCIPLES OF INORGANIC CHEMISTRY

The power of sulphuric acid to cause the elements of water to
combine is the cause of a number of chemical reactions^ where one
of the substances contains among other things hydrogen, and the
other substance oxygen and something else. The hydrogen of the
one compound and the oxygen of the other combine, and the remain-
der of the first comiK)und frequently combines with the remainder
of the second compound, giving rise to a new substance. Sulphuric
acid, thus, apparently by its contact, effects many reactions which
would not take place without the presence of a dehydrating agent.
Its reaction is, however, not catalytic, since it enters into the reac-
tion in the sense that it combines with one of the products of the
reaction.

8in(;e 8ul])huric acid has such a remarkable power to combine
with water, we would naturally ask, Does it simply mix with the
water niochanically, or does it form compounds with it? Sulphuric
acid combincH with water, forming two compounds, H2S04.H,0 and
H2S04.1^IIj,(). These compounds are usually regarded as having
the following formulas: —

and

being respectively sulphur combined with one oxygen and four hy-
droxyls, and with six hydroxy Is — the limit if sulphur is hexivalent.

There is no satisfactory evidence that sulphuric acid can combine
with a larger number of molecules of water.

Sulphuric acid has the power of driving more volatile acids out
vf th«Mr Halts, combining with the metal of the salt. Thus, when a
ilrv ohiorido or nitrate is treated with sulphuric acid, the hydro-
oUU»iio or nitric mud is driven out and the sulphate of the metal is

•J NaCl + H,S04 = Na,S04 -}- 2 HCl ;
•J NhN(\ 4- H,S04 = Na^SO^ + 2 HNO3.

V\\\ik itiik^Ut inviM to argue that sulphuric acid was a stronger acid
Uukw xMiUK^k hwUAvhlorio or nitric. Such, however, is not neces-
»^aul\ ihi» \uvto. llvdrtH'hloric and nitric acids being volatile are
di:3|il;M ahI I»> iituoU w^H^kor, non-volatile acids, in accordance with
lUv <\\ ui lal |nnio»|»to \\\i.\X i«'Anir»vra volatile compound can he formed
Hit J ' .vii^ . wvm '^. 'ki't \f/NiWioN» it isforrmd,

i tu I i \.w iiUH4.^uio A \\s» rvUtive strengths of acids is their relative

SULPHUR

191

dissociation, as we have seen ; t.e. the relative concentrations of the
hydrogen ions in their solutions. If we study the dissociation of
sulj)huric acid, we shall learn that it is much weaker than either
hydrochloric or nitric acid.

r

Mr(28«)

«

2

300.0

M.7%

32

490.0

68.7

1024

697.0

97.7

4096

713.0

100.0

8102

713.0

100.0

Sulphuric acid combines with most of the metals or base-forming
elements, forming sulphates. In all of these compounds the sul-
phuric acid is bivalent, combining with two atoms of a univalent
metal, with one atom of a bivalent metal, and so on. The sulphates
are very stable, well-crystallized compounds. The sulphates of the
heavy metals are generally only slightly soluble in water. The in-
solubility of its barium salt even in dilute acids furnishes us with a
means of detecting sulphuric acid and determining it quantitatively.
When sulphuric acid in very small quantity is added to any solu-
ble barium salt, white, insoluble, barium sulphate — BaS04 — is
precipitated.

Physical Properties of Sulphuric Acid. — Sulphuric acid, or oil of
vitriol, which is the monohydrate of sulphur trioxide (SOj.HjO
= H2SO4) is, as its name implies, a liquid. It is a thick, oily liquid
with a specific gravity of 1.84. It boils at 338®, undergoing partial
decomposition. When the vapor is heated to higher temperatures it
breaks down into sulphur trioxide and water. When sulphuric acid
is cooled below zero, it solidifies, the solid not melting imtil a tem-
perature of 10^5 is reached. When this solid is melted, it must be
cooled again to zero before it will resolidify. This is evidently
simply a case of an undercooled liquid, since, if a crystal of the solid
is added to the liquid at zero, more of the solid will separate until a
temperature of 10°.5, its true freezing-point, is reached.

Sulphuric acid dissolves readily in water with large evolution of
heat and a considerable contraction in volume.

Dissociation of Sulphuric Acid. — Sulphuric acid is a typical
dibasic acid, forming two well-defined classes of salts — the normal
sulphates and acid sulphates. The former have the composition
M,S04 and the latter MHSO4.

192

FRINCIPLES OF INORGANIC CHEMISTRY

Like dibasic acids in geueral Bulphuric acid dissociates in two
stages. At first it breaks down, thus ; —

H^80| ^ H, HSO,.

When the dilution of the solution is increased, i.e. wlien more
water is added, the ion HSO4 dissociates, thus : —

HSO^ = H, SO,.

Sulphuric acid is a strong dibaaic acid, and, therefore, the ion

HSOi dissociates into H aud SOj before any very great dilation is
reached. Sulphuric acid at a dilution of from 1000 to 2000 litres,
I.e. in solutions containing a gram-molecular weight of the acid in
one or two tliousand litres, is completely dissociated into two hydro-
gen ions aud the ion SO4. This is shown by the fact that the mo-
lecular eoutliictivity of sulphuric acid does not increase beyond these
dilutions, and is sufficiently large to show that the molecule has
dissociated into two hydrogen ions and not one.

The two classes of sulphates dissociate quite differently. The
normal sulphates dissociate as we would expect :^ —

MsS04 = M, M, SO,;

the acid sulphates thus i —

MnS04 = M, HSO4,

When the dilution is sufficient the ion HSO4 dissociates further
as follows : —

HS04=H, S0|.

A dilute solution of an acid siilpliate, therefore, contains hydrogen
ions and should react acid.

Such is the fact A solution of an acid sulphate of even such
strong base-foruiing elements as smlium and potassium, is distinctly
acid. The concentration of the hydrogen ions in solutions of acid
sulphates has been measured- There are methods for detecting the
concentration of one kind of ions in the presence of other kinds of
ions. Thus, cane sugar is inverted as we say, t.e, broken down into
dextrose and f ructrose only by hydrogen ions, and the velocity of the
inversion is a function of the concentration of the hydrogen ions
present. This reaction has actually been used to determine the con-
centration of the hydrogen ions in a solution of acid salts, where
other ions are always present.

SULPHUR 193

It is obvious that the conductivity method could not be used in
such cases, since all kinds of ions take part in conducting the current.

Scientific and Technical Uses of Sulphuric Acid. — Sulphuric acid
is used very frequently in the scientific laboratory, and far more
frequently in technical processes than any other acid. In scientific
operations it is used as a dehydrating agent, as a drying agent, to
liberate volatile acids from their compounds, and in many other pro-
cesses. In the arts it is used on every hand, and crude sulphuric
acid is manufactured by the hundreds of thousands of tons annually.
It is used to render normal phosphates soluble in water by convert-
ing them into acid phosphates, which can be assimilated by plants.
These are the basis of most of the commercial fertilizers. It is also
used in connection with the manufacture of chlorine from sodium
chloride, and in the preparation of soda. When sodium chloride is
treated with a molecular quantity of sulphuric acid, the following
reaction takes place : —

NaCl + H^04 = NaHSO^ + HCl.

The acid sulphate acts at a higher temperature upon another
molecule of sodium chloride thus : —

NaHS04 + NaCl = Na^SO^ + HCl.

Sulphuric acid is at present extensively used in connection with
the generation of electrical energy in accumulatorSy or storage cells.
In such cells the electrodes are plates of lead and the electrolyte
dilute sulphuric acid. When the electric current is passed through
such cells, a change takes place which we shall consider under lead.
If the electrodes are joined after the cell is " charged," an electric
current flows in the direction opposite to that of the charging current.

Other Compounds of Sulphur with Oxygen and Hydrogen. — The
two acids already considered, sulphurous and sulphuric, are the most
important compounds of sulphur with oxygen and hydrogen. Several
other compounds, however, are known, and these must be considered
briefly. These compounds, which are all acids, are the following : —

Thiosulphuric Acid HjSaOj

HydrosulphurouB Acid . . . . H8Ss04

Pyrosulphuric Acid HjSjOt

Persulphuric Acid HsS^Os

Dithionic Acid HtS^Oe

Trithionic Acid HjSsOe

Tetrathionic Acid HsSiOe

Pentathionic Acid H2SsO«

Hexathionic Acid HfSeOe

o

194

PEINCIPLES OF INORGANIC CHEMISTRY

TMoBulphnric Acid, KS^Oy — Salts of this acid are formed hy

boiling sulphites with sulphur^ —

Ka,SO, + S = KaaS,0»
or by the action of iodine on a mixture of sulphide and solphite.
Ka^S -h Na^SO^ +21^2 NaT + Ka^S A-

The free aeid is very unstable, existing only in dilutej aqueous
solution, and under these conditions for only a short time<

The sodium salt, which should be called sodium thioaulphate,
but which ia frequently called hyposulphite, or in the arts simply
" hyjKi,^' is important in connection with photography, It45 solution
dissolves the halogen salts of silver, and it is, therefore, used for
"fixing" photographs.

Sodium thiosulphate is easily oxidized to the sulphate, and is,
therefore, a good reducing agent It is consequently used to remove
the last traces of chlorine in bleaching, and has come to be known
as anikhlor*

When the thiosulphates are treated with a dilute solution of an
acid, the following reaction takes place : —

Na,S A + 2 HCl ^2 KaCl + B^O + SO, + S.

HydrostilphtLroiiB Acid^ SLS^O^. — The sodium salt has the com-
position NajS A * 2 H/). Therefore the acid is H.SA* The acid
and its salts &y<^ strong reducing agents.

Pyrotulphnric Acid or DinUphnric Acid, H^S^O^ — A salt of
this acid can Ije obtained by heating an acid salt of sulphuric acid: —

21vHS0, = H,0 + KA0r.

The free acid is prepared either by dissolving sulphur trioxide
in Bulphuric acid, —

HsSO4-hS03=nsSa0r,

or by heating ferrous sulphate in the preaence of water- vapor, —

4 FeSO, + H,0^2 SO, + 2 FeA+ HaS„0,.

This is known as Nordhausen mtiphuric acidj or fuming sulphuric
acid.

Periulphurio Acid, H^A- — This acid is obviously the hydrate

of sulphur septoxide, SA*

SA4-H,0=HaSA

Its salts are prepared by the electrolysis of cold, concentrated
eolutiona of sulphates.

SULPHUR 195

Most of the salts of persulphuric acid are easily soluble in water,
including even the barium salt. These are, as would be expected,
excellent oxidizing agents. Potassium persulphate dissolved in
sulphuric acid has been shown to have remarkable oxidizing
properties, and is known from its discoverer as Carols liquid. This
liquid has come very much to the front in the last few years, and
has been the basis of a number of important investigations in the
laboratory of the German chemist, Baeyer. According to him it
probably contains a substance, HjSOs, which we may call pemionoaul'
phuric acid.

FolytMonio Acids. — These include di-, tri-, tetra-, penta-, and hexa-
thionic acids, having the respective compositions, H^SjO^, HjSsOe,
H,S A, H,SA, and H,SeOe.

These compounds are formed in general by the action of iodine
in different quantities on sulphites or thiosulphates. The free acids
are in general unstable and easily decomposed.

COMPOUNDS OF SULPHUR WITH THE HALOGENS AND

OXYGEN

Compounds of Sulphur with Chlorine. — Sulphur combines with
chlorine, probably forming several compounds. One of these is a,
fairly stable substance, having the composition SgClg, and is called
sulphur monoddoride. This compound is formed when dry chlorine
gas is conducted over molten sulphur. It is a reddish-brown liquid
boiling at 137°. Its vapor-density shows that it has the double
formula S2CI2, and not the single. Sulphur monochloride readily
dissolves chlorine at low temperatures, probably forming the com-
pounds SCI2 — sulphur dichloride, and SCI4 — sulphur tetrachloride.
These compounds are, however, still somewhat in doubt.

Sulphur combines also with bromine, iodine, and fluorine.

Compounds of Sulphur with Chlorine and Oxygen. — There are
two well-known compounds of sulphur with oxygen and chlorine.
When sulphur dioxide is treated with phosphorus pentachloride, the
compound SOCl, is formed, boiling at 78°. This is known as thionyl
chloride. When thionyl chloride is treated with water the following
reaction takes place : —

SOCI2 -f. 2 H,0 = HjSOs + 2 HCl.

Since sulphurous acid is formed from thionyl chloride by the
action of water upon it, it is sometimes known as .the chloride of
sulphurous acid.

196 PRINCIPLES OF INORGANIC CHEMISTRY

When chlorine gas is allowed to act on snlphur dioxide, another
compound of sulphur with chlorine and oxygen is formed : —

S0, + Cl, = S0,Cl3^

This is known as sulphuryl chlonde, and is a liquid, boiling at 69®.
When treated with an excess of water, sulphuryl chloride breaks
down into hydrochloric and sulphuric acids : —

S0,C1» + 2 H,0 = H^SO^ + 2 HCl.

When one molecule of sulphuryl chloride is treated with one
molecule of water, the following reaction takes place : —

S0,C1, + H,0 = HCl + S0,C1(0H).

The compoimd, SOjClOH, chloraulphuric acid, is also formed by
the direct union of hydrochloric acid and sulphur trioxide : —

HCl + S0s = S02Cl(0H).

When treated with water it decomposes into hydrochloric acid
and sulphuric acid : —

SOjCl(OH) + H,0 = H2SO4 + HCl.

There is another compound of chlorine, oxygen, and sulphur
known. It is obtained by dehydrating chlorsulphuric acid by phos-
phorus pentoxide : —

2 S02C1(0H)= H,0 + SACl,.

It is known as pyrosulphuryl Monde.

CHAPTER XIII

SELXSNIUM AND TEZiLURlUM

There are two elements occurring in comparatively small quan-
tity, which closely resemble sulphur in their properties. These are
selenium and tellurium. A few of their compounds will be consid-
ered very briefly.

SELENIUM (At. Wt. = 79.2)

Selenium was discovered in 1817 by the Swedish chemist Ber-
zelius. It occurs in the same general associations as sulphur,
and frequently along with sulphur. It occurs in combination with
silver and copper as definite minerals. It is frequently found in
the dust of flues where sulphides are roasted, or in the chambers in
the manufacture of sulphuric acid. Like sulphur, selenium occurs
in more than one modification. A number of allotropic forms have
been described. If amorphous selenium is dissolved in carbon disul-
phide and the solution evaporated to crystallization, red crystals
separate, which melt at 175.** Selenium in the amorphous condition
melts at 217**. When kept at an elevated temperature, say 125® to
140®, for a considerable time, the amorphous variety becomes crys-
talline, gray in color, and has somewhat of a metallic lustre. In
this condition, known as metallic selenium, it has veiy different
properties from ordinary selenium. It is insoluble in carbon disul-
phide, and thus resembles flowers of sulphur. It differs from all
the varieties of sulphur in being able to conduct the'electric current.
The amount of its conductivity depends on the intensity of the light
to which it is exposed, varying considerably in a very short time
with the degree of the illumination to which the selenium is exposed.
It has been proposed to utilize this property of metallic selenium in
transmitting sound by means of light, and an instrument known as
the photophone has been constructed for this purpose, but has never
met with any great success.

Selenium boils at 650®, the vapor-density decreasing with rise in
temperature. When a temperature of 1400® is reached the vapor-
density becomes constant, and corresponds to a molecular weight of

197

198 PRINCIPLES OF INORGANIC CHEMISTRY

about 164. This is about twice the atomic weight of selenium,
showing that the molecule of selenium, like the molecule of sulphur,
contains at this temperature two atoms.

Compounds of Selenium. — Selenium combines directly with
hydrogen, forming the compound HgSe — hydrogen selenide — which
is strictly analogous to hydrogen sulphide. This compound is also
obtaineil when metallic selenides are treated with a strong acid.

Selenium combines with oxygen, forming selenium dioxide, the
analogue of sulphur dioxide. This is a crystalline solid, which,
when dissolved in water, forms selenious acid, of which it is obvi-
ously the anhydride. This is the only compound of oxygen and
selenium which is known.

«SW(*HioiM acid, formed by the union of selenium dioxide with

' SeOs + H,0 = Hj^eOa,

iriNti^ubloH in many respects sulphurous acid. It forms two series of
«ialU» \\\^ aoid selenites and the selenites, having the compositions,
r^jHHaivoly, MllSeOj and MjSeOs.

It ditTor», however, from sulphurous acid in not being a strong
Vi^hloiu|( a^t^nt. Indeed, it is not a reducing agent at all, but readily
$i\^ up ita oxygen, and is therefore an oxidizing agent. When
ik^t^uiouM aoid is treated with sulphur dioxide, the former is reduced
to iH»lDuiuu\, and the latter oxidized to sulphuric acid.

While the compound selenium trioxide is not known, the acid
PiUPrenpoudiuff to this anhydride is known. When selenium is
Ireateil with ntrong oxidizing agents, such as chlorine or bromine
waWr» iM' metallio selenites treated with bromine or fused with
)HiUitAiuu\ uitmto, »elenic acid or its salt is formed. The acid is a
*olid, uuiltiug at A8*. Like sulphuric acid it combines with water,
(lU'iuiiig a hydrate, H,Se04.H20. Unlike sulphuric acid it is a strong
oxiiU^iug ngeut, readily giving up its oxygen and passing over to
neleuiim^i uoid k\v Helen ium.

Htdeuiuui eiuubiueM with chlorine, forming two chlorides; selenium
uuauK^hlmide, Siv,('l„ and selenium tetrachloride, SeCl4. The latter
iti a iHuu|uiratively nUible aul)stance and thus differs from the corre-
lipoudiug ehlitvide of sulphur. Selenium combines with sulphur,
fvuuuiug the e(UU)Hmml SeH^

TKUaMlirM (At. Wt. = 127.6)

Tellurium. — Tellurium is a much rarer element than selenium,
VH^euri'iug ooiubiued with such metals as lead, bismuth, silver, and

SELENIUM AND TELLURIUM 199

gold. Tellurium forms grayish- white crystals which resemble a
metal. It conducts electricity and thus resembles one modification
of selenium. Its melting-point is about 450°. Its boiling-point is
1400°. Its vapor-density shows a molecular weight which is twice
its atomic weight. At this temperature there are, therefore, two
atoms in the molecule, in this respect resembling sulphur and
selenium.

Compounds of Tellurium. — Tellurium combines with hydrogen,
forming hydrogen telluride, having the composition H^Te. This is
analogous to hydrogen sulphide and hydrogen selenide. It is a gas
with a very disagreeable odor like the former compounds.

Tellurium combines with oxygen, forming the compounds TeO,
Te02, and TeOs. The last two are analogous to sulphur dioxide and
sulphur trioxide, while the first has no analogue among the sulphur
compounds. These oxides, however, show very little tendency to
combine with water, and thus differ markedly from the correspond-
ing oxides of sulphur and selenium.

Tellurium, however, forms two acids with hydrogen and oxygen.
These are tellurious and telluric acids, having the compositions, re-
spectively, HjTeOs and HxTeOi.

Tellurium, unlike sulphur and selenium, also shows certain basic
properties. It forms with nitric acid a basic nitrate, and thus differs
fundamentally from sulphur and selenium.

Tellurium can also combine with chlorine and bromine, forming
the compounds TeClj and TeCl^, and TeBr, and TeBr^

CHAPTER XIV

NITROaEN (At. Wt. = 14.04)

We now pass to another group of elements, the nitrogen group.
This is group V in the Periodic System. The members of the
nitrogen group are nitrogen, phosphorus, arsenic, antimony, and
bismuth. We shall first take up nitrogen and study it in some
detail on account of its importance chemically.

Oocnrrenoe and Preparation. — The chief source of nitrogen is the
atmospheric iiir, which consists approximately of one-fifth oxygen
and four-fifths nitrogen. Nitrogen exists also in many forms of
living nuitter, and in the waters and soil in the form of compounds
which are in^H)rtaut for the growth of plants.

It can be prepared in fairly pure condition by removing the oxy-
gen from atmospheric air. This can be accomplished by means of
phosphorus. When moist phosphorus is brought in contact with
the air, the oxygen combines with the phosphorus, forming phos-
phorus pentoxide, PjO^, and the nitrogen remains behind.

The oxygen can be removed from atmospheric air also by certain
metals at an elevated temi)erature. Thus, when metallic copper is
heated to redness in the presence of atmospheric air, the oxygen com-
bines with the copper, forming cupric oxide, CuO, and the nitrogen
remains behind. This methoil is used quite frequently in preparing
fairly pure nitrogen upon the large scale, since the oxygen can be
removed from a considerable volume of air in a comparatively short
time by this methoil. The air is allowed to pass over the copper,
which is heated to reilness in a glass tube, and the nitrogen is
coUectetl as it escai>es from the end of the tube.

Neither of the above methods is capable of yielding very pure
nitrogen, since there is present in atmospheric air small quantities
of many other substances, as we shall see; and none of these are
removed by the phosphorus. They therefore remain and contami-
nate the resulting nitrogen. It is possible to prepare fairly pure
nitrogen from atmospheric air, but this is a difficult and tedious
operation.

To prepare nitrogen of a high degree of purity, certain chemical

200

NITROGEN 201

reactions are made use of. When ammonium nitrite, a compound •
having the composition NH4NOS, is heated, the following reaction
takes place: —

NH4NO, = 2H,0 + N^

This is an excellent method of preparing pure nitrogen.

Another convenient method of obtaining pure nitrogen is by the
action of nitrous acid, a compound having the composition HNO»
upon urea, an organic compound containing nitrogen, and having the
composition CON2H4 ; —

2HN0, + C0N,H4 = C0, + 3 H,0 + 2N^

CSiemical Properties of Hitrogen. — Nitrogen is characterized by
its inertness, not only at ordinary temperatures, but even at elevated
temperatures. If we consider its chemical inactivity alone, we
would be surprised that Rutherford discovered it as early as 1772.
When we remember, however, that it constitutes four-fifths of our
atmosphere, and that the oxygen can be separated from it, we can
understand why it should have been discovered so early.

A few substances, however, combine with nitrogen at elevated
temperatures, forming compounds known as nitrides. These include
magnesium, boron, lithium, and silicon. On account of its chemical
inactivity nitrogen cannot support combustion, except in the very
few cases of substances which combine directly with nitrogen.

It cannot support life, all animals dying in a very short time
when compelled to breathe only nitrogen. Nitrogen is taken into
the lungs with every breath in quantities about four times as great
as oxygen. On account of its chemical inactivity it does no harm
to the organism, simply serving to dilute the oxygen.

Physical Properties of Hitrogen. — Nitrogen is a tasteless, odor-
less, colorless gas. Its critical temperature is — 146°, and its critical
pressure is 35 atmospheres. It can, therefore, be liquefied, but a very
low temperature must be employed. It forms a colorless liquid,
boiling at — 195®. Nitrogen is liquefied by the same general methods
as air. Indeed, liquid air is four-fifths liquid nitrogen. When liquid
air evaporates, the nitrogen boils off first, as we have seen. The
reason for this is now apparent. Nitrogen liquefies about thirteen
degrees lower than oxygen, which is the same as to say that its
boiling-point is thirteen degrees below that of oxygen. When a
mixture of the two is allowed to evaporate, the lower-boiling liquid
passes off more rapidly and leaves the higher-boiling liquid behind.
Of course some of the liquid oxygen evaporates also, but the nitro-

202

FRINCIFLES OF IXORGAXIC CeOIISTRT

gen eraporatiog more lapidly^ finally leaves behind almost pore
liquid oxygen. Liquid nitrogen, when allowed to eTaporate under
yery small pressure, is an excellent refrigerating agent. It boils in
a vacuum at from —225* to —230*,

When liquid nitrogen is cooled to —214°, it solidifies. The
melting-point is^ therefore, above the boiling-point in a vacuum.
When solid nitrogen is warmed in a vacuum, it would, therefore,
j>ass at once into a vapor, without assuming the liquid state.

vCOMPOUNBS OF NITROGEX WITH HYDROGEN

Azmnonia, HH^. — The besl>known compound of nitrogen and
hydrogen is aiamonia. Ammonia occurs in nature in small quantities
in the free condition. It occurs in certain waters, in very small
quantity in the atmospherCj and in certain minerals. In the form of
its salts it occurs in many soils, and on account of their great solu-
bility the salts of ammonia exist largely in solution in water. The
salts of auiioonia are very vaUiaLle in the soil in connection with the
growth of plants, and efforts are contimially being made to cause
their accumulation in soils used for agricultural purposes. Ammonia
is liberateil in considerable quantity by decomposing organic matter.
This is readily detected by the odor of the gas escaping from decom-
posing animal remains or decaying vegetable juatter.

Ammonia can be formed in tlie laboratory by a number of meth-
ods. When nitric oxide is treated with nascent hydrogen, ammonia
is formed : —

6 Hj + 2 NO = 2 H,0 + 2 ^U^.

Ammonia can he formed by the direct union of hydrogen and
nitrogen. When one part of nitrogen is mixed with three parts of
hydrogen and electric sparks passed througli tlie mixture^ a part of
the hydrogen and nitrogen combine, forming ammonia. The com-
bination is far from complete, unless the ammonia is removed as fast
as formed. In the latter case all of both g^es can be made to
combine : —

K,4-3H, = 2Nn^

The volume of the ammonia formed is just half the sum of the vol-
umes of the nitrogen and hydrogen whirh have disap|)eared. If one
V€^!ume of mtrofjf^n^ cojnbmes wUh ihfee voluttms of htfdrogen, there urc
two voi u me » of u m m o n (a formed.

This shows again the simple relations by volume in which gases

NITROGEN 203

combine, and the simple relation between the volumes of the original
gases and the volume of the product.

Ammonia is most conveniently prepared by the action of a base
on an ammonium salt. When ammonium chloride, nitrate, or sul-
phate is boiled with an aqueous solution of a strong base like sodium
hydroxide, ammonia gas is liberated : —

NH4CI + NaOH = NaCl + H^O + NH3 ;

NH4XO3 + NaOH = NaNOs + H^O + NH3 ;

(NH4)j SO4 + 2 NaOH = Na,S04 + 2 H^O -f 2 NHj.

In the laboratory ammonia is prepared most conveniently by
mixing slaked lime with ammonium chloride and warming the
mixture. The reaction is —

2 NH4CI + Ca(OH), = CaCl, + 2 H,0 + 2 NH3.

Ammonia was formerly obtained from decaying organic matter,
and from ammonium salts which occurred in certain arid regions of
the earth. The ammonium chloride which occurred in the neighbor-
hood of the temple of Jupiter Ammon was termed sal ammonictCy
whence the origin of the name ammonia. Ammonia to-day is ob-
tained mainly from the dry distillation of coal in the manufacture
of illuminating gas. The ammonia liquor from the gas-works is
treated with sulphuric acid, when ammonium sulphate is formed.
In this form ammonia can be readily transported, and can be ob-
tained in free condition from the sulphate by treating the latter
with a strong base.

CSiemioal Properties of Ammonia. — Ammonia in the pure, dry
condition is not active chemically. When perfectly dry ammonia gas
is brought in contact with perfectly dry hydrochloric acid gas, there is
not the slightest reaction between the two substances. If there is a mere
trace of moisture present, the two gases react at once, forming the
solid ammonium chloride.

Certain metals like sodium react with ammonia. When ammo-
nia gas is passed over metallic sodium, the following reaction takes
place : —

2 NH3 -f Na = 2 NH^Na -^ Hj.

Ammonia dissolves in water with the greatest ease, forming a
compound which neutralizes acids and which is, therefore, a basic
substance. From the study of a large number of basic substances
we are led unmistakably to the conclusion that all bases contain the

204

PRrNCIPLES OF mORGANIC CHEMISTRY

group (OH), known as liydroxyl; and when bases are dissolved in
water this group splitii off as the anion, and gives the basic charac-
ter to the solution of the substance in question,

WTien ammonia dissolves in water, it must, therefore, combine
with the water, forming the compound KH^OII : —

The corapo^md NH^OH, which, however, has never been isolated, is
then acted on by more water, and dissociated thus r —

Tlie hypothetical group NH^ is called amnwnium^ While this group
lias not been isolated, there is little doubt as to its existence. It
plays the same r6le, as wti shall see, that a metal atom does m the
forraation of compounds.

Composition of Ammonia. — We have seen that one volume of
nitrogen cotuhioes with three volumes of hydrogen, forming ammo-
nia. From Avogadro*s law, there are just as many ultimate parti-
cles or molecules in one volume of hydrogen as in one volume of
nitrogen, therefore three times as many in three volumes of hydro-
gen. From a study of the densities of hydrogen and nitrogen, we
have seen that the molecule of eat^h substance is composed of two
atoms. Therefore, in three molecules of hydrogen we have six
atoms, and in one molecule of nitrogen two atoms. Since one vol-
ume of nitrogen combines with three volumes of hydrogen, then, to
form ammonia, the molecule of ammonia must be NH^, or some mul-
tiple of NHj. Ky a vapor-deusity determination, w^e decide this
part of the question and find that ammonia is I^Ha.

This is the synthetical method of determining the composition of
ammonia. There is also the analytical method*

When ammonia is treated with chlorine, it is decomposed into
hydrochloric acid and nitrogen : —

2NH3-|-3C1, = 6HC1 + Nj,

Chlorine combines with hydrogen volume for volume. It is there-
fore only necessary to know the volume of the chlorine used up, and
the volume of the nitrogen set free when chlorine acts on ammonia,
to know the volume of the hydrogen which was combined with the
Tolume in fjuestion of the nitrogen tjo form ammonia.

A glass tube closed by means of a stop-cock^ and containing above
the stop-cock a reservoir for holding concentrated ammonia, is filled
with pure chlorine. This tube is divided into three equal parts.

NITROGEN 205

The concentrated ammonia is allowed to flow through the stop-cock
drop by drop. When it comes in contact with the chlorine, the
action is so vigorous that there is a flash of fire as each drop of
ammonia enters the tube containing the chlorine. When ammonia
has been admitted to the tube until all the chlorine is used up,
shown by the fact that when more ammonia is added there is no fur-
ther evidence of chemical action, some more ammonia is run in to
combine with the hydrochloric acid which has been formed as the
result of the reaction. When all the hydrochloric acid has com-
bined with the ammonia, forming ammonium chloride, which dissolves
in the aqueous ammonia, and the gas in the tube allowed to come to
normal pressure by admitting water as long as the pressure of the
air will drive it into the tube, the tube will be found to be exactly
one-third full of nitrogen gas. The tube which was full of chlorine
at atmospheric pressure has a volume which is just equal to that of
the hydrogen which was combined with the nitrogen set free. This,
we have seen, is one-third of the volume of the chlorine originally
used, and, therefore, of the hydrogen with which the nitrogen was
combined in the ammonia. Ammonia consists, then, of one volume
of nitrogen combined with three volumes of hydrogen.

The equations expressing the reaction of ammonia on chlorine^
and then on the hydrochloric acid formed, are —

3Cl,-f2NH3 = 6HCl + N,;
6HCl-f6NH3 = 6NH4Cl.

Physical Properties of Ammonia. — Ammonia is a colorless gas
with a very penetrating odor. One litre of ammonia at 0° and 760
millimetres pressure weighs 0.775 grams. The critical temperature
of ammonia is 130®, so that it can be easily liquefied. At 10° it is
converted into a liquid when subjected to a pressure of 6.2 atmos-
pheres. It boils under atmospheric pressure at — 33°.7, and is con-
verted into a solid which melts at — 78°.3. Liquid ammonia is a
very interesting substance. It has been shown to have considerable
dissociating power. Solutions of salts dissolved in liquid ammonia
conduct the electric current very well, and in some cases better than
solution in water at the same concentrations. This does not mean
that salts in liquid ammonia are dissociated to a greater extent than
in water. Dissociation, as we have seen, depends upon the molec-
ular conductivity /n^, at any dilution, v, and also upon the molecular
conductivity at complete dissociation, fi^. fi^ may be larger in
liquid ammonia, and fi^ still larger for any substance in the ammo-
nia, when the dissociation, a, which is the ratio between the two^

206

PRINCIPLES OF raORGANTC CHEMISTRY

would be sinaller in the ammonia than in water. Such being the ease,
fi^ is larger for a given sub9tam*e in liquid ammonia than in water.
The question arises, however, Why is fi^ larger in ammonia than in
water? ^t^ depends upon two quantities^ the number of the ions
present and the velocity with which they move. We have just seen
that the number of ions present in the ammonia is less than in water,
and must conclude, tbereforej that the ions move with greater veloc-
ity in liquid anuuonia than in water. There are methods available
for measuring the relative velocities of ions, but these have not yet
been applied to liquid ammonia.

Liquid ammonia has a ven/ hkfh Sjtecific heat. According to some
authorities, slightly higher even than water. On account of its high
heat of vapori nation it is an excellent refrigerating agent, and is used
extensively for this purpose, especially in connection with the artiji-
cial prtparathn of ke. Ammonia gas is liquefied by pressure, a large
amount of heat imng, of course, set free during the process. This
heat is removed by a cur rent of cold water flowing around the vessel
in which the liquefaction is taking place. The liquid ammonia then
Hows into tulles which closely surround the vessels containing the
water which is to be frozen, and is allowed to vaporize in these tubes.
In vaporizing it must obtain heat from somewhere, and takes it
from surmundhig objects. The water loses its heat, is cooled below
the freezing temperature, and solidifies. The ammonia, having
passed into the form of a vapor, is not lost, but is pumped into the
liquefying chamber, subjected again to pressure and liquefied, the
heat set fi-ee being again removed by the current of cold water-
The process is thus a continuous one, tha same ammonia being used
over and over again.

Maehines for freezing water by means of liquid ammonia, were
early devised by Carre and were known as CVirr^ ice machines. Many
of the modern devices are modifications of these machines of Carre,
utilizing exactly the same principles.

A few years ago most of the ^' artificial ice ^' was made by the
ammonia process. Now considerable ice is obtained by allowing
water to evapoi^ate into a apace under diminished pressure.

Ammonia dissolves in water, as we have already seen. It is one
of the most soluble of all known gases, one volume of water, at 0°,
dissolving about UoO volumes of ammonia. As the temperature
rises, the solubility of the ammonia decreases very rapiiUy. Aqueous
ammonia has a much smaller specific gravity than water, the sj>ecifio
gravity decreasing as the concentration of the solution increaeea.
A few examples will make this clear: —

NITROGEN 207

Pebcxmtage of Ammonia

Specific Okavitv

1 per cent

o.»9a

6 per cent

0.979

10 per cent

0.959

26 per cent

0.911

30 per cent

0.898

36 per cent

0.884

At first it may not be perfectly clear how a concentrated solution
of ammonia in water could be lighter than pure water. When
ammonia dissolves in water there is a large increase in volume, and
this more than compensates for the addition of the ammonia.

Ammonium^ HH4. — When ammonia dissolves in water, we have
seen that it combines with a molecule of the water, in the sense of
the following equation: —

NH3H-H,0=NH,0H.

This compound, ammonium hydroxide, dissociates in the pres-
ence of more water into the ions ammonium and hydroxide : —

NH40H = NH4,OH.

While the group ammonium has never been isolated, it acts as a
unit in compounds which ammonia forms with acids. In its chemi-
cal properties it so closely resembles the alkali metals, that it is
classed with them.

Ammonium Amalgam, HH^Hg. — While the compound NH4has
never been isolated, its amalgam, or compound with mercury, is
readily prepared. When sodium amalgam, a compound of sodium
and mercury having the composition NaHg, is treated with a con-
centrated solution of ammonium chloride, the amalgam swells up,
occupying a relatively large volume. The product has a metallic
lustre, and is probably ammonium amalgam. The reaction probably
takes place in the sense of the following equation : —

NaHg + NH4CI = NaCl + NH,Hg.

It would seem that ammonium amalgam was a hopeful substance
from which to obtain the group ammonium. It is, however, unstable,
breaking down at ordinary temperatures into ammonia, hydrogen,
and mercury.

This is, apparently, the nearest that we have come thus far to
obtaining the group ammonium in the free condition, but it is obvious

208

PRINCIPLES OF INORGANIC CHEMISTRY

that the group is unstable under all tbe conditions to which it has
thus far lieen subjected.

Hydrazine, N^^* — A number of methods have been devised for
preparing hydrazine. Some of these, however^ involve a knowledge
of organic cheniiiitry and cannot be taken up in this plat^e. One
method of preparing hydrazine can, however, bo referred to.

When potassium nitrite, a compound having the composition
KNOt, is treated with sulphurous acid, the two combine forming
the compound KaSOaNjO^. When this is reduced bj nascent hydro-
geoj hydrazine is formed : —

KaSOaXjOa + 4 H, = K,SO, + 2 H,0 + K.H|.

It is also formed bj the reduction of hyponitrous acid N^OaHo by
nascent hydrogen : —

NaO^Ha + 3 H, :^ 2 HaO + l^jH^.

Properties of Hydrazine, — Hydrazine is a liquid boiling at llS***
It forms a crystalline solid which melts at VA. It combines with
water, forming the hydrate NjHi^H^O. Like ammonia, It has basic
properties forming salts with acids.

Triazoic Acid, or Hydrazoic Acid, HIT3- — This remarkable com-
pound w^as discovered a few years ago by the German chemist
Curtius. The compound is remarkable on account of its composition
and properties. It is surprising that we should have a compound
containing three nitrogen atoms and a hydrogen atom, and nothing
else. It is still more surprising that such a compound should have
strongly acid properties.

Hydrazoic acid is prepared most simply by the action of nitrous
oxide, N3O, upon soda amide NaNHj, formed as we have seen by
the direct action of ammonia on metallic sodium: —

N,0 + 2 NaNH, = l^H^ + KaOH + NaK^

When the sodium salt is treated with a strong acid, hydrazoic acid
is formed : —

NaN, + HCl = KaCl -(- H^^

Hydrazoic acid is formed also by the action of nitrous acid on
hydrazine ; —

-R^O, + K,H, =! 2 H,0 + HKj.

Also by the action of an oxydizing agent on a mixture of hydrazine
and hydros yl amine : —

K,H, + KH|0 + 0, = 3 H,0 + HK^

NITROGEN 209

Hydrazoic acid is a colorless liquid, boiling at 37®. It is very-
explosive in concentrated solution, and its fairly dilute, aqueous solu-
tion must be dealt with carefully or explosion will result. It is a
strong acid, its aqueous solutions readily conducting the electric
current. It dissolves many of the metals, forming salts, which re-
semble in general the chlorides, differing from them, however, in
being very explosive. The composition of the salts is as remarkable
as that of the acid itself. The salts with the univalent metals con-
sist of a metal atom united with three nitrogen atoms. When we con-
sider the inertness of nitrogen, it is surprising that such a compound
should be capable of existence. The ammmiium salt has the compo-
sition HNs-NHs = N4H4, and is another compound of hydrogen and
nitrogen. The base hydrazine combines with triazoic acid, forming
the compound N^H^.HN^ = NJI^, a fifth compound of nitrogen and
hydrogen- The five compounds of hydrogen and nitrogen which are
thus far known have, respectively, the compositions : —

NH, N,H

N,H4 N4H4 * *•

CHAPTER XV

NEUTRALIZATION OF ACIDS AND BAS£S8

Ammoniam Hydroxide. — We have seen that when the compound
ammonia is dissolved in water, it combines with the water, forming
ammonium hydroxide : —

NH3 + H2O + NH4OH.

Ammonium hydroxide dissociates as follows ; —

NH4OH = NH4, OH.

Ammonium hydroxide when dissolved in water dissociates into
the cation ammonium and the anion hydroxyl. The hydroxyl ion
and not the ammonium ion gives the characteristic basic property
to the solution. This is shown by the fact that there are many
compounds which, when dissolved in water, dissociate yielding the
ammonium ion, and these solutions have no basic properties. On
the other hand, every compound which yields hydroxyl ions is a
basic substance.

Bases are Hydroxyl Componnds. — That the statement is correct
that bases are hydroxyl compounds, can be seen at once by examin-
ing the composition of a number of basic substances.

Ammonium hydroxide NH4OH

Sodium hydroxide NaOH

Potassium hydroxide KOH

Calcium hydroxide Ca(OH),

Strontium hydroxide Sr(OH),

Barium hydroxide Ba(OH)j

Aluminium hydroxide A1(0H)3

Ferric hydroxide Fe(OH)s

This list of basic substances could be greatly extended. It will
be observed that they all contain a metal combined with one or
more hydroxyl groups. When these substances dissociate, the
hydroxyls split off as anions, and the metal forms the cation. A
few examples will make this clear : —

210

NEUTRALIZATION OF ACIDS AND BASES 211

NaOH = Na, OH.
KOH = K, OH.
Ca(OH),= C^ OH, OH.
Ba(OH),= Ba, OH, OH.
A1(0H)3 = Al^ OH, OH, OH.

Acidity of Bases and Basicity of Acids. — We observe in the
above examples that some bases dissociate yielding one hydroxyl
ion, other bases yield two hydroxyl ions, and others still three
hydroxyl ions. If we take a gram-molecular weight (the molecular
weight of the substance in grams) of a monobasic acid and dissolve
it in water, diluting the solution to a litre, and take a gram-molec-
ular weight of any one of the above bases which yield one hydroxyl
ion and dissolve it so as to form a litre of solution, the litre of the
acid would exactly neutralize the litre of the base. Such a base
which yields on dissociation one hydroxyl ion is known as a monacid
base: — + ^ + ^ + ^

H, CI + Na, OH = H,0 + Na, CI.

If we prepare a solution of a base which dissociates into two
hydroxyl ions, containing a gram-molecular weight in a litre, it will
require just two litres of a solution of an acid such as that referred
to above to neutralize the one litre of the base : —

C^ OH, OH + H, CI + H, CI = 2 H^O -f- CI, CI, Ca.

Such bases are known as diacid bases. To neutralize a gram-
molecular weight of a base which dissociates into three hydroxyl
ions, requires just three litres of the above solution of acid. Such
a base is termed a triacid base : —

Ai^ OH, OH, OH + H, Cl + H, CI + H, Ci =
SHjO + aCcI, ci, CI.

A solution containing a gram-molecular weight of a substance in
a litre is known as a molecular nonnal solution. A solution which
contains in a litre a gram-molecular weight of the substance divided
by its valence is known as an equivalent normal solution. When we
are dealing with a monacid base, the two solutions are identical.
When the base is diacidic, we must divide its molecular weight by
two and dissolve this niunber of grams so as to form a litre of solu-
tion, in order to have an equivalent normal solution. In such a case

212

PRlNCirLES OF INORGANIC CHEMISTRY

the equivalent normal solution is just half as strong as tlie molecular
normal. In the case of a triacld base, the equivalent normal is one-
third of the molecular normal, and so on.

The terms molecular normal and equivalent normal solutions are
used continually^ and their meaning should be clearly understood*

Just as we have mono-, di-, and tri-acid bases, just so we have
mono-, di> and tri-basie acids. An acid which dissociates, yielding
one hydrogen ion, is tnonobasic : ~

HCl^H, CL

If the molecule of the acid yields two hydrogen ions, it is
dibaaic: — + + =

H^04=H, H, SO^

If the molecule of the acid dissociates j yielding three hydrogen
ions, it is inboHic : —

H3A804=H, li, H, AaOij
and 80 on.

Indicators. — In neutralizing acids with bases, we must use some
means to determine when there is no longer any of the acid presentii
or any of the base jJiesent. We make use of certain color changes
which are pro<luced in certain substances by acids, on the one liaud^
and by bases ou the other. If to a solution which contains an acid
a little of the vegetable dye, litmus, is added, tlie litiEUis turns red, j
while in an alkalijie solution it is blue. By adding cautiously &]
little acid to the alkali, or a little alkali to the acid, until the excess i
of the other is just neutralized, we have the neutral tint of the
litmus, which is purple.

Similarly, methyl orange is colorless or slightly yellow in
alkaline solution, and deep red in a solution which is acid. Phenol*
phthalein is red iu the presence of an alkali, and colorless in the
presence of an acid. Cyanine is blue in the presence of a kise and
colorless in the presence of an acid.

We understand pretty thoroughly the action of these indicators,
now that we have the theory of electrolytic dissociation.

Theory of Indicators. — Chemical molecules may \je colored op
colorless, and ions may be colored, giving their color to completely
dissociated solutions. A molecule may have the same color as the
ions into which it dissociates, or it may have a different color. Ai
colorless molecule may dissociate into ions, one or more of which is j
colored, and a colored molecule may dissociate into colorless ions.

Upon these facts is based the use of indicators in quantitative

NEUTRALIZATION OF ACIDS AND BASES 213

analysis. An indicator is a compound which shows a change of
color when the solution passes from the acid to the basic condition,
and vice versa. An indicator is always either a weak acid or a weak
base, which, on dissociation, yields an ion which has a different color
from the molecule itself. Indicators fall then, naturally, into two
classes, — acidic indicators and basic indicators. As an example of
an acidic indicatory we will take first phenolphthale'in. This is a weak
acid, which means that in the presence of water it is very slightly
dissociated, if it is dissociated at all. The molecules of phenol-
phthalein are colorless, as is shown by the fact than an aqueous or
alcoholic solution of this substance is colorless. If a solution of a
strong base is added to phenolphthale'in, the salt of that base is
formed. This salt, like most salts, is readily dissociated in the
presence of water. The salt of phenolphthale'in dissociates into the
cation of the base and the complex organic anion; e.g, the sodium
salt dissociates into the cation sodium and the complex organic
anion ; and it is this latter which gives the characteristic color of
this indicator.

In using this indicator, a small quantity is brought into the pres-
ence of the acid, which is to be titrated against a strong base. The
indicator, in the presence of pure water, is almost completely undis-
sociated. In the presence of the strong acid, which contains many
free hydrogen ions, it would be dissociated even less than in pure
water, as we shall learn. An alkali is added and the strong acid is
all neutralized. The moment an excess of alkali is present, it forms
a salt with the phenolphthalei'n. This salt dissociates at once, and
the colored anion gives its characteristic color to the solution.

PhenolpJUJiale'in cannot be used with weak acids nor weak bases. If
the acid is so weak that its salts, even with strong bases, are hydro-
lyzedy i.e. broken down by water into the free acid and the free base,
the free base would begin to react with the phenolphthale'in long
before enough base had been added to completely neutralize the acid.
The result would be the appearance of a faint color on the addition
of a little alkali, and this color would increase in intensity as more
and more alkali was added. There would, then, be no sharp change
in color when all the acid had been neutralized, and the indicator
would be practically worthless in such cases. Thus, carbonic and
phosphoric acids and the phenols cannot be titrated with phenol-
phthaleln as an indicator. If a weak base is used, such as ammonia,
there will also be a certain amount of hydrolysis of the salt. This
will leave some free base present, which will react with the phenol-
phthale'in and give rise to a gradual change in color. But even

214

FHIXCIPLES OF IXORGAXIC CHEMISTRY

if the ammonium salt of the acid which is being titrated is not
hycirolyzed by water, ammonia cannot be used with phenolphthaleM,
Aiumonia h a weak base, and phenol phthalet'n is a weak acid, and the
salt of tlie two would itself be hydrolyzed by water The indicator
would, therefore, not act sharply when amnion ia was used as a baee.

It is well known that the facts agree very satisfactorily with the
theory. Phenol phthaleln cannot be used as an indicator with either
weak acids or weak hiiBea.

Another example of an acid indicator whoso molecules are nearly
colorless and whose anion is colored, is p-nUrophenoh In alcoholic
solution, in wliieh the substance is almost un dissociated, it is uearly
colorless. Water dissociates it slightly, and consequently the aqueous
solution is slightly colored. If an alkali is added^ the salt of this
weak acid is formed, and this dissociates into the metallic cation,
and into the anion CflH((NO,)0, which is deep yellow in color. The
action of this substance as an indicator will be understood at once
from the above description of the action of phenol phthalein.

Lifmns is an example of an acid indicator whose molecules are
colored, but whose anion has a different color. The molecules of the
weak litmus acid are red. When an alkali is added the salt is
fonned, and this dissociates giving the free litmus anion, which is
d^p blue. Litmus, like phenol phthalein, cannot be used satisfac-
torily with weak bases. These would form salts with the litmus,
which woidd be hydroly:£ed and prevent a sharp color reaction ; or
tifceir salts, with any Imt tbe strongest acids* would undergo some
In^ialysis and prevent a sharp appearance of colon In order that
BtaM sbould be used in titrating weak a^inds, only the strongest
^mm cvi be employed.

tm w&i indicator which can, however, l>e used with weak ba^es
^ i^llni mmst^' This is a considerably stronger acid than the indi-
^gl^wMk wt have already considered. The molecules of the free
MS^^Vi^ *^ imioiis yellow. In the presence of a strong acid
, tbe charactenstic red color j while in the presence
» farmed, and this dissotnates» yielding the yellotr
t-Mi be used with weak basesj provided they
mlds* In these cases there is but slight
■sedf ftiid also but slight liydrolysis of the
mud »nd the weak base, since the indi-

t<l wM indicators it will be seen that
i with strong bases, and a weakly

NEUTRALIZATION OF ACIDS AND BASES 215

Weak bases, on the other hand, must be titrated with strong
acids, and a strongly acid indicator must be used.

Basic indicators are but little used in practice. As an example of
this class we may take cyanine. This is a weak base, and, therefore,
but little dissociated. The molecules are deep blue in color. In the
presence of an acid a salt is formed, which dissociates into the anion
of the acid and the cation of the base. This very complex cation is
colorless ; consequently, the indicator is blue in the presence of a
base, and colorless in the presence of an acid.

The examples considered above suffice to illustrate the different
types of indicators, and to show how satisfactorily their action is
explained in terms of the theory of electrolytic dissociation.

Salts. — When a dilute solution of an acid acts on a dilute solution
of a base, what takes place and all that takes place is the formation
of a molecule of water : —

CI, H + OH,Na = H,0 + Ci,Na.

The sodium ion remains after the process of neutralization in
exactly the same condition as before, and, similarly, the chlorine
remains in the ionic condition. The hydrogen and hydroxyl ions,
however, unite and form a molecule of water. It is a general rule that,
whenever we have hydrogen and hydroxyl ions in the presence of one
another uncombinedy they unite and form vjater. There is an abundance
of direct experimental evidence in favor of this conclusion.

If, however, we evaporate the solution containing the sodium and
chlorine ions, they unite and form a molecule of sodium chloride.
This is a salt. We would define a salt as follows : A salt is a com-
pound formed by the union of an anion of an acid with a cation of a base.
This takes place generally, as already stated, only when the solution
containing these ions is evaporated and at least a part of the water
removed.

The salts are named after the acids from which they are derived.
Salts of hydrochloric acid are called chlorides, those of nitric acid
nitrates, and those of sulphuric acid sulphates. In general, salts of
acids which end in " ic " are termed " ates" ; salts of sulphurous acid
are called sulphites, salts of nitrous acids nitrites, and so on. In
general, salts of acids which end in " ous " end in " ite,-^

So much for the nomenclature of salts in terms of the acids.
Since the cation also enters into the salt, we must be able to di.v
tinguish the salts of one cation from the salts of another cation.
The name of the cation is used before the name of the acid with
whose salt we are dealing. Thus, the chloride of sodium is known as

216 PRINCIPLES OF INORGANIC CHEMISTRY

sodium chloride, the chloride of calcium, calcium chloride, and so on.
When we come to metals which show different valence the case is a
little more complicated.

Take the case of copper. It forms two chlorides — CuCl and
CuClj. The former, in which the copper is monovalent, is known
as cuprous chloride, and the latter cupric chloride. Take the two
chlorides of iron — FeCl, and FeCl,. The former is known as fer-
rous chloride and the latter as ferric chloride. It is a general rule
that the name of the salt in which the metal has the lower valence,
i.^. carries the smaller electrical charge, ends in oiw, while the name
of the salt in which the metal has the higher valence ends in tc.

One further point in connection with the nomenclature of salts
must be mentioned. If we are dealing with a dibasic acid, there are
two possibilities. We may have a salt still containing one of the
hydrogen atoms of the acid, as KHSO4, KHSO^, etc. These are known
as acid potassium sulphate and acid potassium sulphite, while the
salts K2SO4 and KjSOj are known as normal potassium sulphate and
normal potassium sulphite. The acid salts are also frequently
known as primary salts — primary potassium sulphate and pri-
mary potassium sulphite.

When we are dealing with salts of tribasic acids, we have three
possibilities, and in many cases they are all realized. Take phos-
phoric acid, H3PO4; we can have three salts with a univalent
cation : —

KII,P04, KjHP04, and K3PO4.

The first is known as monopotassium phosphate, or primary potas-
sium phosphate ; the second as dipotassium phosphate, or secondary
potassium phosphate; and the third as normal potassium phos-
phate.

There is still a class of salts which we have not considered.
Just as we may have acid salts in which part of the acid hydrogen
remains, so we may have basic salts, in which part of the unneutral-
ized hydroxyls remain. Take bismuth hydroxide; there are three
nitrates having the compositions : —

Bi(0H)oN03, BiOH(N03)5^ and BiCNOaV

These are known as bismuth mononitrate, bismuth dinitrate, and bis-
muth trinitrate or the normal nitrate of bismuth.

Having considered the nomenclature of salts at sufiBcient length,
we shall pass to the study of the energy changes which take place
when an acid is neutralized by a base.

NEUTRALIZATION OF ACIDS AND BASES 217

Heat of Hentralization. — When solutions of acids and bases are
brought together, heat is liberated. Quantitative measurements of
the amounts of heat set free brought out a simple and very impor-
tant relation. This can best be seen from the following results for
strong acids and bases. Gram-molecular weights of different acids
were brought together with a gram-molecular weight of a given base,
both the acid and base being present in very dilute solution. The
amounts of heat set free by a number of acids when neutralized with
the base sodium hydroxide, were : —

Hydrochloric acid and sodiam hydroxide
Hydrobromic acid and sodium hydroxide
Nitric acid and sodium hydroxide
Hydriodic acid and sodium hydroxide
Chloric acid and sodium hydroxide
Bromic acid and sodiam hydroxide .
Iodic acid and sodium hydroxide

Hbat op
Nbutkalizatiox

13,700 cals.

13J00 cals.

13J00 cals.

13,ft00 cals.

13,700 cals.

13,780 cals.

13,810 cals.

The remarkable fact comes out that the heat of neutralization of
these strong acids with a given base, sodium hydroxide, is a
constant.

This suggests a further question very closely correlated to the
above. Suppose we neutralize a given acid with a number of bases,
will the heat liberated be a constant ? and if so, will this bear any
close relation to the above constant where the base was the same
and the acid changed ? This can be answered by the following re-
sults, in which hydrochloric acid was neutralized by a number of
bases : —

Heat or

NsUTBALIZATIOir

Hydrochloric acid and lithium hydroxide .... 13,700 cals.

Hydrochloric acid and potassium hydroxide . . . 13,700 cals.

Hydrochloric acid and barium hydroxide .... 13,800 cals.

Hydrochloric acid and calcium hydroxide .... 13,000 cals.

The heat of neutralization of a given acid with a number of bases is
also a constant, provided the acid and bases are present in very
dilute solution. But what is even more surprising, the constant in
this case has the same value as in the preceding case where the base
was unchanged, and the nature of the acid varied.

These facts when they were first discovered were very perplexing.
Indeed, no satisfactory explanation of them could be furnished, and
it was not until the theory of electrolytic dissociation was proposed
that we could account for them at alL

218

PRINCIPLES OF INORGANIC CHEMISTRY

Explanatioii of Hie CoEstant Heat of B'entralisatioii of Strong Acids
and Strong Bates. — It is one of the crow u in g glories of the theoiy
of electrolytic iiissociation, that it not only explains all of the facts
in eonneetion with the neutralisation of strong acids and b^es in
dilute aqueous solation ; but these fgcts are a necessary consequence
of the theory-
Take, as an example, hydrochloric acid and sodium hydroxideL
In a very dilute, aqueous solution of hydrochloric acid all the mol^
cules are dissociated into hydrogen and chlorine ions thus: —

HCl ^ H H- CL

Similarly^ in dilute aqueous solution the molecules of Bodiura hj-
droxide are completely broken down into ions : —

NaOH ^ Na, OH.

When the dilute aqueous solutions of the base and acid are bronght
together, the following reaction takes place : —

Na, OH + K, Cl ^ Xa, CI + H,0.

The cation of the base, sodium, and the anion of the acid, chlorine,
remain in aoltition as ions after tlie process of neutralization in
exactly the same condition as before neutralization took place. The
anion of the base, hydroxy], and the cation of the acid, hydrogen,
combine and form a molecule of water.

It may l>e urged that the sodium and chlorine ions combine, since
sodium chloride is formed as the result of the neutralization. The
salt is formed if the solution is evaporated j Le. if the solution is con-
centrated. Tiut it can be shown by several se])artite and inde|)endent
methods, that a dilute solution of sodium chloride contains only ions
and no molecules. The sodium and chlorine, then, remain as ions*

The hydrogen and hydroxyl combine and form a molecule of water
This is proved by the fact that water is always formed as the result
of the process of neutralization; and further, it has been shown hj
a half-dozen different methods that hydrogen and hydroxy 1 ions can-
not remain in the presence of one another uncombined to any ap-
preciable extent. This is the same as to say that water is practically
undtssociated.

Since hydroxyl is the anion of every base, and hydro^n the
cation of every acid, the process of neutralization of any strong acid
with any strong base in dilute solution, consists in tlie union of the
hydroxyl ion of the base with the hydrogen ion of the acid, forming
a molecule of water.

NEUTRALIZATION OF ACIDS AND BASES 219

The process of neutralization of any acid by any base is, there-
fore, exactly the same as the process of neutralization of any other
acid by any other base. The total heat that is liberated when a gram-
equivalent of a completely dissociated acid acts on a gram-equivalent
of a completely dissociated base, is the heat set free by the union of a
gram-equivalent of hydroxyl ions with a gram-equivalent of hydrogen
ions. Thus : —

H aq -f OH aq = 13,700 cals.

Since all processes of neutralization of completely dissociated acids and
bases are the same, the heat of neutralization of all sucli acids and bases
must be a constant, and must be the heat of combituUion of a gi'awr
equivalent of hydroxyl and hydrogen ions.

Hentralization of Weak Acids and Bases. — If either the acid or
base is what we term weak, the heat of neutralization is not 13,700
calories, but differs from this value. Thus, take the following ex-
amples : —

Formic acid and sodiam hydroxide
Acetic acid and sodium hydroxide
Dichloracetic acid and sodium hydroxide
Valeric acid and sodium hydroxide .
Phosphoric acid and sodium hydroxide

H»AT or

Neutkalizatioh

13,400 cala.
13,300 cals.
14,830 cals.
14,000 cals.
14,830 cals.

In these cases the acids are weak and the base is strong ; neverthe-
less, there are considerable differences between the heats of neutral-
ization and the constant 13,700 calories.

Similar results were obtained when weak bases were neutralized
with a strong acid. If, however, both acid and base are weak, the
heat of neutralization differs still more from the constant 13,700
calories. A few examples of this condition are given below : —

Heat or
Xevtralizatioit

Formic acid and ammonium hydroxide .... 11,900 cals.
Acetic acid and ammonium hydroxide .... 11,900 cals.
Valeric acid and ammonium hydroxide .... 12,700 cals.

When the weak base ammonia is neutralized by the weak organic
acids, the heat of neutralization differs very widely from the con-
stant 13,700.

Explanation of the Results with Weak Acids and Bases. — If the
acid or base is weak, we shall learn that it is only little dissociated
by water, even in dilute solutions. When only a part of the acid or

220 rUINX'IPLES OF INORGANIC CHEMISTRY

base is dissf)ciate(l, the process of neutralization could proceed only
until ail the disscKuated substance had reacted; were it not for the
fiu^t that as soon as the ions already present begin to react, more
ions wouhl In* formed from the undissociated molecules, or, in a word,
the ])r(K'ess of dissociation would continue as the reaction continued
until all the molecules had dissoc^iated.

Wht'u molecules dissociate into ions, heat is either evolved or
consumed. The thermal change which accompanies the dissociation
of the umlissociated molecules, either increases or diminishes the
amount of heat set free due to neutralization alone. If the heat of
dissociation is positive, it adds itself to the heat of neutralization ;
if negative, it diminishes the heat of neutralization. Thus, the heat
which is liberated when a weak acid acts on a weak base, may be
either greater or less than the constant 13,700 calories — greater, when
the heat of dissociation is positive, less, when it is negative. It could
be eipial to the constant only when the heat of dissociation is zero.

The facts, then, agree with the tlieory, not only when the acid
and base are completely dissociated, but when the dissociation is not
complete. We could predict from the theory of electrolytic disso-
ciation that the h(»ats of neutralization of weak acids and bases
wouhl not be a constant, with the same certainty that we could pre-
dict the constant value of the heats of neutralization of completely
dissiviated aiMds and bases. The api>arent exceptions presented by
the weak acids and leases furnish as strong confirmation of the theory
as the cases which c*mform to rule.

Explanation of the Law of the Thermoneatrality of Solntiona
ef Salts. — The theory of electrolytic dissot^iation furnishes us with
the tirst rational explanation of the law of the ihermoneuirality of
salt solutions. This law. wliich was disrovennl by Hess, states that
vhen dilute solutions of salts are mixed, there is little or no change
in the heat tone. This is a necessary consequence of our theory.
Take two salts, sixiinm chloride and jK»tassium bromide. In dilute
s^ivs>v.s solutions these exist entirely as ions : —

XaCl = Nkci,
KUr = K, Br.

T"-:!!: ri^ solutions of these salts are mixe«l. all of the parts
"■•^t.'* - «-'~'>:'T. as ions. There is no chriiiic;ii artiuu whatso
^^^ i-.^ *rns.^.r:v^T!: T»-maining in the same cuiMiiiiun iifri-r mixing
^ >^T^ Tft*^ ?v then. absolut4?ly no rt-asi-u to exi-ott any

COMPOUNDS OF NITROGEN WITH OXYGEN, ETC. 221

We can now begin to see the importance and wide-reaching sig-
nificance of the theory of electrolytic dissociation. This theory fur-
nishes us with the explanation of the constant heat of neutralization
of acids and bases, and of the law of the thermoneutrality of salts ;
and this is but the beginning. We shall see as our subject develops,
that it has thrown an entirely new light on a great number of chemi-
cal problems which, without its aid, were simply empirically estab-
lished facts, whose meaning was entirely shrouded in darkness. We
shall see that this theory is fundamental, if we hope to raise chem-
istry from empiricism to the rank of an exact science.

COMPOUNDS OF NITROGEN WLTH OXYGEN AND HITDROGEN

Ammoniam Hydroxide, HH4OH. — We have already seen that
ammonium hydroxide is a basic substance. This is a compound
of nitrogen with oxygen and hydrogen, and must be considered
here. The most characteristic property of this substance is its basic
nature. Being a base, it readily neutralizes acids, forming salts.
Ammonium hydroxide unites readily with hydrochloric acid, forming
a well-characterized, beautifully white salt, ammonium chloride : —

NH4, OH + H, ci=xH,0 + NH4, CI.

As the water is removed, the ammonium and chlorine ions com-
bine, forming ammonium chloride : —

NH4, Ci=NH4Cl.

Ammonium chloride or sal ammoniac, it will be remembered, is
of special interest in connection with the determination of vapor-
densities. It was one of those substances which gave abnormally
low vapor-densities, and was for a long time regarded as an excep-
tion to the law of Avogadro. It will be recalled how it was proved
experimentally that when ammonium chloride is heated it breaks
down in the form of vapor into ammonia and hydrochloric acid, and
was shown to present no real exception to the law of Avogadro. Am-
monium hydroxide reacts readily with nitric acid, forming ammo-
nium nitrate, and with sulphuric acid, forming ammonium sulphate.
In the last case there are two possibilities. If there is sufficient
ammonium hydroxide present, the normal sulphate is formed : —

2 NH4OH + H3^04 = (NH4)j^04 + 2 H,0.

222

FKINCJPLK8 OF INOEGANIC CHEMISTBT

If there la ouly one equivalent of axnmouiuin hrdroxide present to
one tiquivalent ol bulphuric acid, the acid sulphate is formed: —

Nif/>Jf 4- li^), = NH,HS(>4+ HA

While aiJUJioniiiiJi hydroxide \i^» basic projierties, it is not a strong
l>ase. Indeed, in eoniparii>on with such substances as sodium hydrox-
ide and pota8()iuni hydroxide, it is a very weak base. The strength
of a ihinf, like the stren}(th of an acid, is measured by its conduc-
tivity. Htrength is proj)ortional to dissociation, and dissociation is

the ratio between tlie molecular conductivity of any dilution, fi^ and

It
the molecular conductivity at infinite dilution, /i^. a = —^*

The valueH of fi^ for several dilutions of ammonium hydroxide
are given lielow : —

2

10

KM)

KKK)

10,1)00

6«)JKMI

Ml*

Mi^CI^-^)

a

1.2

0.5 per cent

3.1

1.4 per cent

0.2

4.4 per cent

26.0

10.2 per cent

fll.l

27.8 per cent

70.0

31.0 per cent

(220)?

To determine the value of fir for a weakly dissociated substance
like amuiouium hydiH>xiile, we cannot prooetnl as already described,
i.<*. iiu'iea.'ie the dilution of the solution until on further increase the
lUoUvuUiv eonductivity di>es not chanj^^. The reason is that the
dilution at which ci»niplete disswiation is reucheil is so great that
%ho coaductivitv luethinl cannot l)e applieil to it. An indirect method
uiJ.si bo cmploNi*il in such cases.

MiMiuwMat of tk# Uttociatiott of a Weak Base, like Ammftwiiim
Bjrtnixiito the ditticulty eui\)untered is iu the determination of
v^V v^iiic v>t .4. . While ammonium hydroxide itself is only slightly
;.^s,^\<,^;i\;, 'ilio siilts of this base are stron>;ly diss»xMatevl, and. in-
.«vvs vi%i|-a';t''»> diss*.>ciated at dilutions to which the conductivity
H^v4v^, ^^ Iv iviklily applied. It is, then, a simple matter to
v\.w^^. u isv \»ilue o( Iks. for an ammonium salt. The question

* ^iv .; ^ u^ : x^ W hut i^nnection exists between the value of ilo f*^r

.^ . «. .<...;..* -x*it and *** for the free btise aiumt)ni;i.*

•%■ . ».vk^^ .X :o Ih: touud in the Luao of Kohlmnach, which say^

COMPOUNDS OF NITROGEN WITH OXYGEN, ETC. 223

that the value of fLti for any compound is the sum of two constants — tlte
one depending on the anion and the other on the cation. The value
of /ix> for a salt like ammonium chloride is, then, the sum of two
constants, a and ft, a depending for its numerical value upon the

cation NH4, and b for its numerical value upon the anion CI. We
determine by the conductivity method the value of /iflo = a + 6.
We know b (70.2) from prevous determinations and obtain a thus : —

a = fix) — ft.

Ammonium hydroxide has exactly the same cation as ammonium

chloride, NH4, and, therefore, the value of a is the same for both
compounds. Ammonium hydroxide is made up of the cation ammo-
nium, whose conductivity constant is a, and the anion (OH), whose
conductivity constant we will call c. We know the numerical value
of c from previous determinations, and determine the value of fioo for
ammonium hydroxide by adding a and c : —

fioo (for ammonium hydroxide) = a + c

If we are dealing with a weak acid, we use in a similar manner
the salt of that acid with a strong base. The value of /i» for the
salt is determined. From this the constant for the cation of the salt
is subtracted, and to the remainder the constant for hydrogen is
added.

Hydroxylamine, B'H3(0H). — The other compound of nitrogen with
oxygen and hydrogen which has basic properties, is hydroxylamine.
Although this compound was discovered in 1865, it was prepared in
the pure condition for the first time in 1891 by Lobry de Bruyn.

Hydroxylamine is prepared by the direct action of nascent hydro-
gen on nitric oxide : ' —

3H, + 2NO = 2NH,OH.

It is also prepared by the reduction of nitric acid by nascent
hydrogen : —

HNOa + 3 H, = 2 H,0 + NH,OH.

It consists of white needles, which, when exposed to moist air,
take up water readily. It is, therefore, a hygroscopic substance.
It melts at 33** and boils under 60 millimetres pressure at 70®.

Hydroxylamine dissolved in water has basic properties. This is
the same as to say that it is dissociated by the water, yielding
hydroxylions. Its conductivity shows that hydroxylamine is, how

224

PRINCIPLES OF LSORGAi^IC CHEMISTRY

ever, only a wciak baaa With acids it forms salts by simple addition,
like ammonium hydroxide;^

NHsOH + HCl ^ KH3OHCI,
2NH3OH H- H3SO4 = (KH,OH)2S04.

Hydroxy lamine is a strong reducing agent. Mercuric cliloride is
reduced to uiercuroua chloride, and an alkaline solution of a copper
salt is reduced to cuprous oxide »

Ilydroxylamino is reduced by nascent hydrogenj forming am-
monia and water : —

KH,OH + Ha = H,0 + mi,,

Compoands of Nitrogen with Oxygen. — Nitrogen forms the follow-
ing compounds with oxygen: Nitrous oxide, NjO; nitric oxide, NO;
nitrogen aesquiuxide or trioxidejNaOaj nitrogen dioxide or tetroxide^
depending upon whether it has the composition KO^ op K^O^j and
nitrogen pentoxidej N^Oj.

We shall now study these compounds in some detaiL
Kitroas Oxide> N^O. — Kitroua oxide is formed when ammonium
nitrate is heated to 2oO^, The following equation expresses the
reaction which takes place : —

NH,N03 = 2H.O + KA

The c^ygen and hydrogen combine and form water, and the nitro-
gen and oxygen form the compound N^Oj which escaf>es.

Nitrous oxide is a remarkable substance, in that it suj^povts com-
bustion almost as w*ell as pure oxygen, rhosphorus and carbon bum
readily in nitrons oxide. Certain substances, however, burn in
oxygen and burn less readily or do not burn in nitrous oxide. The
products of combustion in nitrous oxide are the same as in pure
oxygen^ showing that the compoimd, Nfi^ is readily broken down,
yielding free oxygen.

When nitrous oxide is inhaled into the lungs, it produces a
remarkable physiological effect, generally throwing the subject
into a hysterical condition. It is, the re fore j know^n m laurjinng gan.
When consumed in larger quantity it produces anesthesia, and
13 consequently used in minor surgical operations.

Nitrous oxide is a colorless gas^ with a sweetish taste, and dis-
solves readily in cold water. It should, therefore, be collected in
cylinders over hot water. Its critical temperature is 39°, and critical
pressure 3G atmospheres. It liquefies under atmospheric pressure at
— 87% and solidities at —115% When boiled under diminished

COMPOUNDS OF NITROGEN WITH OXYGEN, ETC. 225

pressure, temperatures as low as — 135° to — 140® can be produced.
In the liquid form it is an excellent refrigerating agent.

In reference to the energy changes which take place during chemi-
cal reactions, we meet hei-e for the first time with a new condition.
When most chemical reactions take place, heat is evolved, — most
reactions are exothermic. In this case the opposite is true. When
nitrogen and oxygen combine to form nitrous oxide, heat is absorbed.
Such chemical reactions, which take place with absorption of heat,
are kncJwn as endothermic reactions,

Vitric Oxide, NO. — ^Nitric oxide can be formed by the reduction
of the higher oxides of nitrogen, or by the reduction of nitrous or
nitric acid. It is prepared most conveniently by the action of nitric
acid on metallic copper. The equation expressing the reaction is : —

8 HNO3 + 3 Cu = 4 H,0 + 3 Cu(N03)2 4- 2 NO.

As quickly as the colorless gas, nitric oxide, is brought in con-
tact with free oxygen, the two combine at ordinary temperature : —

2NO-f-02 = 2NO,.

The gas NOj, as we shall learn, has a yellowish brown color, and
as quickly as nitric oxide is brought in contact with the air, the
above reaction takes place, giving the characteristic colored fumes.
These are the fumes which always appear when nitric acid acts on
metallic copper.

Nitric oxide not only has a remarkable power to combine with
oxygen, forming a higher oxide of nitrogen, but also the power of
giving up some of the oxygen which it already possesses, — of being
an oxidizing agent. Certain substances, like phosphorus and magne-
sium, if once ignited, will continue to bum in nitric oxide, forming
oxides with the oxygen obtained from nitric oxide. All of the
oxygen is removed by metallic potassium or metallic sodium, free
nitrogen remaining. Nitric oxide produces a dark, violet color when
brought in contact with a warm solution of a ferrous salt. This
reaction is used to detect nitric oxide.

Nitric oxide is a colorless gas, whose critical temperature is
— 93®.5, and whose critical pressure is 71.2 atmospheres. Its boil-
ing-point is — 153®.6. The critical temperature being so low, it is
much more difficult to liquefy than nitrous oxide.

Nitrogen Sesqnioxide or Nitrogen Trioxide, N^O,. — Nitrogen
sesquioxide is obtained by the action of arsenic trioxide, As^Osi upon

Q

226

PRINCIPLES OF INORGAXIC CHEMISTRr

nitric acid* Also by the action of nitric oxide upon nitrogen perox-
ide, the temperature not being above — 21° : —

Further, it is obtained by the action of a strong acidj like
sulphuriCj upon nitrites; — ,

2 NaNOj + H8SO4 = Ka^SO* + H^O + NjOa,

It is, therefore, sometimes called nitrous anhydride^ since it is
nitrous acid mituis water; and, by the addition of an alkali, it forms
nitrites, Nitmgen sesquioxide is stable only at low temperatures-
Above — 20^ or —15°, it begins to decompose into nitric oxide and
nitrogen dioxide :^ —

At very low temperatures it passes over into a deep-bine liquid.

Kitrogen Bioxide or Kitrogen Peroxide, HO^ — ^^^itrogen dioxide
is conveniently formed by heating dry lead nitrate : —

2 Pb(KO,)s = 2 PbO + 0, + 4 KOt-

Also by a method which we have recently studied, — the action of
oxygen on nitric oxide ; —

2KO-hO,^2NOa.

It can also be prepared by the direct union of nitrogen with
oxygen. When an electric spark is passed through a mixture of
these two gases, containing one volume of nitrogen and two vol-
umes of oxygen^ they combine to a certain extent^ forming nitrogen
dioxide : —

Nt + 20,==.2KOj-

Nitrogen peroxide is a strong oxidizing agent It is also very
poisonous. When treated with cold water it decomposes into nitrous
and nitrio acids t —

2 NO, + HjO = HKO, -h HNO^
When treated with hot water, it yields nitric acid and nitrio oxide ; —

3 KOt + H,0 = NO + 2 HNOa.

Nitrogen dioxide liquefies at 22*, forming a reddish-brown liquid.
As the tenii>erature is lowered the color gradually disappears, until
at — 20° it passes over into a colorless solid which melts at — 12*.

The physical properties of the vapor of nitrogen peroxide are

COMPOUNDS OF NITROGEN WITH OXYGEN, ETC. 227

unusually interesting. The rapor-density varies with both the tem-
perature and pressure. The lower the temperature and higher the
pressure, the larger the specific gravity of the vapor, and, conse-
quently, the larger the molecular weight calculated from the specific
gravity. At comparatively low temperatures (20**), and especially if
the pressure is high, the molecular weight is 92, which corresponds
to the formula Nj04. Above 100®, if the pressure is only a few centi-
metres, and about 140** at atmospheric pressure, the molecular weight
calculated from the vapor-density is 46, corresponding to the formula
NO^

These changes in the vapor-density are accompanied by corre-
sponding changes in the color of the gas. At low temperatures the
vapor is only a little colored, being somewhat yellowish brown. As
the temperature is raised the color becomes darker and darker, until
finally, at an elevated temperature, it becomes quite dark. When
the vapor is cooled again, the original color is restored. We obvi-
ously have to deal here with two substances of the composition N0|,
the one, N2O4, being a polymer of the other. At high temperatures
and low pressures only the former exists, having a dark, reddish
color. At low temperatures and high pressures we have the com-
pound N2O4. As we ordinai-ily have to deal with the gas, it is a
mixture of these two isomeiic substances.

It is obvious that we have to do here with a condition of equi-
librium between the two substances NOj and Nj04 — a condition
which can be changed by varying either the temperature or press-
ure, and still more by varying both. The higher the temperature*
and lower the pressure, the more NO2 is present ; the lower the tem-
perature and higher the pressure, the more N2O4 is present If we
keep the temperature constant, the pressures of the NOs and ISfii in
the mixture conform to the equation

p2

— = constant,

P

where pi is the pressure of the NO, and p is the pressure of Vfi^
Pure N2O4 melts at — 9®, forming a colorless liquid which freezes

to a colorless solid. When warmed it breaks down into NOj^ with

its characteristic reddish-brown color.

Vitrogen Pentoxide, HA- — Nitrogen pentoxide is fonned bj the

action of strong dehydrating agents, like phosphimis pentoxide,

upon nitric acid : —

2 HNO3 + P2O, = (PfO^.HjO) + TSfi^

228 PRINCIPLES OF INORGANIC CHEMISTRY

Nitrogen pentoxide readily combines with wate^, forming nitric
acid: —

N A + H,0 = 2 HNOs.

It is, therefore, the anhydride of nitric acid. It is a powerful oxi-
dizing agent, as would be expected from the large amount of oxygen
which it contains, decomposing into nitrogen dioxide and oxygen : —

2NA = 4NO, + 0^

Nitrogen pentoxide forms colorless crystals, melting at 30**. The
liquid boils with partial decomposition at 50®.

Acid Compounds of Nitrogen with Oxygen and Hydrogen. — Nitro-
gen forms three acids with oxygen and hydrogen ; hyponitrous acid,
HNO or HjNA; nitrous acid, HNO,; and nitric acid, HNOj. We
shall study these somewhat in detail.

Hyponitrous Acid, HNO or K.JS[jOi. — The salts of hyponitrous
acid are formed by the careful reduction of nitrates or nitrites by a
mild reducing agent, such as sodium amalgam. By the reduction of
sodium nitrate, sodiimi hyponitrite is formed : —

2 NaNOs + 4 HjO + 8 NaHg = 8 NaOH 4- 8 Hg -f- Na^N^O^

The difficultly soluble silver salt is readily prepared from the
sodium salt, and the free acid obtained from the silver salt by means
of hydrochloric acid dissolved in ether, so as to exclude water : —

AgjNsOs 4- 2 HCl = 2 AgCl -f- H,N,Oj.

It Is also prepared by the action of nitrous acid on hydroxylamine : —

NHjOH + HNOa = H,0 -f- H^N^Oa.

Mthouj^h hyponitrous acid is a weak acid, shown by the small
v^M^sluotivity of its aqueous solution, it forms both normal and acid
!ii^U* Tiu* fornu»r have the composition MgNoOj, the latter, MHNjO^

U\|sM\itnniM acid in aqueous solution decomposes readily into
uj<'k\»M4 sAulo «nd water: —

lT,NA = H20-f-NA

\ii4,Mn .»\ulo »^» thon^foro, the anhydride of hyponitrous acid.

ll\ is'u»iu^u^ twid forms white crystals, which, when free from
\\.»ui. 4iw' \*^i\ oxpUvnivo. The molecular weight of hyponitrous
I it. . . ii u tiuitioU l*v tho frcozing-point method, corresponds to the

Vn . ..M4ii 110 .^i uJ Ua.H Ihhmi described, having the same composi-
Ufii ml ukuio luoUvulur weight as hyponitrous acid, and, therefore,

COMPOUNDS OF NITROGEN WITH OXYGEN, ETC. 229

metameric with it. The difference in the properties of these two
substances is supposed to be due to the different arrangement of the
atoms in space in the two substances. This kind of isomerism is
known as stereoisomerism, which apparently plays a more important
r51e in organic chemistry than in inorganic.

Vitroiu Acid, HNO,. — The salts of nitrous acid, the nitrites, are
obtained by removing oxygen from the nitrates. When a nitrate is
carefully heated, it loses oxygen and passes over into a nitrite.

2NaN03 = 0,4-2NaN02.

If a mild reducing agent, such as metallic lead, is fused with a
nitrate, the reduction to nitrite takes place far more easily and
completely : —

KNO3 4- Pb = PbO -f KNO^

When a nitrite is treated with a strong acid, such as stdphuric,
nitrous acid is set free : —

2 KNO2 + H,S04 = K2SO4 -f- 2 HNOj.

Nitrous acid can exist only in solution. When an attempt is
made to remove the water, the nitrous acid loses water and passes
into the anhydride NjO, : —

2HNO,= H,0-f-Ns08.

Nitrous acid is an excellent reducing agent, since it readily com-
bines with oxygen, forming nitric acid. When brought in contact
with a substance rich in oxygen, like potassium permanganate, it
takes oxygen away from the compound, converting it into colorless
substances. The destruction of the beautiful purple color takes
place very rapidly.

Nitrous acid can also act as an oxidizing agent, giving up some of
the oxygen which it already possesses. Thus, it oxidizes hydriodio
acid to iodine : —

HI + HNO, = H,0 + 1 4- NO.

Vitric Acid, HNOj. — This is not only the most important acid of
nitrogen, but one of the strongest and most important of all known
acids. It was early prepared from nitre, which is potassium nitrate,
whence its name. It can be formed by passing electric sparks
through a mixture of oxygen and nitrogen, as Cavendish showed.
A far better method, however, of preparing nitric acid is by the
action of sulphuric acid on some nitrate.

H^04 -f- 2 KNO3 = K2SO4 + 2 HNO»

230

PRINCIPLES OF INORGANIC CHEMISTRY

A similar reaction takes place when sodium nitrate is treated with
sulphuric acid. These methods are used almost exclusively for the
preparation of commercial nitric acid^ on account of the ahiin dance
of these nitrates which occur in nature* Potassium nitrate^ knowu
as saltpetre, la formed where organic matter is decomposing in the
presence of potassium salts. It occurs in the form of a solid only in
arid regions, since, on account of its great solubility, it would pass into
solution if it came in contact with any appreciable amount of water.
Sodium nitrate occurs in abundance in the arid regions of Chili, and
is known as OldU saiipetre. When sodium niti'ate is treated with
sulphuric acid, the first reaction which takes place is ; —

NaNO, + H,SO, = KaHSO* + HKO^

If the temperature is raised sufSciently, the acid sulphate acts oo
more of the nitrate, decomposing it in the seEsa of the following
equation : —

NaHSO^ + NaNOji = NajSO, + HXOa.

In order that this second reaction may take place, such a high tem-
perature must be used that much of the nitric acid is decomposed.
Only the first reaction is, therefore, allowed to take place in the
preparation of nitric acid,

Chemo&l Froperties of Nitric Acid. ^ The most characteristic
chemical property of nitric acid is its strong oxidising power-
When brought in contact with substances which can take up oxygen,
nitric acid readily gives up its oxygen and passes over into lower
oxides of nitrogen.

When a metal is treated with nitric acid, the hydrogen ion of the
acid gives up its charge to the metal, converting the latter into an
ion, while the hydrogen becomes an atom. Thus far nitric acid acta
just like the other acids which we have studied. The hydrogen, in
the case of nitric acid, however, does not escape, but acts on more
nitric acid, reducing it to lower oxides of nitrogen, or to nitrogen
itself, or even to ammonia, depending npon conditions. When nitric
acid 18 added to metallic silver it loses one molecule of oxygen^ x>ass-
ing over into nitrous acid. When treated with metallic copper, nitric
acid is reduced to nitric oxide, XO, as will be remembered ; while
nitric acid upon zinc is still further reduced, yielding hyponitrous
acid, NOH. In the presence of zinc and sulphuric acid nitric acid
is reduced to ammonia. The powerful oxidizing action of nitric acid
manifests itself towards substances which cannot take the poaitiTB
electric charge from the hydrogen and become an ion. Thus, phos-

COMPOUNDS OF NITROGEN WITH OXYGEN, ETC. 231

phoTUS is oxidized by strong nitric acid to phosphorus pentoxide
or phosphoric acid, and carbon to carbon dioxide. Metallic tin
which does not form a nitrate is oxidized to stannic acid, Sn(0H)4.
The salts of nitric acid — the nitrates — are, without exception, very
soluble in water. They are excellent oxidizing agents. Nitric acid
is still used extensively in the preparation of sulphuric acid. It is
also used in the manufacture of certain dyestuffs, and of such explo-
sives as nitroglycerine, nitrocellulose, etc.

Physical Properties of Vitric Acid. — Nitric acid is a liquid,
boiling with partial decomposition at 86**. The liquid solidifies
at-47^

Nitric acid and water are miscible in all proportions. The acid
having a specific gravity of 1.1 contains 17.1 per cent of nitric acid;
that having a specific gravity of 1.2 contains 32.4 per cent of acid ;
that having a specific gravity of 1.3 contains 47.5 per cent of acid; that
having a specific gravity of 1.4 contains 65.3 per cent of acid, while
the pure acid has a specific gravity of 1.53.

All mixtures of nitric acid and water boil higher than pure nitric
acid. The relations here are similar to those observed with hydro-
chloric acid. When any mixture of nitric acid and water is boiled,
it tends towards the composition of 6S per cent of the acid. This
mixture has a constant boiling-point, which is 120**.5. If the solu-
tion of nitric acid in water is more concentrated than 68 per cent,
acid will distil over until this concentration is reached. If it is
less concentrated than 68 per cent, water will distil over until this
concentration of acid remains behind. This composition corresponds
approximately to the acid HNOS.2H2O =K(0H)5. That this is a
mixture of nitric acid and water and not a definite chemical com-
pound, is proved by the fact that when a different pressure is used
the composition of the mixture changes.

I>etection of Vitrio Acid. — Nitric acid is readily detected by the
dark-purple color produced when it is mixed with a concentrated
solution of ferrous sulphate, both solutions being warm. The test
for nitric acid is made as follows : The nitric acid or the nitrate is
treated with a little concentrated sulphuric acid, and warmed until the
containing vessel feels quite warm to the hand. Another test-tube
is filled about one-third full of crystals of ferrous sulphate, and dis-
solved in just as little water as possible, the solution being heated
until it feels warm to the hand, but not heated to boiling. The
solution containing the nitric acid is now added drop by drop to the
solution of ferrous sulphate, when the dark color will make its ap-
pearance in the form of a ring where the two liquids come in contact.

232 PRINCIPLES OF INORGANIC CHEMISTRY

Dissociation of Vitric Acid and titrates. — Nitric acid dissociates
in the sense of the following equation : —

HNO,^H, NOa.

It is therefore a monobasic acid, and can yield only one series of
salts. These are of the general type MNOa, and dissociate thus : —

MN08=M,Nb8.

The conductivity of nitric acid shows that it is one of the very
strongest acids known.

¥

Mr (18°)

a

1

20

100

Moo = 600

299.1

832.8

842.1

= 842.7

87.3 per cent

97. 1 per cent

99.8 per cent

100.0 per cent

Fuming Vitric Acid. — Fuming nitric acid is formed in the
preparation of nitric acid from sodium nitrate and sulphuric acid,
if the temperature is sufficiently high to cause the acid sodium sul-
phate to react with more sodium nitrate. It is apparently a solution
of nitrogen dioxide in nitric acid.

It is a much more energetic oxidizing agent than ordinary con-
centrated nitric acid. When warmed in its fumes many organic
substances will take fire and burn. When a piece of iron has been
dipped in fuming nitric acid for a moment, and is then removed and
dipped in ordinary concentrated nitric acid, the latter does not act
upon the iron. Iron in this condition is known as in the passive
state. It was supposed for a long time that the iron became covered
with a layer of oxide, which protected it from further action. It is
now known that this is not the explanation of the phenomenon, it
having been recently shown by the German, Hittorf, that the passive
state is purely an electrical phenomenon.

Aqua Eegia. — Certain metals, like gold and platinum, do not dis-
solve in nitric acid, but when treated with a mixture of nitric and
hydrochloric acids they dissolve readily. The mixture which is
most efficient consists of one part of nitric acid and three parts of
hydrochloric acid. This is known as aqua regia^ The nitric acid in
the mixture oxidizes the hydrochloric acid and liberates chlorine.
There is probably also formed one or more compounds containing
nitrogen, oxygen, and chlorine. These probably have the composi-

COMPOUNDS OF NITROGEN WITH OXYGEN, ETa 233

tions NOCl and N0,C1, and are known, respectively, as nitrotyl and
nitryl chlorides. The action of these various substances is to con-
vert the metals into chlorides, even platinum being transformed into
platinic chloride by aqua regia. The name was derived from the
fact that this mixture can dissolve gold.

COMPOUNDS OF NITROGEN WITH THE HALOGFA^S

Compounds of Vitrogen with Chlorine and Bromine. — When chlo-
rine acts upon ammonium chloride, the trichloride of nitrogen, NClg,
is formed, probably in the sense of the following equation : —

3 01, -h NH4CI =4 HCl + NCl^

Nitrogen tricJdoride is a yellow liquid, which explodes very
violently and often with the slightest provocation. The reason
for its instability b doubtless closely connected with the endother-
mic nature of the reaction which produces it. When one nitrogen
atom combines with three chlorine atoms, about 42 calories of heat
are absorbed. This is set free again when the decomposition of the
compound takes place, heating the gases which are formed, and
causing them to exert a great pressure. Nitrogen also combines
with bromine.

Compounds of Vitrogen with Iodine. — Iodine combines with
nitrogen, forming apparently several compounds known as nitrogen
iodides. When ammonia is treated with iodine at low temperatures,
the compound N^,!, is formed; while at ordinary temperatures
we have N3H5I3 produced. The compound INj is formed when a
solution of iodine in ether is allowed to act on the silver salt of
triazoic acid, AgN,. All of these compounds are characterized by
their explosive nature.

COMPOUNDS OF NITROGEN WITH OXYGEN, HYDROGEN, AND

SULPHUR

Vitrosyl-fnlphnric Acid, 802(0H)H0,.— There is one compound
of nitrogen with oxygen, hydrogen, and sulphur — nitrosyl-sulphuric
acid — which must be considered on account of its importance in the
manufacture of sulphuric acid. It will be remembered that this
compound is formed by the action of nitrous acid on sulphur
dioxide in the presence of the oxygen of the air: —

2 SO, + 2 HNO, + 0, = 2 SO,(OH)NO,.

234 PRINCIPLES OF INORGANIC CHEMISTRY

Nitrosyl sulphuric acid crystallizes very frequently in the lead
chambers, and is then known as chamber crystals. When these come
in contact with water-vapor, they decompose in the sense of the
following equation : —

2 SOs(OH)NO, 4- H^ = 2 SO^OH), + NjO,,

the proilucts being sulphuric acid and nitrogen sesquioxide.
This compound is also known as nitrosulphonk aeicL

CHAPTER XVI

THE ATMOSPHERIC AIR AND CERTAIN RARE ELEMENTS
OCCURRING IN IT

THE ATMOSPHERIC AIR

It was stated when we were studying nitrogen that the chief
source of that element was the atmospheric air. Indeed, it com-
prises nearly four-fifths of the atmosphere. In addition to nitro-
gen, we find an abundance of oxygen in the atmosphere. This
amounts to nearly one-fifth of the whole. In addition to these two
elements we find many other substances, both elementary and com-
pound, in the atmosphere, so that we must study this mixture of
gases with some thoroughness.

The meaning of the term cUmospJieric air is well understood.
It is that mixture of gases which surrounds our globe, and which is
carried along with it as it sweeps through space. The importance
of the atmosphere can be seen at once, if we recall that without it
the forms of life which are now extant upon the surface of the earth
would be at once exterminated.

Composition of the Atmosphere. — In order to determine the exact
composition of the atmosphere, we must make a quantitative analy-
sis of it. The oxygen in the air can be determined in several ways.
A measured volume of air can be passed over heated copper. The
oxygen combines with the copper, forming copper oxide. By weigh-
ing the tube containing the copper before the experiment, and
weighing the tube containing the copper and copper oxide after the
experiment, we know from the gain in weight the weight of the oxy-
gen in a given volume of air. Knowing the weight of a litre of
oxygen or of a litre of air, we can calculate at once the percentage
of oxygen in the atmospheric air.

Again, the oxygen can be removed from the air by inserting a
piece of phosphorus. This will combine with the oxygen and form
phosphoric acid. By measuring the original volume of the air, and
the volume after all the oxygen has been removed, we have the per^
centage of oxygen by volume in the atmospheric air.

285

PEIKCIPLES OF INORGANIC CHEiCSTRY

A third method of determining the ainouiit of oxygen in the air,
is to mix with a known volume of air a given volume of hydrogen,
and explode the mixture. All the oxygen will combine with the
hydrogen and form water, which, at the temperature of the experi-
ment, will be precipitated in the liquid form. From the contraetioa
in volume after the explosion, the amount of oxygen present can be
calculated. This last method is known as the eudiomelrk method.
Results by the different methods show that the oxygen in the air i^j
about 20.8 per cent by volume, and 23.0 per cent by weight

The question arises, Does the amount of oxygen pi*eaent remain
constant, or does it vary from place to place or from time to time?
While slight variations have been detected, pure air from different
parts of the globe, and in different altitudes, varies but slightly in
composition.

The remainder of the atmospheric air is nearly all nitrogen, a
number of other substances occurring in it in very small quanti-
ties* There are traces of atrbon dioxide in the air. The amount
can be determined by passing the air through a solution of barium
hydroxidCj and weighing the amount of barium carbonate precipi-
tated.

The small quantity of ammonia In the air can be determined by
passing a given volume of air through a solution of a standard acid, j
and determining how much of the acid is neutralized.

The air under all conditions contains water-vapor. The amount,
however, varies greatly fr^mi time to time and from place to place.
In certain regions far i-e moved from the sea, and over desert land,
the amount of water-vapor in the air is comparatively smalL Over
regions which are cluse to large bodies of water, the aniount of water-
vapor in the atmosphere may be quite considerable. To determine
the amount of waier-vapor in the atmosphere, it is only necessary to
pass a measured volume of air over some goo<l drying agent, such
as phosphorus pentoxide, and determine the increase in the weight
of the pentoxide.

Other substances may occur in th© atmosphere in very minute
quantities, such as ozone, liydrogen dioxide, oxides of nitrogen,
and the like, but the quantities are so small that they can, for all 1
pi'actical purposes, be disregrirded.

In addition to the constit uents already named, there are a num*
ber of rare elements which occur in the air in very small quantities*
These are the newly discovered elements, argon, helium, neon, krj^])- '
ton, and xenon. These elements we shall consider briefly a little
later.

THE ATMOSPHERIC AIR 237

Is the Air a Mixture or a Compoiind? — The question naturally
arises, Is the atmospheric air a chemical compound or a mechanical
mixture ? The fact that it has so nearly the same composition the
world over, would argue in favor of the oxygen and nitrogen being
in combination, forming a definite compound. This line of argu-
ment, however, is by no means conclusive, since we might easily
have the two gases mixed in essentially the same proportion in all
regions. Gases diffuse so rapidly that if there was any appreciable
difference in composition, it would soon become equalized by diffu-
sion from the region of greater to the region of less concentration.
There is, however, direct evidence which shows that the air is sim-
ply a mechanical mixture of oxygen and niti-ogen, and not a chemical
compound.

When air is shaken with water, the part which dissolves has a
very different composition from ordinary air. The latter contains
in round numbers four parts of nitrogen to one of oxygen, while air
which has been dissolved in water contains only 1.9 parts of nitro-
gen to one of oxygen. This is due to the fact that oxygen is much
more readily soluble in water than nitrogen. If air is a com-
pound of oxygen and nitrogen, the compound would dissolve as such,
and the air which would be dissolved by water would have the same
composition as ordinary air.

Again, chemical union is always accompanied, as we express it,
by thermal change. Oxygen and nitrogen mix in the proportion to
form air without any thermal change, and air is, therefore, not a
chemical compound.

Physical Properties of Atmospheric Air. — The specific gravity of
air varies slightly, just as the composition changes slightly. Under
the average conditions of zero degrees and 760 mm. pressure, one litre
of air weighs 1.293 grams. The pressure of the air, however, de-
creases very rapidly as we rise from the level of the sea, and a litre
of air on the top of a high mountain would weigh much less.

The question as to whether the air has an upper limit, or extends
indefinitely into space, has been much discussed. From the general
law of the apparently unlimited expansion of gases, in terms of
which a gas will occupy the entire space placed at its disposal, it
would seem that the atmospheric air must extend out indefinitely
into space, the density, however, becoming very small at no great
distance from the surface of the earth, and decreasing almost to the
infinitesimal at a comparative short distance.

Certain work, however, which has been done on the expansion of
very dilute gases, shows that when a certain dilution of the gas has

PRINGirLES OF INORGANIC CHEMISTRY

been reached it does not obey the ordinary law of expansion, but
its power to expand is greatly diminished. From this it is highly
jvrobEible that the atmosphere does not extend to an unlimited dis-
tance into space, but that there is an upper boundary to the earth's
atmosphere, which is perhapa only a few hundred miles or less from
the surfai^e of the earth.

Liquid Air, — \\g have seen that both oxygen and nitrogen can
be liquetied, aiul would expect, therefore, that atmospheric air, which
is essentially a mixture of these two gases^ could also be liquefied-
Such is the fact. The methml employed is based on exactly the
same principles which were maile use of to liquefy oxygen and
similar substances, Tbe nujst economical method consists in com-
pressing tlie air and removing the heat set free by a stream of cold
water. The compressed air is allowed to expand, when its tempei-ap
ture is very much lowered. It is then allowed to cool other com-
pressed air, which, in turn, is allowed to expand, and a still lower
temperature is produced. This is continued until a temperature is
reached at which the compressed air, when allowed to expand, be-
comes partly liquefied. In this process the air is allowed to expand
through a fine opening known as a veedle valve, when part of the
compressed air is liquefied and the remainder passes off as gas.

Liquid air has a slightly bluish color. When filtered from solid
carbon dioxide and ice it is transparent. It boils at — 190^ Aa
already stated, the liquid nitrogen, liaving a lower boiling-point than
liquid oxygen, boils off more rapidly, and the liquid remaining after
liquid air has been allowed to evaporate for a considerable time^ la
almost pure liquid oxygen.

Since liquid air is now manufactured on a commercial scale, it is
possible to use it on the lecture table for experiments at very low
temperatures* Indeed, it is by far the best means at our disposal
for protiucing temperatures in the region of — 180** to — IIK)^ At
these tenqieratures chemical activity is greatly diminisheiJ, and many
of the properties of many substances are greatly changed. Some due*
tile metals iH-come quite brittle, and can be easily broken. Flesh
lH*come9 brittle, and can be broken like thin glass. Xearly all liquids
ai'o converted into solids when immersed in liquid air. Mercury can
bo frozen lu a mould in the form of a hammer sufficiently hard to
rive a naih Alcohol is readily converted into a solid which resem-
bles semi'traiisparent ice,

In varuiim-jacketed b«lbs liquid air can be preserved for quite a
time. Its vapor-pressure is, however, so great, that vessels which
contain it muit be left open.

THE ATMOSPHERIC AIR 289

ARGON, HELIUM, KRYPTON, NEON, XENON

Argon (At, Wt. = 39.9). — These five elements have all been dis-
covered in the atmospheric air since the summer of 1894. Just
before this time Lord Rayleigh had observed that nitrogen obtained
from atmospheric air by removing all known constituents was
slightly heavier, volume for volume, than nitrogen prepared by heat-
ing ammonium nitrite. A litre of nitrogen obtained from the air
weighed 1.2572 grams, while a litre of nitrogen from ammonium ni-
trite, which was known to be chemically pui-e, weighed 1.2521 grams.
Ko one knew what this meant, but the fact was established beyond
question. The most probable explanation seemed to be that the
nitrogen from the air contained some impiirity which was heavier
than nitrogen. Acting upon this line of thought, Rayleigh and
Ramsay took up the problem from the chemical side. They deter-
mined to remove the oxygen from the air, then the nitrogen and
other known constituents, and see if anything remained.

They removed the oxygen from the air by passing it over red-hot
copper. The nitrogen was removed from the residue by passing
it over red-hot magnesium, the ordinary impurities having been pre-
viously removed. There remained a residue which spectrum analy-
sis showed to be a new substance, and which was a little less than
one per cent of the atmosphere. Rayleigh and Ramsay were not
able to make it combine with any known substance, and from its
chemical inactivity called it argon. Its vapor-density showed that
its molecular weight was 40. When cooled in liquid oxygen and
subjected to a pressure of 50 atmospheres, it liquefied at — 187^ It
solidified at — 189**.5. All of the facts known point to the element-
ary nature of argon, and there is not the slightest reason for sup-
posing that it is a compound.

Rayleigh and Ramsay next attempted to determine the number
of atoms in the molecule of argon.

ITumber of Atoms in the Molecnle of Argon. — There are several
methods for determining the number of atoms in a molecule of a gas.
One method is based upon the ratio between the specific heat of the
gas at constant-pressure and the specific heat at constant-volume.
That it would require more heat-energy to raise the temperature of
a given mass of gas, a certain number of degrees at constant-pressure
than at constant-volume, is obvious. When the gas is kept at con-
stant-pressure as the temperature is raised, it expands, doing work
driving back the atmosphere. If we represent the specific heat at
constant-pressure by C,, and the specific heat at constant-volume by

240

PRINCIPLES OF INOKGANTC CHEMISTRY

C^, wlien the ratio between theie two is ^SG, it has been sbown from
tlie kinetic theory of gases that the molecule must be inonatomic : —

§ = 1.66.

It would lead ns too far to deduce here this relation from the kinetic
theory.

The above described method is, on the whole, the one be^t known
in connection with the determination of the number of atoms in a
molecule of a gas,

Eayleigh and RamBay, however, made use of a method which \e
more convenient, especially when the quantity of substance at dia-
posal is not large. Instead of measuring the two specific heats of
argon — at coustant-presaure and at constant-volume — they simply
measured the velocity of Bound in tlie gas. There ia a comparatively
simple relation between velocity of sound in a gaa and the ratio
between the two specific heats of the gas, so that knowing the
former, the latter is easily calculated.

Kayleigh and Eamaay did not measure the velocity of sound in
the gas directly, but measured the wave-length of sound in the gas
by placing some hjcopodlum powder in the glass tube filled with the
gaS| through which the sound was passing. The lycopodium collects
at the points of rest, the nodes, and by measuring the distance
between two nodes we have the wavedength of sound in the gas
with which the tube is filled. Knowing the wave-length of sound
in the gas, and the pitch, we calculate at once the velocity. This is
KuiuU*3 method of determining the ratio of C^ to 0„ for any gas,

Kayleigh and liamsay found that the molecule of argon is mon-
atomic. The atomic weight of argon is, therefore, the same as its
molecular weight, 40»

Argon has also been found in certain minerals, and in the waters
of certain springs.

^ Helium (At Wt. = 4), Keon (At Wt. ^ 20), Krypton (At Wt

H =81.8), and Xenon (At Wt = 128). — Since the discovery of

^H argon, Kamsay has carried his investigations on the atmospheric air

^H much farther, and has discovered four new substances, all of which

^M appear to Ijc elementary. When air Is liquefied two of these escape,

^m - being very %*olatile^ helium and neon.

^B IMiftmf so called because it had been recognized by means of the

^m spectroscope as occurring in the sun, has also been discovered in the

^H waters of certain springs, and in certain ores of uranium. When a

^^^^ mixture of helium and neon is cooled in liquid hydrogen the neon is

THE ATMOSPHERIC AIR 241

liquefied while the helium remains a gas. Helium does not combine
with any known substance, its molecule is monatomic, and its boiling-
point somewhat lower than that of hydrogen. It has, then, the
lowest boiling-point of any known substance, and has thus far not
been liquefied. Its atomic weight, which is identical with its moleo-
idar weight, is 4.

Neon has an atomic weight of 20.

Krypton and xenon boil higher than air, and were, therefore,
found in the residue from the evaporation of a large amount of
liquid air. They were separated by the difference in their boiling-
points.

The atomic weight of krypton is 81.8, of xenon 128.0.

CHAPTER XVII

PHOSPHORUS (At. Wt. = 31.0}

Oconrrenoe and Preparation. — Phosphorus^ discovered by Brandt
in 1669, derives its name from the fact that it emits light, or, as we
say, is phosphorescent. It does not occur in the free state, but mainly
in the form of phosphates, and especially in combination with calcium
as the calcium salt. This is the compound of phosphorus which
occurs as apatite, phosphorite, etc., and the gieat " phosphate beds "
in the southern part of the United States are mainly calcium phos-
phate. Phosphorus also occurs in the bones of animals in the form
of the calcium salt, and most of the phosphorus of commerce is
made from this source.

Phosphorus is widely distributed through the soil in the form of
its salts, and especially of its calcium salt. This comes in part from
decomposing rocks which contain phosphates, and also from decom-
posing animal and vegetable remains. Plants in general take phos-
phates from the soil and build them up into their own structure.
Animals live largely upon vegetables, or upon other animals which
live on vegetable food, and thus secui'e the phosphates which they
so much need. The great phosphate beds are supposed to be the
remains of animals once living upon the earth.

Phosphorus is of fundamental importance to our highest func-
tions. It occurs in the brain, albumen, etc., and is essential to
mental activity.

Phosphorus is prepared from tricalcium phosphate, which has
the composition Ca8(P04)2 ; phosphoric acid, as we shall see, having
the composition H3PO4. If bones are used as the source of the cal-
cium phosphate, the organic matter is first destroyed by burning.

The tricalcium phosphate is treated with sulphuric acid, when
monocalcium phosphate is formed: —

Ca^O^)^ 4- 2 H2SO4 = 2 CaS04 4- CaCH^POOj.

This is then heated, when it passes over into calcium metaphoa-
phate : —

CaCH^OO, = 2 H,0 + Ca(POa),.
242

PHOSPHORUS 243

The calcium metaphosphate is then heated with a mixture of sili-
con dioxide (SiOg) and powdered charcoal. The following reaction
takes place : —

Ca(P03)j + 5 C + SiO, = 5 CO + CaSiO, + 2 P.

In heating the phosphate with carbon and sand the electric fur-
nace is now frequently used. In this case it is not necessary to
transform the phosphate into metaphosphate in advance; but the
phosphate can be heated at once with carbon and sand, when the
reaction expressed by the following equation takes place : —

2 CasCPOO, 4- 10 C + 6 SiO, = 10 CO 4- 6 CaSiO, + 4 P.

The phosphorus obtained by the above method is contaminated
with various substances. To remove the impurities it is filtered
through chamois skin while liquid under water, redistilled, and cast
into sticks, in which form it appears on the market.

Properties of Phosphoriu. — Phosphorus is a soft solid, with a
slightly yellowish tint. In contact with the air it combines readily
with the oxygen, forming an oxide of phosphorus. When phosphorus
is brought in contact with oxygen a part of the latter is transformed
into ozone, as we saw when we were studying ozone.

Phosphorus combines with most of the elements, and with such
elements as iodine and bromine with great vigor.

Phosphorus is an extremely poisonous substance, and in working
with it precaution must be taken not to inhale its vapors.

While phosphorus when warm has a soft, waxy consistency,
when cold it is quite brittle.

Phosphorus melts at 44°.5, forming a yellowish liquid. When
heated in an atmosphere free from oxygen it boils at 290**. When
heated in contact with oxygen it takes fire at about 50°.

Phosphorus in the form of vapor at low temperatures is com-
posed of molecules of P4. As the temperature rises these break
down into molecules of P^

Phosphorus exists in more than one form, there being no less
than four allotropic modifications. Ordinary yellow pliosphorus has
already been briefly described. It dissolves readily in carbon disul-
phide, from which it crystallizes when the solvent is evaporated.

When yellow phosphorus is allowed to stand under water for a
long time, exposed to the light, it passes over into a red modification.
Red phosphorus is easily prepared by heating the yellow phosphorus
to 250** in an atmosphere free from oxygen, or to 300** in a vacuum
for a few minutes. Red phosphorus is an amorphous powder, and

2U

PRINCIPLES OF INORGANIC CHEMISTRY

fi

to the eye resembles in no respect the ordinary variety. The differ-
ence between the two is really deep-seated. Red phosphorus is
much less active cheniically than yellow. When heated to 200° in
the air, it does not take fire. When brought in contact with elements
and compounds with which yellow phospliorus unites at once, it does
not combine with thera. Red phosphorus is not soluble in carbon
disulphide, and is much less poisonous than the yellow variety.

When red phosphorus is heated to ^GO'' in an atmospliere of carbon
dioxide, it passes over quantitatively into the yellow modification.

We have in these two varieties of phosphorus a case somewhat
analogous to that met with in the two modifications of oxygen and
suli>bur. Certain differences are, however, obvious. We saw in the
case of sulphur that the real difference between the properties
of the two rnoflifications was to be sought for in the different
amounts of intrinsic energy present in the two modifications*

Exactly the same relations were discovered in the case of oxygen
and ozone. We slionld, thereforej naturally ask whether there is any
similar relation between the two modifications of phosphorus. Do
the different modifications contain different amounts of intrinsic
energy? This can be answered by burning the different modifica-
tions in oxygen, when they yield the same end product, phosphorus
pentoxide, PjO^. The results of thermochemical measurements show
that when yellow phosphorus is transformed into i^d there are
27,300 calories of heat set free, and this is approximately the thermo-
chemical equivalent of the difference between the intrinsic energies
of these two modifications of phosphorus.

When red phosphorus is heated in evacuated tubes to 360"*, or
ixed with metallic lead and highly heated for a considerable time^
another motlification of phosphorus appears. The molten lead when
allowed to cool is covered with black crystals, and these are also
contained within the solidified mass of lead. This form of phos-
phorus is kno^m as cnfslalUz<ul, m<*/al/ic, or hktck phosphorua.

Another modification of phosphorus has been prepared by con-

'Sensing vapors of phosphorus by means of ice-water in an atmosphere

o! hydrogen. The water becomes covered with a white powder, and

this is Khite phoHphonis. It has projjerUes quite different from

ordinary yeUow phosphorus*

It \& iiuposBible to say at present whether '* black phosphoma "

aad "white phoaphonia" have different amounts of energy in their

TSwAecnies, and each ^ different amount from all other mmUfications,

i^iii»lVtv«C!iissary thermochemical measurements have not yet been

"•46. l^m mhalt is kaown, however, in general concerning the

PHOSPHORUS

245

energy relations which obtain for allotropic modifications of an ele-
ment, it seems very probable that different amounts of heat would
be set free by burning the same amount of these different modifica-
tions of phosphorus to the same end product.

A characteristic of ordinary yellow phosphorus is, that it emits
light when placed in the dark. This is, undoubtedly, closely con-
nected in some way with the oxidation of the phosphorus, since
substances which hinder or prevent the oxidation, reduce or prevent
the light-giving power of the element phosphorus.

Compounds of Phosphorus with Hydrogen. — Phosphorus forms
three compounds with hydrogen, having, respectively, the following
compositions : PH^, PH„ and PjH. At ordinary temperatures the
first is a gas, the second a liquid, and the third a solid.

Gaseous hydrogen phosphide, or phosphinCy is prepared by the
action of caustic potash on phosphorus in the presence of water : —

3KOH-f.4F-f 3HjO=3HjKPO,+ PH,.

Phosphine produced by this method always contains a little of
the liquid compound PHj, which renders it spontaneously inflamma-
ble. The preparation of phosphine by the above reaction is a very
beautiful experiment.

Arrange a flask A, as in Fig. 28, and introduce a few grams of
caustic potash, dissolved in 15 or 20 cubic centimetres of water. Add

Fio. 28.

a few small fragments of phosphorus. Connect an escape tube as
shown at B, allowing it to dip beneath the water in the vessel C

246

PRINCIPLES OF INORGANIC CUEMISTRY

This water should be kept warm, in order that the end of the tube
may not become stopped up with phosphorua which will distil
over from the flask- The flask ^4 is €onuect«d with a hydrogen
generator by means of the glass tube D. When the apparatus has
become filled with hydrogen from the generator, the solution of
caustic potash is gently heated, and phosphine quickly begins to
escape. The bubbles, as they come in contact with the air, take fire
spontaneously and burn, the phosphorus being oxidized to an oxide
of phosphorus, and the hydrogen to water. The products of com-
bustion rise in beautiful rings, which increase in diameter as tbey
ascend, and all together the effect is very beautiful,

Phosphine is obviously the phosphorus analogue of ammonia,
Phosphine PH^j, ammonia NH^. Like ajumonia, it can combine with
the hydrogen acids of the halogens, Ibe combination taking place by
direct addition. The compound with hydriodic acid is formed as
follows:- PH, + HI = 1'HJ,

the compound, pho^yfionium iodide^ being a white, beautifully
crystalline substance, which is not very stable even at ordinary
temperatures.

The gas phosphine, which is not spontaneously inflammable, can,
however, be burned. It is poisonous, and ail work with it sliould be
done under the hood. It is liquefied at 85", and passes over into a
solid at ^ 132".5.

The liquid compound PHj or (PHa)^ is formed along with the
gaseous when the latter is made by the method Juat described. It is
the presence of this compound which makes the gas set free in the
above experiment spontaneously inflammable. The liquid readily
decomposes into the gas PH^,, and the solid P.jH or (l^H),,

Compounds of Fhoiphorue with Oxygen and Hydrogen. — Phos-
phorus formtf a iuiml>er of compounds with oxygen, indeed four in
^L These are phosphorus suboxide, P^O \ phosphorus sesquioxidCj
PjO^; phosphorus tetroxide, P^O^; and phosphorus pentoxide, PjO^.
We have seen that when phosphorua is oxidized on the air, the
pentoxide is formed. The other oxides result from tlie incomplete
oxidation of the phosphorus,

PhoHphoru^ suboxide is formed by the action of soilium hydroxide
in a mixture of water and alcohol on phosphorus. When the solution
is acidified after the action is over, the suboxide is precipitated.

The tetroxide is formed by heating the sesquioxide to about 400',
Under these conditions, the sesquioxide breaks down into the
tetroxide and phosphorus*

PHOSPHORUS 247

The sesquioxide of phosphoruSy PjOg, is formed by the incomplete
oxidation of phosphorus. When phosphorus is burned in a slow
current of air, which does not furnish enough oxygen to convert it
into the pentoxide, it forms the sesquioxide, which has the compo-
sition PjOa, but may have the formula Vfi^, The sesquioxide readily
takes up oxygen and passes over into the pentoxide.

Phosphorus pentoxide, P2O5, is formed by the oxidation of phos-
phorus in the presence of an excess of oxygen. It is a beautifully
white compound, which has remarkable power to combine with
water. Indeed, it is the best drying agent at the disposal of the
chemist. Phosphorus pentoxide is the anhydride of an acid. When
it combines with the maximum amount of water, it forms phosphoric

^^^ • — PA + 3 H,0 = 2 H3PO,.

This brings us to the acids of phosphorus, of which there are
several.

The Acids of Phosphorus. — Phosphorus combines with oxygen
and hydrogen, forming no less than seven compounds which are
acids. These are : —

Phosphoric acid . . . . . H3PO4

H,P,0,
HPO3

H,PA
HaPOa
HPO,
HaPO,

Pyrophosphoric acid
Metaphosphoric acid
Hypophosphoric acid
Phosphorous acid .
Metaphosphorous acid .
Hypophosphorous acid .

The most important of these, by far, is ordinary phosphoric acid,
or orthophosphoric acid.

Orthophosphoric Acid, H3PO4. — Orthophosphoric acid is formed,
as already stated, by dissolving phosphorus pentoxide in water. It
is also formed by the direct oxidation of phosphorus by strong
oxidizing agents such as nitric acid. It is in the form of salts of
this acid that phosphorus occurs in nature, the calcium salt, Ca3P04,
being the compound in which phosphorus occurs in the great phos-
phate beds. Wlien this salt, which is insoluble in water, is treated
with an excess of concentrated sulphuric acid, it is converted into
soluble compounds, the compound formed depending upon the amotint
of sulphuric acid present. When normal calcium phosphate is treated
with one molecular weight of sulphuric acid the following reaction
takes place : —

CaaCPO^), + H,S04 = CajH,(P04), + CaS04.

The salt formed is seconda)^ calcium phosphate.

248 PRINCIPLES OF INORGANIC CHEMISTRY

When two equivalents of sulphuric acid are used the salt formed
is the primary calcium phosphate, and in the sense of the following
equation : —

Ca^CPOO, + 2 H,S04 = 2 CaSO^ + CaH^CPOjr.

When three equivalents of sulphuric acid are used the following
reaction takes place : —

CaaCl^OOa + 3 H,S04 = 3 CaS04 + 2 H3PO4,

giving free orthophosphoric acid.

The nlK)ve reactions are extensively made use of to render ordi-
nary normal calcium phosphate soluble in water, so that plants can
oMmu it and take it up into their tissues. These are the funda-
ni^^uUl rt^aotions employed in the manufacture of commercial fertil-
tw^r oithor fiH)m phosphate rock or from animal bone.

\% xm obvious from the above that phosphoric acid is a tribasic
tkC'Mx which forms three series of salts : —

*lMu» li«Wiarv or normal phosphates, having the composition M3PO4,
whor^ M i» a univalent metal; the secondai-y phosphates, having
\\<p \HMU|HViiti<m nMjP04; and the primary phosphates, having the
^H^u|Hvi\tiou n|MP04.

IttMOOilltlOA of Phosphoric Acid. — Since phosphoric acid is a tri-
kiAakio m'uU it must dissociate into three hydrogen ions. The com-
y\^}> \Usk!t\H^iatiou of phosphoric acid in the following sense is very
UiWoult t\> t>ffwt : —

lI,rO, = H,H,H,P04,

niiUH^ |vh\^l4uu*\o aoid is a comparatively weak acid, as is shown by
it\\^ MK^Nviug wmluotivity results: —

4

,ly (250)

72
146
297
855

l*liu^plu»iio uoi\l vUasooiatea first in the following sense: —
U,rc\=iH, H,P04.

PHOSPHORUS 249

When more water is added, or when these hydrogen ions have
been used up by a base, the second hydrogen begins to split off in
the ionic state : — ^

H^04 = H,HP04.

It is not until these hydrogen ions have been used up, or very
great dilution has been reached, that the third hydrogen ions begin
to split off: —

HP04=H,P04.

We can now understand why it is quite an easy matter to pre-
pare mono- or primary sodium phosphate by adding sodium hydrox-
ide to phosphoric acid, and also why the secondary salt can be
readily prepared. It is, however, not as simple a matter to obtain
the tertiary salt in pure condition.

It is a general rule that the salts of weak acids are acted upon
by water to a greater or less extent, being broken down into the cor-
responding acid and base. Take tertiary, or normal sodium phos-
phate, Na3P04. When this is acted upon by water the following
decomposition takes place to some extent : —

NasP04 + H,0 = Na, Na, HPO4 + Na, OH.

This kind of dissociation is known as hydrolytic diMociation.
This takes place to a greater or less extent whenever the salt of a
weak base with even a strong acid, or even a strong base with a
weak acid, or still more when the salt of a weak base with a weak
acid is brought into the presence of water. This is the explanation
of the alkaline reaction shown by such compounds in water. The
hydroxyl ions set free as the result of the combined action of
hydrolytic and electrolytic dissociation, give their characteristic
alkaline reaction with all indicators sensitive to alkalies. Even the
secondary sodium phosphate is hydrolyzed to a slight extent, and
shows a feebly alkaline reaction.

Some of the salts already met with undergo hydrolytic disso-
ciation in the presence of water, notably the sulphites. We shall
meet mauy more examples of this kind of dissociation before the
subject is ended.

Detection and Determination of Phosphoric Acid. — Phosphoric
acid forms a number of insoluble salts with the heavy metals.
Some of these have characteristic color. The silver salt, Ag,P04, is
yellow. Phosphoric acid is detected when present in very small
quantity by adding a nitric acid solution of ammonium molybdate

250 PRINCIPLES OF INORGANIC CHEMISTRY

(a compound which we shall study later), when a complex, yellow pre-
cipitate is formed, known as ammonium phospho-molybdate. This
compound is soluble in ammonia, and when a mixture of ammonium
sulphate and magnesium sulphate is added to the ammoniacal solu-
tion, all the phosphoric acid is precipitated quantitatively as the
ammonium magnesium salt — NH4MgP04. When this is heated it
loses ammonia and water 2 NH^MgPO* = 2 NH, + H,0 -f Mg,P A,
and forms the pyrophosphate of magnesium, which is a stable sub-
stance and can be easily weighed.

In determining the phosphoric acid in a commercial phosphate,
which contains the primary, secondary, and tertiary salt, three
determinations are necessary. The primary salt is soluble in water,
the secondary salt in an aqueous solution of ammonium citrate, while
the tertiary salt is soluble only in acid. The phosphoric acid in
each solution is determined as described above. The water soluble
plus the citrate soluble constitute the "available" or "soluble"
phosphoric acid, while the remainder is "insoluble" phosphoric
acid.

Pyrophosphoric Acid, B^flj. — Pyrophosphoric acid is formed
from phosphoric acid by loss of water : —

2H3P04 = H,04-H4PA.

This reaction takes place between 250** and 300^ Salts of this acid
are easily obtained by heating secondary phosphates : —

2 HM,P04 = H,0 4- M4P A.

When the lead salt of this acid is treated with hydrogen sulphide,
the lead sulphide is precipitated and free pyrophosphoric acid is
formed. The presence of this acid is readily detected since its sil-
ver salt Ag4PA is pure white, while the silver salt of orthophos-
phoric acid is yellow. Pyrophosphoric acid, since it contains four
hydrogen atoms, might yield salts in which one, two, three, and four
of these hydrogens were replaced, as we say, by metals. There are,
however, only two classes thus far known; those in which two
hydrogen atoms are replaced, and those in which four are replaced.
Salts in which one and three hydrogens are replaced, if capable of
existence, have not thus far been prepared. There is no reason on
the face of it why they should not be made.

Metaphosphoric Acid, HPOs. — Metaphosphoric acid is formed
when normal phosphoric acid is heated higher than is necessary to
form the pyroacid. When a temperature of about 400"* is reached

PHOSPHORUS 251

the second molecule of water passes off from the normal acid, and
the metaacid results: —

H,P04=H,0 + HP0,.

It is also formed when phosphorus pentoxide takes up one molecule
of water : —

PA + H,0 = 2HP0,.

Salts of this acid are formed when primary phosphates, MHsP04,
are heated : —

MH^O^ = H,0 + MPO3.

Metaphosphoric acid, on account of its vitreous appearance, is known
as glacial phosphoric acid. When allowed to stand in contact with
water it takes up the water, forming orthophosphoric acid.

It is detected by the fact that its barium salt is a white, insoluble
solid.

Hypophosphoric Acid, H4PsOa. — Hypophosphoric acid is formed
as one of the products of the action of phosphorus on an insufficient
supply of air. When the insoluble barium salt is treated with sul-
phuric acid the free acid is formed.

Phosphorous Acid, HsPOj. — Phosphorous acid is formed by the
action of phosphorus on moist air. Also by the action of water on
a chloride of phosphorus with which we shall soon become familiar,
phosphorus trichloride : —

PCla + 3 HjO = 3 HCl -f H3PO3.

The acid can be obtained from the solution in the form of crystals
which melt at 70**.

Phosphorous acid contains three hydrogen atoms, and would,
therefore, be expected to be a tribasic acid. The fact is, it is only
dibasic, the salts richest in metal having the composition MjHPOa.
This is to be explained in terms of its dissociation as follows : The
first stage in the dissociation of this substance is represented by the
following equation : —

H8PO, = H,H,P03.
The second stage is represented thus : —

H,P03 = H, HPO3.

It is impossible to go farther and split off the last hydrogen atom
as an ion. This is not wholly unlike phosphoric acid. We saw

£52

PRINCIPLES OF INORGANIC CHEMISTRY

th^t the first hydrogen atom readily passed into the ionic condition
in the pN^sence of water; the second split off much less easily;
while the third was converted into an ion only with the greatest
ditBoultv. In the case of phosphorous acid, it is impossible to cause
the thinl hyiirogen to pass into the ionic state.

■tt^^Ptphorous Aoid, HPO,. — Metaphosphorous acid is stated
to W fi>rmotl when phosphine undergoes slow oxidation.

HjfpophPtphorous Acid, HsPOi. — Hypophosphorous acid is formed
b^ tlu^ Action of an alkali on phosphorus. The reaction was re-
l^r^l t\^ wli^u we were dealing with the preparation of phosphine.
Thf» frtHi^ aoid is obtained by treating the barium salt with sulphuric
l^n\lv Hypophosphorous acid readily takes up oxygen, forming phos-
)4hMrio <i^oid. It is, therefore, a strong reducing agent.

Hy iHi»))hosphorous acid contains three hydrogen atoms, and might,
Ih^wfoirii^* be supposed to be tribasic. The fact is, that it is not even
4iW.Hiv\ It is only monobasic. The salts have the composition
AlH^kHV The acid must, therefore, dissociate as follows: —

H^O, = ii,H,PO^

iKi^ iou H4H\ not being capable of further dissociation.

iUfl|tka iit the Acids of Phosphorus. — The relative strengths
\4 iS uxuuWr of the acids of phosphorus can be seen by comparing
Ihi^ii^ ooiuhiotivities. Take them in the order of increasing amounts
uf oxyuvu in the molecule: —

)6

Hi
U^4

H.ro.

n»PO,

H,PO.

Mi^

fiF

M^y

lai

121

60

im

176

90

»u

241

146

au

298

225

aat)

329

297

m\

339

335

'Ihovo IH \^ r«tht»r rtMuarkablo relation brought out by the above
i>\iUU\»lo>4. U iH gouorully true that increase in the amount of oxy-
^\^\\ \\\ the UioUnnilo inormsos the strength of the acid. Here, ex-
\w\\\ I ho op|HW4ito in truo; the more oxygen in the molecule, the
\^oak\ I \\\\\ uoi\li os|KHMally in the more concentrated solutions.

PHOSPHORUS 258

COMPOUNDS OF PHOSPHORUS WITH THE HALOGENS

Photphorns Trichloride, PCI,. — When chlorine gas is passed over
an excess of phosphorus, in an atmosphere free from oxygen, the two
combine and form phosphorus trichloride : —

2P-f 3Cl, = 2PCl3.

This compound, when brought in contact with water, decomposeSi
forming phosphorous acid and hydrochloric acid : —

PCls + 3 H,0 = 3 HCl + HaPO,.

Phosphorus trichloride is a colorless liquid boiling at 76**, and
passing into the solid form at — 112°. Phosphorus has the power
to take up more chlorine and form phosphorus pentachloride, and to
take up oxygen and form phosphorus oxychloride.

Phosphorus Pentachloride, PCI5. — The pentachloride of phospho-
rus is formed, as stated above, by the action of chlorine on the tri-
chloride of phosphorus ; also by the direct action of an excess of
chlorine on phosphorus. Like the trichloride it is readily decom-
posed by water, forming phosphoric and hydrochloric acids ; —

PCI5 + 4 HjO = 5 HCl + H3PO4.

With a small amount of water it undergoes partial decomposi-
tion, yielding the oxychloride of phosphorus and hydrochloric
acid : —

PCI5 + H,0 = 2 HCl 4- POCV

Phosphorus pentachloride is a white solid, which, at an elevated
temperature, passes into vapor without melting. The vapor of phos-
phorus pentachloride is especially interesting in connection with the
validity of Avogadro's law. Its vapor-density is less than would
correspond to the formula PCI5. This substance and ammonium
chloride were held up as especially prominent examples of com-
pounds which did not obey the law of Avogadro, as we saw when we
were considering exceptions to this law.

The explanation which was furnished by experiment was, how-
ever, entirely satisfactory. When phosphorus trichloride was vola-
tilized it underwent partial decomposition : —

PCl5 = PCl3 4-Cl2.

The presence of chlorine in the vapor was shown by its green cplor.

ii4 FR1>'C1PLES OF INORGANIC CHEMISTRY

AYheu phosphorus pentachloride is Tolatilized, howeyer, in the
pn^^wiKV of an exet^ss of either of its decompositioD products (PCI3
\v l'l<.\ its vajH>r has a deusity which corresponds to the formula
IH^I^; ju*5 a;j when ammonium chloride is volatilized in the presence
of jiu 0XK*t»*s of either of its decomposition products (NH, or HCl), it
jjive* normal uuJwular weight.

Tlu**o aiv excellent examples of the effect of mass on chemical
an'tivuv. When the mass of any of the dissociation products is
inon^^iHthl. tho dissiK^iation is driven back as we say, i.e. the constitu-
ont8 nMuain uniteil or, if once separated, unite again to form the
ori)(inal iHMU|Hnuul. The law governing the influence of mass has
alroaily Uhmi takon up.

Phosphomt Oxyohloride, POCls. — Phosphorus oxychloride, a liquid
with a vory |HMietrating and nauseating odor, is formed, as we have
alroaily stHMU by tlie action of water on the pentachloride of phos-
phorus, tlus^t as it is the product of the first stage in the decom-
jKwition of phosphorus pentachloride with water, so when treated
with waic»r it undergoes further decomposition, yielding hydro-
chloric aoid and phosphoric acid: —

rOCl, -h 3 H,0 = 3 HCl -h H3PO4.

Phosphorus also forms compounds with bromine, iodine, and
fluorine.

CHAPTER XVIII

ARSENIC (At. Wt = 75.0)

Ooonrrence and PreparatioxL — An element closely allied chemi-
cally to phosphorus is arsenic. That there are marked differences,
however, will appear as the following chapter develops.

Arsenic does occur in the free condition. It is generally in com-
bination with the metals, either directly as in the compound with
iron, FcjAs^ or with sulphur as in arsenical pyrites, Fe AsS.

Arsenic is generally obtained from its compounds by simply
heating them ; the arsenic, being volatile, passes off as vapor.

Properties of Arsenic. — Arsenic is a solid at ordinary tempera-
tures, gray in color and very brittle. It combines with oxygen,
slowly at ordinary temperatures, but rapidly at elevated tempeni-
tures, with evolution of light and heat. In an atmosphere of
chlorine it burns readily, forming a chloride of arsenic. Arsenic
in the form of vapor is composed of molecules of AS4.

Compound of Arsenic with Hydrogen — Arsine, AbRq. — Arsenic
forms with hydrogen the compound AsH,, which is analogous to
ammonia, NH,. It is formed by the action of nascent hydrogen on
compounds of arsenic. When we have a compound of arsenic in the
presence of zinc and an acid, arsine is formed. If the compound is
ordinary arsenic trioxide, As^Os, this is reduced by nascent hydrogen
as follows : —

AsjO, -t- 6 Hj = 3 HjO -t- 2 AsHg.

When arsine is heated it is broken down into its elements, arsenic
and hydrogen. W^hen arsine is burned and a cold object introduced
into the flame, arsenic is deposited upon the object. These reactions
are made use of for the detection of arsenic. Marsh's method for
detecting arsenic consists in reducing the arsenic compound to arsine
and burning the arsine.

The apparatus used is shown in Fig. 29. Into the flask A some
zinc which is perfectly free from arsenic is introduced. Upon this
is poured some pure, dilute, sulphuric acid, which acts upon the
zinc, generating hydrogen. The tube C, containing calcium chloride,
is introduced for the purpose of drying the gas. When the appa-
ratus has become filled with hydrogen, the gas is ignited as it escapes

266

256

PRINCIPLES OF INORGANIC CHEMISTRY

fviiru i\w end of tlje tube i?. The substance which is supposed to
coutaiti arsimiu is ilbsolved in hydroeliloric acid, and added to the
contents of ttiu t!ask A If avsetiic is present, arsine will be formed
and will escape niixod with the hydrogen. If arsine is present, the
color of til© hydrogen flame will change very perceptibly. The

5

a:

KmL

ilniHli oolorless flame of the hydrogen will become milky-white in
tolor with a greenish-blue tint. If now the tube is heated, the
amitie will be decomposed and a mirror of arsenic will be de[>osited
Wfim the walls of the tube* A cold evaporating dish inserted into
ib» tanit^ will become covered with a layer of arsenic.

la Older that this test should be of any value» all the materiala
IMHil Mist be perfectly free from arsenic. In testing for arsenic the
gpitllil pre^'AUtion should be taken to secure this result.

VHMnHJKI^ OF ARSENIC WITH OXYGEN AND HYDROGEN

fftHfiiiillt of Arsanio with Oxygen. — Arsenic forms two com-
^m>4AIhIii wilh ^utyijcn* One of these has the composition AsjO,, and
l4 )liH^^ II A* *r$<Miic Irioxide, or arseniotis oxide ; the other has the com-
)4M(MUi^i Is* V i^i*d u known aa arsenic pentoxide» or arsenic oxide,

^xmikk^ Trio^side. Mfi^^ is formed when arsenic is oxidized either
ly iHiiiuti^ ia ih« air, or by some strong oxidizing agent such as

»i** \\\ide or arsenic trioxide is a white solid, which passes

ARSENIC 267

gradually into the crystalline condition. Arsenic trioxide crystallizes
in more than one form, and thermal changes take place when one
form is transformed into another. It, therefore, exists in allotropic
modifications. When heated, arsenic passes at once into vapor with-
out melting. If, however, it is subjected to a higher pressure it can
be melted. In this case, as in so many others, the melting-point is
higher than the boiling-point under atmospheric pressure.

The vapor-density of arsenic trioxide varies with the temperature.
Below 800° the vapor-density corresponds to the double formula,
ASfOa. As the vapor becomes heated higher and higher the vapor-
density becomes less and less, until at from 1700® to 1800° the vapor-
density corresponds to the formula As^O,. This is another substance
whose molecule in the form of vapor is more complex at lower tem-
peratures, and dissodatea into simpler molecules as the temperature
rises. The molecular weight of arsenic trioxide in nitro-benzene has
been determined by the boiling-point method, and found to correspond
to the double formula, AS4O6.

Arsenic trioxide is in a sense the anhydride of araenious acid,
just as phosphorus trioxide is the anhydride of phosphorous acid.
Arsenic trioxide is, however, only slightly soluble in water.

Arsenic trioxide is the form in which arsenic comes most fre-
quently on the market. It is known as white arsenic, and is very
poisonous. The best antidote for arsenic poisoning is a mixture of
magnesia and ferric hydroxide. The arsenic is precipitated by this
mixture probably in the form of ferric and magnesium arsenite, which
is only slightly soluble.

Arsenic Pentoxide, A&A. — Arsenic pentoxide cannot be formed
like phosphorus pentoxide by burning the element in oxygen. Indeed,
if arsenic pentoxide is heated to a fairly high temperature, it breaks
down into arsenic trioxide and oxygen. It is prepared by removing
water from arsenic acid, and is, therefore, the anhydride of this acid.
It is also prepared by heating arsenic trioxide with some strong
' oxidizing agent such as nitric acid.

Arsenious Acid, H^AsOs. — This acid is not known in the free con-
dition. There are, however, three classes of salts known, depending
upon whether one, two, or three of the hydrogen ions have given
their electrical charges to the metal atoms which have entered the
compound. This acid can apparently lose the elements of water and
form metaarsenious acid : —

HjA^sOs = H,0 -t- HAsO,.

At least salts of this acid are known.

258 PRINCIPLES OF INORGANIC CHEMISTRY

Anenio Add, H9A8O4. — When ordinary white arsenic, or arsenic

trioxide, is heated with some strong oxidizing agent sueh as nitric

acid or aqua regia in the presence of water, it is oxidized to arsenic

acid. The reaction consists in the direct addition of oxygen and

water : — ^

AsjOs 4- O2 + 3 HaO = 2 H8ASO4.

The acid is known in solution as a syrupy liquid, and in the solid
form as white needles. Like phosphoric acid it forms three series of
salts, primary, secondary, and tertiary. It must, therefore, dissociate
in the three stages : —

H8As04 = H, HjjAsG^,
H,As04 = H, HASO4,
HAs04 = H, Ar04.

Arsenic acid, like phosphoric acid, loses water in stages forming
the pyro-, and meta- acids. \Vhen two molecules of the acid lose one
molecule of water the pyro- acid is formed : —

2 H,As04 = HjO + H4ASA.

When one molecule of the acid loses one molecule of water the
mttaaraenic acid results : —

H8ASO4 = H,0 + H AsOg.

When primary salts of arsenic acid are heated the following
reaction takes place : —

H,M ASO4 = H,0 4- I^LA^sOa.

When secondary salts are heated they yield pyroarsenaJtea : —

2 HM,As04 = H,0 + MjAsjOy.

Here again the resemblance between arsenic and phosphorus
appears.

Oompoundt of Anenio with the Halogens. — Arsenic forms a num-
l)er of compounds with the halogens. The best known is the tricMo-
ride^ formed by the action of hydrochloric acid on arsenic trioxide,

AsA -h 6 HCl = 3 H,0 + 2 AsClj,

or by the direct union of arsenic and chlorine, which readily takes
place. Arsenic trichloride is decomposed by water as follows : —

AaCU -f 3 H,0 - 3 HCl + H.AsO,.

ARSENIC 259

This reaction seems to be exactly the reverse of the above. "When
arsenic trioxide is heated with hydrochloric acid, arsenic trichloride
and water are formed. On the other hand, when the trichloride
is heated with water, hydrochloric acid and ai*senious acid are
formed.

Here we have again a good example of the effect of mass on
chemical activity. In order to have the first reaction take place^ a
large amount of hydrochloric acid must be used. In order to effect
the second reaction a large amount of water must be present. The
effect of mass here is such as to condition the way in which the
reaction proceeds.

Compounds of Arsenic with Sulphur. — Arsenic forms no less than
three compounds with sulphur.

Arsenic disulphide As^S^ occurs in nature as the mineral realgar*
It can also be prepared by fusing together the two elements in
equivalent quantities.

Arsenic trisulphide ASiS^ occurs in nature as the mineral orpi-
ment. It is formed by the action of hydrogen sulphide on arsenious
acid : —

2 H, AsOs + 3 HjS = 6 H2O + As,S,.

It can be formed also by fusing two equivalents of arsenic with
three of sulphur. This compound, on account of its fine yellow
color, was formerly used as a pigment.

Arsenic pentasulphide As^n is formed when hydrogen sulphide is
conducted into a cold, hydrochloric acid solution of arsenic acid : —

2 H3 ASO4 4- 5 H3S = AS2S5 4- 8 HA

It is also obtained by fusing arsenic with an excess of sulphur, when
the two combine and form the pentasulphide : —

2 As -f 5 S = AS2S5.

The excess of sulphur can be dissolved in carbon disulphide and
removed. On account of its insolubility in water, it is the form in
which arsenic is usually precipitated and weighed in quantitative
analysis.

Snlpho-salts of Arsenic. — Arsenic forms with sulphur and the
alkali metals, salts of acids having the composition HgAsSs and
H8ASS4. These acids are the sulphur analogues of arsenious acid
HjAsOg and arsenic acid H8ASO4. The sulpho-acids or thio-acids
are themselves not known, but certain salts are well-characterized
substances.

200 PRINXMPLES OF INORGANIC CHEMISTRY

Thus, when arsenic trisulphide is treated with sodium, potassium,
or anuuoniuni sulphide, the two combine as follows, M representing
tho alkali metal : —

3 MjjS + As^s = 2 MjAsSj.

This is obviously the salt of sidpharaenious or thioarsenious acid, and
is known as a sulpharsenite or thioarsenite.

Similarly, when the pentasulphide of arsenic is treated with an
alkaline sulphide, the two combine : —

3 MjS + AsA = 2 M3ASS4.

There is formed the salt of sulpJiarsenic acid or thioarsenic acid,
and this is known as a sulpharsenate or thioarsenate.

The analogy between the sulphur acids of arsenic and the oxygen
acids can be carried still farther. Just as w^e have salts of sulphur
acids corresponding to arsenious and arsenic acids, so, also, we have
salts of a sulphur acid corresponding to metaarsenic acid, IIAsOj.
When the trisulphide of arsenic is treated with a polysulphide of an
alkali metal, we have : —

M,Ss + AsA = 2MAsS;„

which is a sulphometaarsenate.

The above compounds of arsenic with sulphur and the alkali
metals, are of fundamental importance in separating arsenic from
other elements. Arsenic is precipitated as the sulphide, along with a
Btonber of other sulphides, by means of hydrogen sulphide. The sul-
phide of arsenic is soluble in the polysulphide of ammonium, forming
e&>^vssftlts; and is thus separated from most other substances.

CHAPTER XIX

ANTZMON7 (At. Wt. = 120.2)

Oeenrrenoe and Preparation. — Another element which presents
many chemical analogies to phosphorus and arsenic is antimony.
Antimony occurs in nature chiefly as the trisulphide, SbjSs, which
is the well-known mineral atihnite. It also occurs in combination
with arsenic and also with oxygen.

Antimony is prepared from stibnite by roasting out the sulphur.
The sulphide is heated in the air, when the sulphur is converted into
the dioxide, and the antimony into the trioxide or sesquioxide, SbsO^.

The oxide is then reduced with carbon : —

2 SbA + 3 C = 3CO2 + 4 Sb.

Antimony sulphide is sometimes heated with iron, when the iron
combines with the sulphur forming iron sulphide, and antimony is
set free.

Properties of Antimony. — Antimony is a bluish-white solid with
metallic lustre. It melts at 630^ and boils at 1450^. It combines
with oxygen at elevated temperatures, but not at ordinary tempera-
tures. Like arsenic it combines readily with chlorine at ordinary
temperatures. When a piece of antimony highly heated in the air
is thrown upon white paper, it continues to run about over the surface
of the paper, leaving tracings which are often very beautiful.

The vapor-density of antimony decreases with rise in temperature.
At the boiling temperature the vapor of antimony probably consists
of molecules of Sbi, which break down into simpler molecules as the
temperature rises.

Compound of Antimony with Hydrogen — Stibino, SbHs. — Anti-
mony forms with hydrogen the compound SbHj, which is analogous
to the compounds of hydrogen with nitrogen, phosphorus, and arsenic.

NHg ammonia

PHg ... . . phosphine

AsHg arsine

SbHg stibine

261

PRTNCrPLES OF INORGANIC CHEMISTRY

Of these sabstances only ammonia has pronounced basic prop-
erties, but, as we have seen, this is not a very atroug base, Stibine
ia formed in a manner strictly analogous to arsine, by the reduc-
tion of antimony com pounds by nascent hydrogen. When a solu-
tion of an antinionj compound in liydrochloric acid ia introduced
into a flaak iu whieh hydrogen is being generated, the antimony
ia reduced to stibine. Stibine like arsine is unstable at an ele-
vated temperature. When heated to 1B0° it breaks down into anti-
mony and hydrogen. When ignited it burns with a characteristic
pale, greenish flame, which de|>osits a mirror of antimony on a
cold object introduced into it. These facts are utilized for detecting
antimony by the method of Marsh, The procedure for detecting
antimony is exactly analogotis to that employed for detecting arsemc,
and the results are very similar if antimony or if arsenic is present.
If either is present the metaMike mirror appears when the tuba is
heated, and the dark ai>ot forms on the cold porcelain when it is held
in the flame, Tlie question arises^ How can we tell whether we are
dealing with arsenic or with antimony, or with both? There are
certain differences between the two dei>08it8 which enable ua to dis-
tinguish the one from the other. The deposit of arsenic is very vol-
atile, reatlily moving ahmg the tube when the flame is placed beneath
it Antimony, on the other hand, is much less volatile, forming
little globules when heated. The arsemc mirror is soluble in so-
dium hypochlorite, while the antimony is not When the mirror ia
treated with hydrogen sulphide, if it is arsenic it is converted into
the yellow sulphide of arsenic; if antimony, into the red sulphide
of antimony,

COMPOUNDS OF ANTIMONY WITH OXYGES? AND HYDROGEH

With oxygen alone antimony forms three compounds : antimony
trioxide — SbjjOg, antimony tetroxide — BbaO*, and antimony pentox-
ide-*Sb,0,.

Oxides of Antimony. ^ — The oxide of antimony containing the
least amOLHit uf oxygen is the tnoxkie or sesquioxide, Sb^O,. It
occurs in nature as ^^itarmontite^ and is readily prepared by oxi-
disising antimony either with some strong oxidizing agent such as
nitric acid, or by burning antimony in the air. In the latter ease
there is some of the higher oxide, SbO^ formed along with
trioxide.

When antimony trioxide is treated with strong acids it
salts, and^ therefore^ has basic properties. When these salts

h the ^H

forms ^H
ts are ^H

ANTIMONY 268

treated with a large volume of water they decompose readily, yield-
ing the free acid and the compound Sb(0H)8. This hydroxide loses
water easily and passes over into the trioxide : —

2 Sb(0H)8 = 3 H2O 4- SbjO,.

It may, however, lose only one molecule of water and form meta-
antimonious acid : —

Sb(OH)3=H,0-hHSbO^

Salts of this acid with certain metals are known, having the composi-
tion MSbO^

The compound HSbOj, which can also be regarded as SbO.OH,
can form salts with strong acids. The group SbO seems to act as
a unit, and is known as the autimonyl group. The best-known
antimonyl compound is the double tartrate of potassium and anti-
mony. This is known as tartar emetic. It has the composition
SbOK.C4HA.

Potassium antimonyl tartrate is an antimony compound which is
readily soluble in water.

Antimony trioxide is a yellow powder which boils at 1560°. At
this temperature the vapor-density has been determined and corre-
sponds to the formula Sb4O0.

The tetroxide of aiUimony, Sh^O^ is formed, as we have seen, by
burning antimony in the air, especially at a high temperature. It is
also formed by highly heating the trioxide in the presence of air.
Like the trioxide, the tetroxide has both acid and basic properties,
depending upon the conditions. Towards strong acids it acts like a
weak base, while towards strong bases it acts, when combined with
water, like an acid. We have salts corresponding to the general
formula MaSbjOs. These are obviously salts of the acid HaSbj^Oj,
which is formed by the union of the compound Sb204 with a molecule
of water : —

SbA+H20 = HjSbA.

It is interesting to ask how a compound can be both acidic and
basic, depending upon the conditions. How does such a compound
dissociate ? In the presence of a strong acid, where there are many
hydrogen ions in the solution, the compound dissociates yielding
hydroxyl ions, which combine with the hydrogen ions of the acid,
forming water. If the compound is brought into the presence of a
strong base where there are many hydroxyl ions, it dissociates yield-
ing hydrogen ions, which combine with the hydroxyl ions forming
water. We can thus see how the same compound can be acidic or

264

PRINCIPLES OF INOKGAOTC CHEMISTRY

basic^ depending upon the conditions. We shall meet other and
better examples of this same kind of action.

Antimony pentoxkle^ Hh/)if is obtained either by heating anti-
mony with a strong oxidizing agent like nitric aeid^ or by carefully
removing the water by heat from antimonic acid. The compound,
we shall learn, has the formula Hj^SbO^ and when heated, —

2 H^SbOi = 3 H,0 + Sbpfl.

When antimony pentoxide ia heated it loses oxygen and passes over
into the tetroxide.

Acids of AiLtlmony. — Antimony combines with hydrogen and
oxygenj forming several compounds which are acids, but these are
not so numerous as in the cases of phosphorus and arsenic. Indeed,
the compound Sb(0H)3has distinctly basic properties towards strong
acidsj as we have seen, and metaanifmomouif acid itself, SbOOH, may
act basic as in potassium antimonyl tartrate. Although the basic-
forming property begins to manifest itself in antimony, yet it forms
certain well-defined acids. The best known of these is antimonic
acidf having the composition HaSb04, This is formed by the action
of strong oxidizing agents such as nitric acid or aqua regia, on anti-
many. It is also formed when the pentachloride of antimony is
treated with water, which is analogous to the formation of phos-
phoric acid from phosphorus pentachloride : —

SbC!, + 4 H,0 = 5 HCl + H,SbO,.

When antimonio acid is heated to 17a% it loses water, passing first
into metaaniinionic acid; —

H,Sb04 = H,0 + HSbO^

Salts of pffroantimonic acid have lieen prepared. We can regard
this acid as being formed from antimonic acid, in the same way as
pyrophosphoric acid is formed from phosphoric acid by loss of

The sodium salt of this acid has the composition, NaaHjSbjO^. A
potassium salt is known having the composition, K^SbjO-, but this
readily breaks down, in the presence of water» into potassium hydroxide
and the compound KiH^SbjO^. The acid H^S^ib^O^ must, therefore, dis-
sociate into H, H, H^b,Oj, and it is only under extreme circum-
stances that more than two hydrogen atoms separate as ions.

Componndft of Antimony with the Halogeni. — Antimony forms
two compounds with chlorine — the trichloride and the pentar

ANTIMONY 265

chloride. Antimony trichloridey SbClj, is formed by the action of
a mixture of hydrochloric and nitric acids on antimony ; also by the
action of chlorine on an excess of antimony at an elevated tempera-
ture. It is a soft solid, which, on account of its consistency, is
known as antimony butter.

When antimony trichloride is treated with water, oxychlorides
are formed, which are, however, decomposed by an excess of boiling
water, losing all their chlorine and passing over into antimony
trioxide.

Antimony pentachloride, SbClj, is formed by the action of an
excess of chlorine on antimony at an elevated temperature. The
antimony bums in the chlorine, with an evolution of light and heat.
It is also formed by the action of chlorine on antimony trichloride.
Antimony pentachloride is a liquid at ordinary temperatures, boiling
at 140** and freezing at — 6°. At its boiling-point the vapor is only
slightly decomposed, thus differing from the pentachloride of phos-
phorus. When antimony chloride is treated with small amounts of
water it combines with the water, forming definite hydrates ; when
boiled with an excess of water it decomposes, forming antimonic
acid. Antimony combines with bromine, forming the tribromidey
SbBr,; with iodine, forming a tr iodide and possibly a pentaiodide,
and with fluorine, forming a tri- and a pentaflnoride.

Compounds of Antimony with Sulphur. — Antimony forms two
compounds with sulphur — the trisulphide and the peutasulphide.
Antimx>ny trisulphide^ Sb^Sj, occurs in nature as antimonj/ blende. It
is formed by the action of hydrogen sulphide on a solution of an
antimony salt, in which the antimony is in the trivalent condition.
When hydrogen sulphide is conducted into a solution of antimony
trichloride, in the presence of a little hydrochloric acid, the following
reaction takes place : —

2 SbCls -t- 3 H^S = SbA -f 6 HCl.

Antimony trisulphide thus prepared has a dark-red color, with
a slightly brownish tint. It is soluble in concentrated hydrochloric
acid, and also in solutions of alkaline sulphides, forming salts of
sulpho-acids of antimony, which will be considered a little later.

Antimony pentasulphides, SbjSs, is formed by the action of hydro-
gen sulphide on an antimony salt, in which the antimony is pen-
tavalent. When antimony peutasulphide in solution in the presence
of tartaric acid is treated with hydrogen sulphide, the following
reaction takes place : —

2 SbCl, -t- 5 H,S =r 10 HCl + Sb^^

266

PRINCIPLES OF INORC

It is also formed when hydrogen sulphide is passed into an actdi^
fied solution of antimonic acid ^ —

2 HaSbO* + 5 H^ = 8 H,0 + S W%

Antimony pentasulphide is an orange-red powder, not soluble in
dilute, but dissolves in concentrated acids* It dissolves when
treated with the sulphides or polys ulphides of the alkalies, form id g
sulpho-salts of antimony^ which will now be considered-
Compounds of Antimony with Bulpbnr and the Metak. — We have
seen that arsenic combines with sulphur and the metals^ forming
salts of sulpho-aeids of arsenio. In an analogous manner, antimony
forma salts of suIpho-acids. When antimony tri sulphide is treated
with the sulphide of an alkali metal, such as potassium sulphide,
ammonium sulphide, or poly sulphide, the antimony trisi^dphide di^
solves, forming a salt of a sulpho-acid, MSbS^t which is a metaaulph-
anHmonite. Salts of sulphantimonions acid, HaSbSa, are also known.
When antimony pentasulphide is dissolved lu an alkaline sul-
phide, salts of snlphantimonic acid are formed; —

Sb^, + 3 NhS = 2 Na^SbS^.

This compoundj which contains nine molecules of water, is known
as Schlippe's salL It is also formed by the action of caustic soda on
antimony trisulphjde and sulphur.

These sulphur acids are the strict analogues of the oxygen acids,
containing sulj>hnr in the place of oxygen. When the sodium salt
of sulpliantimonic acid dissociates, the ions are : —

NajSbS, =^ Na, Na, Na, SbS^

s

The ion 8bS< cannot be regarded as very stable, since when the above
salt is treated with an acid, which is the same as to add a large
number of hydrogen ions, the acid HaSbS^ is not formed, but this
breaks down into hydrogen sulphide and antimony pentasulphide.

Hard Lead. — \V)ien antimony is fused with lead, the alloy is
much harder than lead, and is known as hard lead. Another alloy
of antimony and lead is known as type-metal*

y

CHAPTER XX

BISBffUTH (At. ^W^t. = 20a5)

The last member of the nitrogen, phosphorus, arsenic, antimony
family of group V in the Periodic System is bismuth. We have
seen that as the atomic weight increases, the elements become less
acidic, and the basic properties begin to manifest themselves. This
condition, which has already appeared in antimony, is intensified in
the element which we are now about to study — bismuth.

Occnrrence and Properties. — Bismuth occurs mainly in the free
condition, but also combined with sulphur as the trisulphide, hifi^
Bismuth is obtained from the sulphide by burning out the sulphur
with oxygen, when it is transformed into the oxide. The oxide is
then reduced by carbon, yielding the element.

Bismuth is a crystallized solid. It forms crystals which are
isomorphous with arsenic and antimony, t.e. the crystals have the
same form, and the two substances can crystallize together. Bis-
muth combines directly with oxygen at an elevated temperature,
forming the trioxide. Like the other members of this group, it
combines directly with the halogens. Bismuth melts at 264®, and
it boils in an atmosphere of hydrogen at about 1600^

The metallic nature of bismuth begins to manifest itself in its
behavior towards the electric current. It shows very marked conduc-
tivity.

Some of the most important substances containing bismuth are
certain of its alloys with other metals. These have the remarkable
property that they fuse far below the melting-point of the lowest-
melting constituent. The well-known Rosens fusible metal consists
of one part of lead, one part of tin, and two parts of bismuth. It
fuses at 93®.8. Another alloy of the same metal, consisting of five
parts of lead, three parts of tin, and eight parts of bismuth, fuses at
94^5. There is an alloy of bismuth still more remarkable than the
above, in that it fuses at 60®.5, and is known as Woo(Vs metal.
This contains two parts of lead, one part of tin, four parts of bis-
muth, and one part of cadmium. It is the lowest-melting alloy of
these substances, and is, therefore, known as their etUectie aUoy —

267

268

PRINCIPLES OF INORGAKIC CHEMISTRY

a eutectic alloj of any two or more metals being the lowest-melting
alloy of those substances. These alloys are used in scientific investi-
gations where a low-melting metal is needed. When Wood's metal is
heated in a test-tube with water it melts long before the water boils.

Compounds of Biamuth with Oxygen and Hydrogen. — ^Miile bta^
muth forms four compounds with oxygeoj bismuth oxide BiO,
bismuth sesquioxide hUOsi bismuth dioxide BiOj, and bismuth
pentoxide BijOj, the only compound of imi>ortance is the sesqui-
oxide. It is formed when bismuth burns in the ain It combines
readily with acids, forming water and the corresponding salt, and is,
therefore, a base.

The corresponding hydroxide, Bi(OH)a, has decidedly basic prop-
erties. It combines with acids forming salts of the general type
BiR|, where R is the anion of a monobasic acid. Thus, with nitric
acid bismuth hydroxide forms the salt Bi(N03)5. The compound
Bi(0H)3 is, therefore, a triacid base, dissociating as follows ; —

Bi(OH)s='BV OH, OH, OH.

The compound BiO.OH derived from BiCOH)^ by loss of water ^ —
Bi(OH)a= HsO -h BiO.OH— ia also basic. Thus, with nitric acid
this buse forms the compound BiO.NO,, and possibly also the com-
pound Bi(NOj)B. The group BiO is known as Usmuihyl^ and its
salts as bismuthyl salts, or basic bismuth salts.

When bismuth nitrate is treated with water it passes over into
a basic nitrata or subnitrate of bismuth. Bismuth hydroxide has
slightly acid properties when in the presence of a strong hise like
potassium hydroxide. The compound formed is, however, very
unstable.

Bismuth Chloride, BiCla^ — That bismuth can form a trichloride
we would expect from the triacid nature of its hydroxide. It is,
however, not formed by treating bismuth hydroxide with aqueous
hydrochloric acid, since the water present would decompose it and
give a basic salt.

It is formed by the action of chlorine on bismuth and has the
composition BiCla. When this is treated with water it passes over
into the oxychloride BiOCl, which is really the chloride of the group
bismuthyl — BiO.

Bismuth Sulphide, Bl^S,, — This compound occurs in nature as
bismuth bletuk. It is readily made in the laboratory by treating
a solution of a bismuth salt with hydrogen sulphide: —

2 Bi(KO,), + 3 H^ = Bi^, + 6 HNO,,

BISMUTH 269

It is a black substance, and is the form in which bismuth is usually
precipitated in qualitative analysis. The sulphide of bismuth is
practically insoluble in an aqueous solution of an alkaline sulphide^
and is thus separated from arsenic and antimony.

Bismuth sulphide is insoluble in dilute acids, but dissolves in
hot, concentrated^ hydrochloric acid.

mAPTER XXI

VANADIUM, COLUMBITJM, NBODYMinM, PHASEODTMrtTM,

TANTALUM

The remaioing members of this uatural group of elements are
vanadiunij columbium, neodymium, prasetxlymium, and tantalam.
These are all very rare substances atid will, tli^refore, be considered
very brietiy.

Vanadium (At, Wt = 51,2)- — VanadiTiin occurs in nature chiefly
as vanadates. These are salts of the acid HsVO^, Vanadium also
forms a metavanadic acid HVO3* Vanadium combines with oxygen
forming the pentoxide, Yfi^ which has weakly basic properties,
It also forms a trioxide or sesquioxide V^^, Vanadium forms
the chlorides YCl*., VCI3, and VCV It *l3'^> forms the oxy chloride
VOCI3, Vanadium combines directly with nitrogen at an elevated
temperature, forming the compound VK. Indeed, vanadium is one
of the few elemetits which burn in nitrogen.

Columbium (At Wt. = 94.0). — This element is frequently known
as niobium. It forms a pentoxide, Cb^Oa (or l^baO^y) which in the
presence of water has w^eakly acid properties. The composition of
the acid is H,,Cb04. It combines with chlorine, forming the penta-
chloride CbCls, which decomposes with w^ater yielding an oxy*
chloride : —

CbCl^ + H3O = 2 HCl -h CbOCla.

Columbium also forma a trichloride, and is thus analogous to
members of the nitrogen gronp, Columbium readily forms a double
fluoride with potassium fluoride, having the composition KjCbF^,
It also forms an oxyfluoride with potassiuni fluoride, having the
composition K/'bOF^j.

Praseodymium and Keodymium (At. Wts, = 140.5 and 143.6). —
These elements, which for a Jong time were regarded as one and
called diilymium, occur in mmarakite, cerite^ vmnmite^ sandf etc.
Until a few years ago they were elasaed among the very rare sub-
stances. In the last few years they have been discovered in consid-
erable quantity in monaxite sand, in connection with the preparation
of the mantles of Welsbaeh lights. Monazite sand has been worked

S70

VANADIUM, NEODYMIUM, PRASEODYMIUM, ETC. 271

over in large quantity by Waldron Shapleigh to obtain pure cerium,
lanthanum, thorium, etc., and during this work much praseodymium
and neodymium have been separated in the form of the double nitrate
with ammonium. More than a thousand tons of material, rich in
these elements, are now in the possession of the Welsbach Light
Company at Gloucester, New Jersey.

The two elements were separated from didymium by Auer Von
Welsbach in 1885, by fractionally crystallizing the double nitrate
with ammonium more than a thousand times.

Praseodymium forms the oxide Pr^Oj. When this is reduced in
a current of hydrogen it passes over into PrjOg. Praseodymium con-
ducts itself in many respects like aluminium, forming the sulphate
Pr2(S04)s, and in general acting as a trivalent ion in forming salts
with strong acids. Its salts are beautifully green in color, whence
the name of the element.

When Von Welsbach separated didymium into its two constitu-
ents, he called the one praseodymium, from the color of its salts, and
the other neodymium, or the new dymium, Neodymium forms the
oxide NdjOj, and the sulphate NdjCSO^)^. Like praseodymium, it
resembles aluminium and the members of the aluminium group in
forming salts in which it plays the role of a trivalent ion. Its
salts, as already stated, are purplish-red in color, and beautifully
crystallized.

Both of these elements form beautifully crystallized double
nitrates with ammonium, having the composition 2(NH4)N08,
PrCNOa), 4 H,0, and 2 (NH4)N0a, NdCNOg), 4 H,0.

With oxalic acid they form oxalates insoluble in dilate nitric
acid, and can thus be separated from all of the more common
elements.

Tantalum (At. Wt. = 183). — Tantalum, so called from the diffi-
culties experienced in isolating it, occurs in nature with columbium,
which it closely resembles in its properties. With oxygen it forms
TajOfi, which in the presence of water is a weak acid. The acid
has the composition H8Ta04. Tantalum combines with chlorine,
forming the pentachloride TaCl^.

CHAPTER XXII

t

CARBON (At, W^t. = 12.0)

We now come to one of the most important elements in the
whole field of chemistry- Thia is the first member of group IV —
ourlxnL Thia element is impot-tant not only as being a great store-
house of intrinsic energ}% winch can readily be converted into heat,
meehauical energy, liglit^ eleetrical energy, etc., bat as being an es-
seutiiU constitnent of every living thing, from the simplest organism
to the most complex. The number of elements which enter into liv-
ing matter is not large, hydrogen, oxygen, nitmgen j sulphur and
phosiihorns in many cases ; but carbon is always present, and is
probably more closely connected with the vital fuDctions than any-
other elcjiient.

AUotropic Forms of Carbon; Bmmond and Graphite. — We know
cui'Ihju in several modifications, both ciystallized and amorphous.
Then* aru two crystalline modifications known resi^eetively as dia-
inoud and graphite. The diamond is carbon and nothing but carbon,
WA is shown by the fact that when the diamond is burned in oxygen,
it is couvertetl completely into a eomix)und of carbon; and when
this numpoinul is collected and weighed, the amount of carbon pres-
mit iu the compound is exactly *equal to the weight of the origiiial
diamond.

The diauiond occurs chiefly in Brazil^ India, and South Africa,
UNually in a mica schist called HacohtmUe. To be of value as a gem
it must be out, as it is said, Le. artificial faces must be ground upon
it, HO an to obtain the highest brilJianuy, The cutting of diamonds
in nuite an art, especially as carried on in Amsterdam. The dia-
iiumd is tlu* hardest of all known substances, with the possible
ricirpiion of hmnu In order to cut it some of its own dust must be
Us<hI, and Un- lliis purpose the smaller and poorer diamonds are |>ow-
dered. IHauiondg as they occur in nature are usually white, but
bhu)k OUMM ftio frequently found.

IManiunds have now l>een made artificially^ — a problem which
has attratitin! great attention in time past. The French chemist
liloiHsau was the first to solve this problem as far as small diamonds

273

CARBON 273

are concerned. In 1893 he prepared diamonds in connection with his
beautiful investigations at very high temperatures, obtained by means
of the electric furnace. His electric furnace is extremely simple,
consisting essentially of two electrodes of carbon, terminating in the
interior of two blocks of lime which fit tightly, forming the crucible
in which substances are heated. Temperatures as high as 2500** can
readily be produced and used, and even 3000** can be secured, but at
this temperature the lime quickly melts.

In such a furnace Moissan saturated molten iron with carbon.
The molten iron was poured at once into a mould which was cooled
by water, and the iron quickly solidified externally. Iron saturated
with carbon expands on cooling, so that as the molten interior solidi-
fies an enormous pressure is produced. Under these conditions
the carbon crystallizes in the form of small diamonds, within the
iron. When the iron is dissolved in an acid, the residual carbona-
ceous matter contains the small diamonds. The largest diamond
which Moissan has thus far prepared has a diameter of only 0.5 mm.
These, however, resemble the natural diamond very closely in their
hardness, their resistance to acids, crystalline form, and even the
striations which occur upon them.

The preparation of large diamonds artificially is as yet an un-
solved problem.

Another crystalline modification of carbon is known as graphite
or plumbago. While the diamond is comparatively rare, graphite
occurs in nature in considerable quantities, especially in Siberia.
Graphite can readily be prepared by heating amorphous carbon such
as ordinary charcoal in an electric furnace, or better by dissolving
carbon in molten metals and allowing it to crystallize.

In order that graphite should combine with oxygen it must be
heated to a very high temperature. Graphite, unlike amorphous
carbon, is a very good conductor of the electric current, and like all
other forms of carbon is very resistant to the action of reagents in
general. Graphite is extensively used in making lead pencils.

All graphites do not seem to be the same. Indeed, there seems
to be a large number of graphites which differ slightly from one
another in properties.

Amorphons Forms of Carbon. — Carbon occurs in the uncrystal-
lized condition in many forms. One of the best known is charcoaly
or wood charcoal as it is usually termed. If a piece of wood is
heated to a high temperature in the presence of an abundance of
oxygen, the carbon unites with the oxygen, forming the well-known
compound, carbon dioxide. If, on the other hand, wood is heated

274

FEINCIFLES OF mORGANIC CHEMISTRY

without free access of air, many products are formed, but the carbon
remains behind for the moat part as charcoal.

Charcoal is prepared in large quantities by what is known as the
deatructive distillation of wood, which consists in heating wood to an
elevated temperature without free access of air. In the charcoal
pita the wood is placed on end in the form of a large circular pile,
with small spaces betweeu tlie separate pieces. The whole is then
covered with earth to prevent free access of air. When the wood is
burned under these conditions, the carbon does not unite with oxy-
gen, but remains behind in the furra of charcoal.

Another form of amorphous carbon is known as coke. This is
obtained by the destructive distillation of ordinary coal without free
access of air. These conditions are realized when eoal ia heated for
the purpose of manufacturing illuminating' gas, and in addition
there are large plants in various parts of the world for preparing
coke. In these coking-ovens coal is subjected to destructive distil-
lation without free access of air, and coke is formed.

Bone-blavk is another form of amorphous carbon obtained by the
destructive distillation of honea> To obtain it in pure form the
inorganic matter contained in the bones is dissolved out by some
strong acid. On account of its great power to absorb certain coloring
matters, bone-black is extensively used to remove these substances
from certain solutions, and especially from solutions of cane-sugar.
In the purification of sugar enormous quantities of bone-black are
used annually, the solution of sugar being slowly filtered through
the bone-black. After the bone-black has become saturated with
the coloring matter of the sugar it is heated again and the color-
ing matter destroyed* The bone-black can then be used again
for purifying more sugar, and this process can be frequently
repeated.

Bone-black and also wooil-charcoal have remarkable powers of
absorbing certain gases, especially carbon dioxide and ammonia.
These substances are, therefore, frequently used to remove objec-
tionable gases from water and other sources.

Sooi or lamp-black is an amorphous form of carbon obtained hy
introducing a cold object into the flame of an ordinary lamp. Under
these conditions some of the carbon, before it combines with oxygen,
is deposited in a very finely divided condition known as lamp-black
or soot. Lamp-black has a number of applications. On account of
its very fine division and intensely black color it is freqiiently used
as a coloring matter. It is also used in preparing carbon inks, which
are very resistant to all chemical reagents.

CARBON 275

CocU or stone-coal is the form in which freo carbon occurs most
abundantly in nature. There are great beds of these deposits in
many places on the earth, and these are of fundamental importance
for the welfare of the human race. In coal we find vast quantities
of intrinsic energy which can readily be converted into other forms,
and our steam-engines, electric motors, electric light plants, etc., are
all dependent upon coal for their utility.

These deposits of coal are chiefly of vegetable origin. In certain
localities where there has been a great accumulation of vegetable
matter, this has undergone decomposition without free access of air,
and the carbon has been deposited in the form of coal. Some of
these deposits are much older than others and have been subjected,
due to geological changes, to greater pressure and higher tempera-
ture. We, therefore, have different varieties of coal. If the coal is
hard and comparatively free from volatile oils, it is called anthracite;
if it is soft and contains much volatile matter, it is known as bitu-
minotia coal. If the process of coal formation is not very far ad-
vanced, we have peat, Ugnitey etc.

The Different Forms of Carbon contain Different Amonnti of
Energy. — It is obvious from the above that many forms of carbon
are known. The question arises. How do these forms differ from one
another? They are all carbon, and materially considered nothing
but carbon, and yet have very different properties. We have met
with analogous cases in the different modifications of oxygen, sul-
phur, and phosphorus, and found in every one of these cases that
the different modifications of the same element contained different
amounts of intrinsic energy. We would naturally look for the same
differences in the case of carbon.

Light has been thrown on this question in the case of carbon by
the experimental work of Favre and Silbermann. They measured
the heats of combustion of the different modifications of carbon and
obtained the following results : —

Heat or Combustion

Charcoal
Retort carbon

Diamond

Graphite

06.080 calories

06.630 calories

r 04.550 calories
\ 03.240 calories

03.360 calories

276 PRINCIPLES OF INORGANIC CHEMISTRY

Since the end product is the same in every case — carbon dioxide
— the differences between the heats of combustion of the various
forms of carbon are a measure of the different amounts of intrinsic
energy in these different forms.

We see that these differences are quite considerable, amorphous
carbon having the largest amount of intrinsic energy and the crystal-
lized varieties the least. The same general results which were
obtained with the allotropic modifications of the other elements, also
appear in the case of carbon.

Phyneal Properties of Carbon. — Carbon, except in the form of
the diamond, is a black solid, hard, and having a more or less metal-
lic lustre in graphite and anthracite, soft in wood charcoal and coke,
and a fine powder in soot or lamp-black.

Carbon remains solid until an enormously high temperature is
reached. In the electric arc, where the temperature is probably in
the neighborhood of 3500**, carbon vaporizes, but even at this enor-
mously high temperature, comparatively slowly.

The specific heat of carbon is anomalous, depending upon the
temperature. According to the law of Dulong and Petit, the spe-
cific heat of an element multiplied by its atomic weight is a con-
Btanti 6.2. If the specific heat of carbon is taken at ordinary
temperatures, this relation does not hold. It has been found, how-
ever, that the specific heat of carbon increases with rise in tempera-
ture, becoming practically constant at about GOO"*. This will be seen
from the following results : —

TSMPUUTURB

SpBCirio Heat op Cabbok

-1(^.6

0.096

68^.3

0.163

140^0

0.222

24r.O

O.JM)3

606^9

0.441

806O.0

0.449

Be heat, which is obtained about 600^, is
. ^^ ,^ atOBiic weight of carbon, the constant is nearly ob-

d Diikng and Petit, then, holds as well for carbon
Branded the specific heat of carbon is
it lias a maximum, constant value.

CARBON 277

COMPOUNDS OF CARBON

Carbon combines with hydrogen, oxygen, nitrogen, and sulphur,
forming such a large number of compounds that a separate branch of
chemistry has grown up around the element carbon. This branch,
which is one of the largest of all the branches of chemistry, is known
as organic chemistry. Indeed, the study of the compounds of carbon
has almost absorbed the attention of chemists for the last forty years,
and much of the best chemical work has been done along these lines.

While the study of the compounds of carbon belongs to organic
chemistry, we shall take up a few typical, fundamental substances
to give an idea of the kind of compounds which carbon forms w'ith
other elements.

Componnds of Carbon with Hydrogen. — Carbon forms with hydro-
gen a very large number of compounds. Indeed, these two elements
form several series of compounds, the individual members of any
series differing in composition by one carbon atom and two hydrogen
atoms. The simplest compound of carbon and hydrogen has the
composition CH4, and is knOwn as marsh gas, or methane. There is
a whole series of compounds closely related to methane, and known
as the methane series. The simpler members are —

Methane CH4

Ethane 0,1^

Propane CsHg

Butane C4H,o

Pentane CjHu

Hexane CflH^

Carbon forms with hydrogen a compound containing just twice as
much carbon in proportion to the hydrogen as methane. This com-
pound has the composition C2H4, and is known as ethylene. This,
like methane, is a fundamental substance, and the first member of
a group of hydrocarbons which have a constant difference in composi-
tion by a constant amount. The first few members of this group are —

Ethylene C2H4

Propylene . . : CjHe

Butylene C^Hg

Such series of compounds as the above, in which successive mem-
bers differ in composition by the group CHj, are known as homologous
series of compounds. The above series, of which several members
are known, is the ethylene series of hydrocarbons.

278

PRTNCTPLES OF INORGAmC CHEMISTRY

Carbon forms with hydrogen another series of compomads con-
taining still more carbon with respect to hydrogen. The first mem-
ber of this series is known as aceiyknej and has the composition
CaHa* A few members are given i —

Acetylene CaHj

Allylene ........ C^H^

Carbon forms with hydrogen still another series of compounds,
containing even more carbon in proportion to hydrogen than acety-
lene. The first member of this series of compounds is known as
benzene^ and has the composition CaHg, A few of the succeeding
members are —

Benzeue CflH^

Toluene ,.*,.... C^Hg
Xylene ........ C^Hia

The members of this series of compounds differ fundamentally in
their properties from those of the three series already considered.
This series is, therefore, not to be regarded as an extension of the
other three series in the direction of more carbon and less hydro-
gen. It would lead too far to discfiss the nature of this difference.

Componnda of Carbon with Oxygen. — Carbon forms three com-
pounds witli oxygen ; carbon mDooxide, CO, containing in the mole-
cule one atom of carbon and one of oxygen j carbon dioxide, C0»
and carbon suboxide, C3O3,.

Carbon Honozide, CO. —Carbon monoxide is formed by the direct
union of the two elements. When carbon is heated in a limited
supply of oxygen J the product is carbon monoxide. It is also formed
by the action of highly heated carbon on carbon dioxide, —

COj + C = 2 CO,

by the action of highly heated carbon on water-vapor, —

H,0 -h C = H3 -h CO,

and in many other reactions* The most convenient method, however,
of preparing carbon monoxide is by heating formic acid, a compound
having the composition H^COj with mlphuric acid* This decom-
poses as follows: —

HjCO, = H,0 + CO^

Carbon monoxide is a colorless, poisonous gas. When breathed
into the system it combines with the hBerooglobin of the blood and
prevents the latter from carrying out its normal function^ which is

CARBON

279

to carry oxygen to the Tarioua organs of the body. Its presence
in the blood can be detected by certain characteristic bands in the
absorption spectrum*

Carbon monoxide Is frequpntly formed in the incomplete com-
bustion of CAfbon in poorly ventilated fun* aces. From such furnaces
it easily escapes into the room^ and is breathed by the inhabitants.
If the room into which carbon monoxide is escaping is poorly ven-
tilated, bad results may follow,

Oartx>n monoxide combines directly with oxygen, forming carbon
dioxide, and is, therefore, a good reducing agent. When carbon
monoxide is brought in contact with the hot oxides of the heavy
metals, they are reduced to the me tall ic condition, and the carbon
monoxide is oxidized to carbon dioxide.

When carbon monoxide is brought in contact with chlorine, and
the mixture exposed to sunlight, the two combine and form the com-
pound COClj, which is known m phoitgene.

Carbon monoxide has the power of combining directly with
certain metals, forming remarkable eomi>ound8. With finely divided
nickel heated to 100* carbon monoxide combines forming tikkd
tetracarbomjl:- jfi + 4 CO = Ni(COV

Willi iron it forma the pentaearbanifl, Fe(CO)|.

Carbon monoxide was one of the substances which remained nn-
liquefied for a long time. This was on account of its low critical
temperatare, — 139*.5. The critical pressnre is only 35.5 atmos-
pheres. Carbon monoxide has a boiling-point of — 190^ which is
very low. At a little lower tem|>prature, —211°, it solidifies.

ThennochemiBtry of Carbon Monoxide. —When carbon is burned
to carl>on monoxide, the amount of heat set free is only about 2000
calories, while WOO calories are liberated when carbon monoxide is
oxidized to carbon dioxide. Carbon monoxide therefore contains a
large amount of intntisic energy which can be converted into heat
by simply oxidising it to carbon dioxide. It is due to this fact that
carbou monoxide is an excellent heating agent^ and, further, is an
important constituent of illuminating gas-

Water-gat. — It has already been mentioned that one of the
methofis for preparing carbon monoxide is to pass wat^r-yapor over
highly heated carbon, the reaction which takes place being —

C + H,0 = CO + K^

This mixture of carbon moooxidB and hydrogen would haire very
little value as an illuminating gas, since both of these gases bum

280

FRrXCIPLES OF INORGANIC CHEMISTRY

with a comparatirely colorless fi^mB, although thej evolve an enor-
mond amount of heat This mixture of gaaes is passed tbrough
highly heated petroleum-vapor, and is thus mixed with hydrocarbons
and other siibstauces vhieh give off an abundance of Tight when
they are bunied. Tins miKture^ known as ivaier^s, is now used
largely for illuminating purposes.

In preparing this gas, ooal ia heated to a very high temperature
in the presence of the air. Water- vapor is then forced over the
highly heated carlion, when the decomposition l^akes place in the sense
of the above equation. When the coal ha^ become cooled to a tem-
perature which is too low, it is again heated in contact with the mv,
steam again passed in, and the process thus continued until the coal
has been used \ip.

Water-gas is now extensively used where illuminating gas, made
by the dry distillation of coal, was formerly employed.

Carbon Biaxidet COj. — The highest product of the direct oxidsr
ticn of carbon ia carbon dioxide* This conipound occurs in a nnniber
of places in the free condition.^ It is one of the constituents, as will
be remembered, of atmospheric air. It also occurs dissolved, in
greater or less quantity, in water. Carbon dioxide escapes from the
earth, in certain localities, either in the free condition or dissolved in
water The famous ** dog's grotto,'* of Naples, is an example. When
a dog or small animal enters this grotto it quickly experiences suffoca-
tion, while a man is not seriously inconvenienced* This is due to
the fact that carbon dioxide is heavier than air and settles to the
bottom of the grotto. It is, therefore, felt raore seriously by the
smaller animals than by man.

Carbon dioxide escapes from certain i^ineral springs, dissolved
in the water of such springs. It is often present in such large quan-
tity as to cause the water to be under uonsiderable pressure. When
the water reajches the snrface of the earth, a part of the gas escapes
and gives the characteristic effervescence.

Carbon dioxide is given off by animals, as can be readily shown
by breathing for a short time Into lime water, or a solution of barium
hydroxide, when insoluble calcium or barium carbonate is formed.

Carbon dioxide is also set free when animal and vegetable matter
decom poses, and also in many processes of fermentation,

Preparaticn of Carbon Bloxide. — Carbon dioxide can be pre-
pared by a number of methotls. Theoretically, one of the simplest
methods is to burn carbon in an excess of air : —

C + 0, = COt.

CARBON 281

Practically, a far more convenient method of preparing carbon
dioxide is to treat a carbonate with an acid. Carbon dioxide in the
presence of water and a strong alkali, forms two series of salts
which have the general composition MHCO^, and MjCOg. When
these salts are treated with an acid, we suppose that the compound
HjCO, is set free. This compound, however, is unstable, and breaks
down at once into water and carbon dioxide, which is liberated.

Carbonates are decomposed, yielding carbon dioxide, not only, by
strong acids, but even by very weak acids, such as acetic. This is
due to the fact that carbon dioxide is a gas, — is volatile, — and it is
a general law of chemistry, that whenever a volatile compound can
be formed, it is formed.

The reaction between acids and carbonates, both acid and neutral,
would be represented thus : —

KjCOa + 2 HCl = 2 KCl + H^O + COj,
KHCOa + HCl = KCl + H,0 + CO^

When certain compounds are heated they readily lose carbon
dioxide, while others lose it only at high temperatures. The carbon-
ate of calcium, or ordinary limestone, or marble, belongs to the former
class. When this substance is heated it breaks down thus : —

CaC03 = CaO + C02-

Chemical Properties of Carbon Dioxide. — The most characteristic
chemical property of carbon dioxide is its power to form salts in the
presence of aqueous solutions of alkalies. When one equivalent of
caustic potash, in water, is brought into the presence of one equiv-
alent of carbon dioxide, the following reaction takes place : —

K0H4-C02 = KHC03.

When two equivalents of caustic potash are used, we have the
following reaction : —

2 KOH 4- CO, = KjCO, 4- HjO.

The first salt is acid potassium carbonate, the second normal potas-
sium carbonate.

Carbon dioxide in the presence of water acts, then, as a dibasic
acid. It dissolves readily in water, the amount dissolved depending
upon the pressure to which the gas is subjected. The aqueous solu-
tion reacts slightly acid, showing that there are a small number of
hydrogen ions present: —

H,C03 = H, HCOs.

282

PEINXIPLES OF INORGANIC CHEMISTRY

The acid is so weak, and Its aqueous solutiotis hare such slight

conductivity, that we are not justified in assuming that there is anjri

dissociation of the ion, HCO3,, in the presence of water alone. When '

an alkali is present, and all the hydrogen ions from thejirst stage of

dissociation are used up, it is probable that the ion HCO^ begins to

dissociate thus ; — _ +

HC03=H, COa,

and this dissociation continues to the end if there are enough hy-
droxy I ions from the base present to combine with all the hydrogen
ions as rapidly as tliey are formed*

The carbonates, like the salts of all weak acids, are hydrolyzed by
water. This is shown by the fact that a salt like sodium carbonate
shows a strongly alkaline reaction, which means that there are by-
droxyl ions present : —

Na4C0»+ H,0=Na, OH + Ka, HCO,.

The hydrolysis of carbonates is by no means complete, only a
comparatively small number of molecules being broken down by
the water as represented in the above equation.

Carbon dioxide is a very stable compound, holding its oxygen
firmly. It can, however, be made to give it up under certain condi-
tions. CertJtin metals, such as zinc, at a very high temperature can
remove half of the oxygen from carbon dioxide, converting it into
carbon monoxide.

REDUCTION OF CARBON DIOXIDE BY PLANTS

Carbon dioxide is being continually reduced hy the green plants
in the sunlight They build up the carbon into complex compounds
with hydrogen, oxygen, and perhaps nitrogen, and these compounds
contain enormous amounts of intrinsic energy. The carlx>n dioxide
obtained by plants comes largely from animals which give itoif whemi
they breatlie, as we have seen. Plants give off oxygen, which is just
what is needed by the animal world.

The complex compounds of carbon ai-e consumed by animals
which decompose these substances into much simpler ones, especially
into carbon dioxide and urea, a compound having the composition
CONjH,. The large excess of intrinsic energy in the complex com-
pounds over that in the sim]iler products which animals excrete, ift|
converted into heat and hy the animal into mechanical work.

The chief source of the energy which animals expend is, then,
the complex compounds of carbon* wliich are built up by plants froB

CARBON 283

the simpler substance carbon dioxide, and which are broken down in
the animal body into simpler substances which contain much less
intrinsic energy.

It is of interest to note that most of the carbon in animal and
vegetable tissues ultimately passes off when these decay, in the form
of carbon dioxide. The carbon dioxide is again taken up by the
plant, converted into complex compounds, consumed by the animal,
broken down into simpler substances, and the cycle is thus completed.

Physical Properties of Carbon Dioxide. — The gas carbon dioxide
can be readily liquefied, since its critical temperature is as high as
31°. Its critical pressure is 73 atmospheres. At lower temperatures
it is liquefied at much lower pressures. At 10® the pressure required
to liquefy carbon dioxide is only about 27 atmospheres, while this is
reduced to 18 atmospheres at — 30^ Carbon dioxide is liquefied on a
large scale by pumping it into thick-walled, steel cylinders, which are
kept cool. Such cylinders are kept in the laboratory as sources of
supply of carbon dioxide.

When the carbon dioxide is allowed to escape from such cylin-
ders through a fine opening, part of it volatilizes and escapes as gas,
while the remainder is converted into the solid condition and can be
caught in a bag placed over the jet. Solid carbon dioxide is a com-
pact, white mass resembling compressed snow. It has been exten-
sively used as a refrigerating agent. When solid carbon dioxide is
mixed with ether it vaporizes rapidly and a low temperature is pro-
duced. When this mixture is allowed to evaporate on the air a
temperature of — 80° results, and when vaporized under diminished
pressure temperatures as low as —100° to —110° can be readily
produced.

The liquefaction of carbon dioxide is of interest and importance
in connection with the liquefaction of gases in general. It was first
converted into a liquid in 1834 by Thilorier, who demonstrated the
refrigerating power of the mixture of solid carbon dioxide and ether.
Such a mixture bears the name of its discoverer and is known as
TliiloHeT^s mixture.

Discovery of Critical Temperature and Pressure. — Cagnairdde la
Tour observed in 1822 that ether and alcohol pass completely into
vapor in a very small space, when the temperature is above a certain
point. Also, that two volumes of ether volatilize at the same temper-
ature as one volume into the same space. This made it probable
that there was a temperature above which these liquids could not
remain in the liquid state, but would pass over into vapor regardless
of the pressure. This observation made but little impression, until

284 PRINCIPLES OF INORGANIC CHEMISTRY

Andwws showeii much later (1869) that there is a temperature for
eveiT gas, al>ove which it cannot be liquefied. This temperature
wiis oadled bv Andrews the critical tempei-atxire of the gas. The
work of Andrews was done largely with carbon dioxide. When the
tiibe <\wt*ining this gas was brought to a temperature of 13^.1, and
}3ttt gas subjooted to a pressure of 48.9 atmospheres, a liquid began
iy^ ap|Wir» and the volume of the gas continued to diminish without
*i\v <\Hwad<*rable increase in the pressure being required. At 21^5 sim-
iW w^«U8 wore obtained. At somewhat higher temperatures, how-
i^wr v^lM and 32**.5), results of a very different character manifested
Ili^M^Wt^ Although there was a marked decrease in volume at a
V^^Ham definite pressure, yet no liquid separated. There was no
i^t^(^u>> that any liquid had been formed. At still higher temper-
^u^Nt th^ abruptness of change in volume at any definite pressure
VtKtuut^ lwi8 and less, and entirely disappeared at 48^.1. These
^^ulla art> stvn conveniently by plotting them in curves; the ab-
^'^(juita» bein^t volumes, the ordinates pressures.

*rh^ o\irvt^ for 13M shows that when a pressure of nearly 50
^liiK^l^^vt^ is reached, the volume diminishes very greatly without
ai<v iMarktn) increase in pressure. This means that the gas has
ytiMjiiKH) ovt^r into liquid at this pressure. The curve for 21*^.1 is
^MiUUlar lo tho alH>ve curve. An abrupt transition from gas to liquid
latlv^a |4a^HS but at a higher pressure. The curves for 31M, 32**.5,
au4 3^"^ v^ Hhow h^ss and less abruptness, but at none of these tem-
y^Mt^aluiv^i U any liquid produced. The curve at 48M shows no break,
Wiug ^H^vf^^tly »nuH)th throughout. The temperature above which
VHmU^ vluv.\uli> oannot be liquefied, was found by Andrews to be
JlO^'Vlii?* auvl thiii i», therefore, the critical temperature of the gas.

'Vhv^ tvu4^)vaturt^ alx)ve which a gas cannot be liquefied has been
M'U40\l l\Y MuuvU^UH^fT the absolute boiling-point of the gas. This is
viU\iv^^> Vh\> muu«« an Andrews's critical temperature.

'4^Kv> V'V\v*MVUH» whioh will just liquefy the gas at the critical tem-
^u'ulkuv KiV4 Uh^u toruuHl the critical pressure. The substance has a
\^\U'Um vU'iiuU^ vb'^iAity under these conditions, and this is its critical
*fc'<*c«*fv. l^^^^ i*iH^i|mHUil of the critical density is the critical volume,

'i%\^ ovitivv^l toiu)H^rature8 and pressures of some well-known
liv|uuU i^»v ij[vv^\ iu tht) following table: —

CARBON

285

CBITIOAIt TSMPKEATUU

CbITICAL PSKSBirBX

Hydrogen ....

— 226°.0

15.0 atmospheres

Nitrogen

— 146°.0

35.0 atmospheres

Carbon monoxide .

— 14r.O

30.0 atmospheres

Argon .

— 120°.0

40.0 atmospheres

Fluorine

— 121°.0

60.6 atmospheres

Oxygen .

— 118^8

50.8 atmospheres

Methane

— 95°.5

50.0 atmospheres

Carbon dioxide

31^0

75.0 atmospheres

Ammonia

130°.0

1 1 5. 0 atmospheres

Chlorine

144°.0

83.0 atmospheres

Bromine

302^.2

The examples given in this table show the great differences in
the critical temperatures of different liquids. It also shows that the
critical pressures of liquids are, in general, not high. If the tem-
perature of the gas is below the critical temi)erature, the pressure
required to liquefy the gas is below the critical pressure. In the
liquefaction of gases, then, low temperature is far more important
than high pressure. Indeed, the temperature must be at least down
to the critical temperature. If the temperature is still lower, very
slight pressure may liquefy the gas. We can now see why the
earlier experimenters were not successful when they tried to liquefy
such gases as oxygen, nitrogen, hydrogen, etc. They used in some
cases enormous pressures, amounting to thousands of atmospheres,
but did not cool the gases down to the critical temperatures. After
these gases were sufficiently cooled they were liquefied at moderate
pressures.

Continnity of Passage from the Liquid to the Oaseous State. — It
will be seen from what has been said in reference to critical tempera-
ture and pressure, that a liquid can be transformed into vapor with-
out becoming heterogeneous at any time. If the liquid is warmed
above its critical temperature, a pressure is produced which is greater
than the critical pressure. The volume may now be increased to
any extent, yet the substance which was originally liquid remains
homogeneous. The passage from the liquid to the gas is thus
perfectly continuous, and it is impossible to say where the liquid
state ends and the gaseous begins. The condition of matter at and
near the critical point has always perplexed men of science, and
many opinions have been expressed concerning it. Andrews dis-
cussed this condition in connection with carbonic acid. He pointed

286

PRINCIPLES OF INOKGANIC CHEMISTRY

r

out that if this gas above the critical temperature is subjected to a
pressure coiisideral>ly above the critical pressure, there is an enor-
mous decrease in volume. The carbon dioxide under this conditiou
is neither gas nor liquid, but occupies a position between the two.

Just as a liquid can be transformed into a gas without aiij break
in continuity, so can a gaa be tmnsformed into a liquid by a continu-
ous process. The gaseous and liquid states^ then, approach as the
critical point 13 reached, and either can be made to pass into the
other without any breach in continuity.

The Kinetic Theory of laquidi. — The close relation which we
ihave just seen to exist between liquids and gases has led to the
.application of the kinetic theory of gases also to liquids. Since the
^passage from a liquid to a gas, and vice vermt under certain condi-
tions is so gradual that we cannot say where the one state of aggre-
gation ends and the other begins, it is highly probable that any
theory w^hieh obtains for the one state irould apply, to some extent
at least, to the other*

The liquid state, aa we have seen^ represents matter in a much
more concentrated condition^ than the gaseous state. There is a
much larger number of molecules in a given volume of a liquid, and^
oonsequeotly, the collisions between the moving molecules are much
more frequent. There would thus result in the liquid an enormous
pressure^ were it not for the attractive forces betw^een the molecules.
These attractive forces hold the moleeules together and prevent them
from flying oBf with explosive violence. Only those molecules which
approach the surface of the liquid with unusually great velocity , can
so far escape from the attractions of the other liquid molecules as to
fly off into the space above the liquid. This explains the existence
of vapor above every liquid. We know, however, that if these mole-
cules fly off into a closed space alK)ve the liquid, the vapor-pressur©
thus produced cannot exceed a certain limit at any given temperar
ture. We can clearly see the reason for this in terms of our theor)^
The molecules of the vapor, in their movements through the oonfin*
ing apace, come in contact with the surface of the liqnid. Some of
these are continually coming within the range of the attractive foi^ces
of the liquid molecules, and are drawn down, as it were, into the
liquid again. There is thus a continual ex<1mnpe going on between
the liquid and the vapor, some liquid particles passing off as vapor,
and some vapor particles condensing aa liquid, until a condition of
equilibrium is reached. Equilibrium is established when the vapor-
pressure has reached such a point that the same number of gaseous
molecules are condensed in any unit of time as there are liquid mole-

CARBON 287

cules converted into vapor. We have seen that it is only the mole-
cules with the greatest kinetic energy which can so far overcome the
molecular attractions as to escape from the liquid as vapor, and this
of course lowers the mean kinetic energy of the liquid. We know
that when a liquid evaporates, the mean kinetic energy of the liquid
molecules decreases, or, as we say, the temperature is lowered. If
the liquid is in such a position that it can absorb heat, it does so ;
and the heat required to effect complete vaporization of a liquid is
very great. This explains why the vapor-tension of a liquid is in-
creased with rise in temperature. The addition of heat increases
the kinetic energy of the liquid molecules, and more are capable of
overcoming the molecular attractions and flying off as vapor in a
given unit of time. The number of molecules in the condition of
vapor is therefore greater, and the vapor-pressure is greater the
higher the temperature.

Carbon Suboxide, C3O2. — This compound has very recently been
prepared by the action of phosphorus pentoxide on ethyl malonate,
CH,(COOCoH5)2 = 2 C2H4 + 2 H,0 4- CjO,, the phosphorus pentoxide
taking up the water. Carbon suboxide is a colorless liquid boiling at
7**. It is unstable, undergoing decomposition at ordinary temperatures.

Compounds of Carbon with Oxygen and Hydrogen. — Thousands
of such compounds are known. While these belong strictly to the
subject of organic chemistry, a few typical substances will be con-
sidered here.

The alcohols are among the simplest of the compounds of carbon
with oxygen and hydrogen. The first member of this series of com-
pounds is methyl alcohol, CH4O, or wood spirit, as it is termed.
The alcohols form a homolc^ous series of compounds which are
analogous to the hydrocarbons. The first members of this series are —

Methyl alcohol . . . CH4O Propyl alcohol . . . CjHgO

Ethyl alcohol . . . C,HeO Butyl alcohol . . . . C4H10O

If we regard methyl alcohol as the first product of the oxidation
of methane, ethyl alcohol is a similar oxidation product of ethane,
propyl alcohol of propane, and so on, the two series running strictly
parallel.

The first step in the oxidation of ethane, CsTIc, is ethyl alcohol,
C^HgO ; the second step in the oxidation is aldehyde^ C2H4O. This is
ethyl aldehyde, the second member of a homologous series of alde-
hydes. The first members of this series are —
Formic aldehyde . HCOH Propyl aldehyde . C2H5COH

Ethyl aldehyde . . CH3COH Butyl aldehyde . . C,HyCOH

288

PRmCIPLES or IKORGANIC CHEMISTRY

Another product of Uie oxidation of a hydrocarbon is an ether.
Take ordinary ethyl ether. This has the composition C4H10OJ and
ia a member of a homologous serves of compoundR. The first mem-
ber is methyl ether, CaHcO, the second methyl-ethyl ether, and so on.

If an aJdehyiie is further oxiilized it passes over into an acidf
and we have homologous series of organic acids, only one of which
will be considered here. If ethyl aldehyde is further oxidized we
have acetic acid, CH/'OOH. This is the second member of the
formic acid series, formic acid being the first; —

Formic acid ,
Acetic acid *
Propionic acid
Butyric acid

HCOOK

CH,UOOH

CjHsCOOH

C,H,COOH

Another homologous series of compounds which carbon forms with
oxygen and hydrogen is known as the ketones. Of these, ordinary
acetone, CH^— CO — CH^j is an excellent example. Still another series
is the Btherml mils or esters^ of which ethyl acetate, CHgOOOCtHa,
is a type j and there are many more such series of compounds, but it
would lead too far to even mention them in this connection.

Compounds of Garbom witlL the Halogens, — Carbon combines
directly with fluorine, forming carbon tetrafluoride, CF*. It does
not combine directly with any of the other halogens, but forms com-
pounds with them by indirect methods. Thus, carboii combines
with chlorine, forming cmifon ietmcUonde^ having the composition
CCI4. This is forjued by the action of chlorine on methane in the
sunlight. The hydrogen of the methane is replaced atom by atom
by chlorine.

(1) CH, ^ C!, == HCl + CH^Cl,

(2) CHaCi + CI, = HCl + CH,C1^

(3) CH,CU -h CI2 ^ HCl + CHCI4,

(4) CHCls 4- CU ^ HCl -h CCI4.

The final product is carbon tetrachloride, the intermediate prod-
ucts being monochlorm ethane, dich I or methane, and trie hi orm ethane
or chloroform.

When carbon tetrachloride is treated with one equivalent of
water, phosgene gas is formed ; —

Ufi + CCl* ^ 2 HCl -h COCl,.

Carbon tetrachloride is a liquid boiling at 77*^^ and solidifying
at - 25°.

CARBON 289

Just as we may have chlorine substitution products of the hydro-
carbons, so we may have bromine, iodine, and fluorine substitution
products, and all are known. The limit in these cases is reached in
the compounds CBr4, CI4, and CF4.

Compound of Carbon with Snlphnr, CSg. — Carbon disulphide, CS,,
is formed by passing the vapors oi sulphur over highly heated carbon.
The two elements unite, forming carbon disulphide, which being
very volatile passes out of the field of action. Carbon disulphide is
easily inflammable, readily uniting with oxygen and forming carbon
dioxide and sulphur dioxide. Carbon disulphide is an excellent solv-
ent, not only for oils, fats, and other complex organic compounds, but
for bromine and iodine and many other substances. It is a liquid
with a highly disagreeable odor, boiling at 46° and solidifying at
— 113**. It refracts light very strongly, having an index of refraction
varying with the wave-length of the light from 1.6 to 1.7.

One further feature in connection with its formation directly from
carbon and sulphur should be pointed out. The reaction in which it
is produced is strongly endothermic, there being considerable heat
absorbed when the two elements unite.

When carbon disulphide is treated with an alkaline sulphide the
two unite : —

M^4-CS2 = N,CS»

forming a salt of trithiocarbonic acid. The acid of which this sub-
stance is a compound has the composition HjCSg, and is obviously
carbonic acid in which' the oxygen is replaced by sulphur. It is
known as trithiocarbonic acid and its salts as trithiocarbonates. The
potassium salt of this acid is used for destroying the louse which is
so injurious to the grape-vine.

Compound of Carbon with Nitrogen — Cyanogen, (Clf),. — When
we consider the inert nature of the element nitrogen, it is surprising
that the compound CN should be capable of existence. The two
elements, however, do not combine directly, but combine readily
with an alkali metal, forming such compounds as potassium cyanide,
KCN. Cyanogen is not so readily obtained from the potassium
compound, but is very easily prepared from mercuric cyanide. This
compound when heated breaks down into mercury and cyanogen: —

Hg(CN),= Hg + (CN)^

Cyanogen is a gas which is characterized by its extremely poisonous
nature. Cyanogen combines directly with potassium at an elevated
temperature, forming potassium cyanide. Cyanogen is quite soluble

^M 290 PEIKCIPLES OF INORGANIC CIIKMISTKY ^^^^|

^P in water. Liquid cyanogeii boils at — 20**7, has a critical tempera- ^M
^m * ture of 124^, and a critical pressure of 62 atmospheres. When cyano ^M
^H gen is formed the action is endothermic. ^M
^M Hydrocyanic Acid, RUN. — Hydrocyanic acid is formed when a ^^H
^H cyanide is treated with an acid : — ^^^|

^^^ MOK + H€l = MCI + HCN. ^^M

^V This acid J which is very soluble in water, is known as prussic aeid- ^H
^H It is characterized by its extremely poisonous nature. Hydrocyanic ^M
^H acid is a very weak a<.'id, as is shown by the small conductivity of ^^^H
^1 its aqueous solutions. A few of these are given below ; — ^^^^H

^^1

1

^^H

^^H

^^M
^^M

^m It dissociates into H, CN, but only to a slight extent When there ^M

^H is a strong alkali present, which is the same as to say an excess of ^H

^V hydroxyl ions, the hydrogeu ions are used up as fast as they are ^M

H ' formed, combining with the hydroxyl ions to form water^ and more ^M

^m of the acid dissociates. I'his progressive dissociation and eombina- ^M

^M tion of the hydrogen ions when formed may c<jntinue until all of the ^M

^m acid is dissociated; and the corresponding cyanogen anions remain ^M

^M in the solution with the cations of the alkali in question. When ^M

^V such a solution is evaporated, t\e. when the water which causes the ^M

cyanide to dissociate is removed, the alkali cations unite with the ^M

cyanogen anions, and a cyanide is formed. Hydrocyanic acid is a ^M

liquid boiling at 2T°, and melting at — 15°. ^M

The cyanogen group shows a very marked tendency to poly- ^M

merize and form comijlex groups. There is a tendency to polymer ^M

ize in groups of three, an4 especially in groups of six. ^M

Cyanic (HOCH) and Snlphocyaiiie (HSCIf > Acids. — TTydrocyanic H

acid can combine with oxygen and form cyanic acid, which has ^M

_ the composition HOCN. This compound shows the tendency of ^|

^H cyanogen to polymerize, since we have also the acids (HOCN)j and ^M

■ (H0CX)3. ■

^m The compound formed by the addition of sulphur to hydrocyanic ^M

^M acid is remarkable in that it is one of the very strongest acids ^M

CARBON 291

known. This is shown by the conductivity of sulphocyanic acid
in water : —

V

Mv

4

337

8

345

16

352

32

858

By comparing the conductivities of sulphocyanic acid with those of
hydrocyanic acid, it will be seen that by introducing a sulphur atom
into the latter, we have passed from one of the very weakest to one
of the strongest acids known. It is generally true, that by increas-
ing the amount of sulphur in the molecule we increase the acidity
of the compound, but not to the same extent as in hydrocyanic and
sulphocyanic acids.

THE r6lE of carbon IN PRODUCING LIGHT

niuinination. — The subject of the production of light or illumi-
nation is one which has attracted attention for a very long time, and
is still doing so. In practically all of the earlier methods of pro-
ducing light, and in many of those used to-day, carbon is employed
in one form or another. Most of the methods of ilhimination owe
their existence to some compound of carbon which is burned or
oxidized, giving out heat and light. This is analogous to what we
saw was taking place in the bodies of animals. The complex com-
pounds of carbon were decomposed into much simpler substances,
which contain far less intrinsic energy than the original substances.
The intrinsic energy which disappeared was converted into heat,
and was the chief source of heat in the animal body.

If this decomposition of complex carbon compounds, or, as we
say, oxidation processes, proceed with sufficient rapidity, there is a
rapid production of heat energy, and light energy results. This is
what takes place in our luminous flames, whether produced by the
candle, oil-lamp, gas-jet, or acetylene light.

Candle and Oil-lamp. — In the candle we have complex com-
pounds of carbon which at ordinary temperatures are solid. These
are made by melting the stearine, tallow, or paraffine, and pouring
the liquid into a mould after a wick has been placed in the centre of
the mould. The object of the wick is to carry by capillarity the

292

PRINCIPLES OF INORGANIC CHEMISTRY

material of the candle, as it is melted, up into the fiame hj capillarity.
The tip of the candle is melted, and the end of the wick ignited*
This melts a portion of the solid hydroearbonss wliich are carried up
hj the wick into the fiame, are %*aporized and burnedj the heat set
free melting more of the solid, and the process is thus a continuous
one. The heat and light are derived from the breaking down of
eomijlex corapounda of carbon into simpler substances, and the
oxidation of the carbon to carbon dioxide*

In our oil-lamps the compounds of carbon which are to be burned
are liquid at ordinary temperatures. These are carried up into the
flame by means of the wick, as in the candle, and the same general
processes are involved in the production of light and heaL

Coal-gflB, Water-^aa. — Cmil-gas is produced, as we have seen, by
the dry distillation of coal, one ton of coal yielding about 10,000
cubic feet of gas. Coal-gas consists largely of compounds of carbon
with hydrogen — hydrocarbons — ^and free hydrogen. These ate the
chief source of the light and heat when coal-gas is burned.

Water-gas is produced by the action of highly heated carbon on
steam, giving, as we have seen, carbon monoxide and hydrogen, and
this mixture is then enriched by abiding to it certain hydrocarbons.
Here, again, the chief source of the light and heat are compounds of
carbon, which are broken down in the flame, and the carbon oxidized
to carbon dioxide.

Flamee and their LuminoBitj. — If we examine
a typical flame, say that of a candle, we observe
three distinct parts : an inner cone a (Fig, 30), of
unburned gases, is surrounded by a zone bj of par-
tially oxidized substances. It is in this zone that
acetylene is formedj which, we shall see, has much
to do with the light-giving power of the flame.
This zone is the chief source of light in the flame.
This is surrounded by a third layer, c, of burning
gasesj and here, where there is an abundant supply
of oxygen from the air, the processes of oxidation
are completed. This part of the flame is relatively
only slightly luminous.

So mucli for the structure of a flame. The ques-
tion remains, What are the causes of the luminosity
of flames ? We have seen that the chief source of
light in a flame is in the middle zone, where the
combustion is not complete* This gave rise to the
theory that the chief source of light in a flame is unburned, solid
particles of carbon, which become heated to incandescence. These

— 5

-a

i

Pio. 30.

CARBOX

293

particles came from the compounds of carbon which are decomposed
by the heat of the flame. This theory accounts for many of the facts
concerning the luminosity of flames, but by no means for all.

Hydrogen gas at atmospheric pressure burns with an almost non-
luminous flame. Hydrogen gas, under high pressure, however, bums
with a luminous flame. The effect of pressure on luminosity is also
shown by burning a candle in a valley and on a high mountain.
Under the former conditions, where the pressure is relatively high,
the luminosity is much greater. Further, gases which bum with a
luminous flame can be made to burn with a relatively non-luminous
flame by mixing them with an indifferent gas, or by simply cooling
them.

These facts cannot by any means be all explained on the solid
particle theory of luminosity. There are undoubtedly many influ-
ences which affect the luminosity of flames, so that probably no
one theory can account for all of the facts connected with this
phenomenon.

According to recent investigations in England, it seems very
probable that the formation and oxidation of acetylene in flames is
vitally connected with their light-giving power.

Bnnsen Burner. — Blowpipe. — Practical use is made of the fact
that complete oxidation, and also the dilution of a gas with an in-
different gas, lowers its luminosity
in constructing the Bunsen burner.
This is a piece of apparatus so
frequently used in the laboratory
that a few words conceming it
will suffice.

A Bunsen bumer is shown in
Fig. 31. The gas enters through
the horizontal tube, into the verti-
cal tube B. Air enters through
the hole C, and mixes with the
gas. The flame consists of two
distinct parts : an inner blue cone,
where the oxidation is far from
complete, and which is known
as the reducing flame, since it
has remarkable power to com-
bine with oxygen and reduce sub-
stances such as the oxides of the
metals ; and an outer, almost non-luminous tip, where the oxidation
of the gas is completed and where the temperature is very high.

—A

Fia. 31.

294

PRINCIPLES OF INORGANIC CHEMISTRY

k

This is known as the osddhing Jtame^ on account of its power to gire
up oxygen to substances which can be oxidized. Metals, for example,
in tills flame are usually converted into oxides.

By means of the Buusen burner very high temperatures can be
secured by the combustion of illuminating gas, without the produc*
tion of any appreciable quantity of light. This is due to the com-
paratively complete oxidation of the coniixjuuds of carbon in the gas^
by the excess of oxygen in the air which is mixetl with the gas.
When a cold object is inserted into the flame of a Bunsen burner, no
carlxm ia deposited upon it in the form of soot, and the lamp is,
theref orey very convenient for heating where cleanliness is absolutely
essential. The Bun sen burner is one of the most frequently used
pieces of apparatus in the chemical latioratory.

The Bhwpijn is a still more efficient means of obtaining a clean
oxidizing, and a clean reducing flame, and of directing these flames

where they are de*
sired. The blowpipe
itself is shown in Fig*
32. It consists of a
tube, ?, into which the
breath is blown from
the mouth, and a tube,
ti, at right angles to
this, through which
the air from the
lungs passes into the
flame. Into the top
of an ordinary Bun sen
burner is inserted a
tube with a narrow
oi)cning, so as to give
a narrow flame. The
blowpipe is placed
upon the upper edgB
of this tube as indi-
cated in the drawing,
and the breath ex*
pel led continuously
through the tube. The combustion of the gases in the flame is
excellent, the flame taking the form showu in the figure. The inner
flame, a, is the reducing flame, and the outer tip, b, the oxidizing
portion of the flame,

*t.

7ra. 33,

CARBON 295

By means of the blowpipe flame very delicate work can be done.
In the reducing flame small quantities of metal oxides can be reduced
to the metallic condition, and identified. The blowpipe is, there-
fore, an aid in detecting the presence of small quantities of sub-
stances, and in skilful hands is of great assistance in qualitative
analysis.

Effect of cooling the Flame. — The effect of cooling the flame can
be readily shown by means of the following experiment: Open an
ordinary gas stop-cock, and a short distance above the orifice hold a
wire-gauze with fine mesh. Light the gas above the gauze, and it
will burn without the gas below the gauze taking fire. This is due
to the fact that the metallic gauze conducts the heat away so rapidly,
that the gas below the gauze is not heated to its kindling temperature,
and does not ignite.

This principle was made use of by Sir Humphry Davy in the
construction of his safety lamp, for use in mines where explosive
gases are liable to accumulate. The flame is simply surrounded by
a fine wire-gauze. If the explosive gases should ignite on the inside
of the gauze, the flame cannot propagate itself through the gauze,
since it is too greatly cooled. The gauze, thus conducting the heat
from the flame, prevents the gases on the outside from becoming
heated to their kindling temperature, and thus explosions are avoided
when lights are carried into an atmosphere containing explosive
gases.

The Acetylene Light. — In the last few years carbon has come to
play a new role as an illuminant. A new method has been discovered
for preparing acetylene, which makes it possible to use this substance
in illumination. The method of preparing acetylene is based upon
the combination of carbon with many of the metals. These com-
pounds, known as carbides, have recently been made in large numbers
in the electric furnace by Moissan. The compound with calcium, or
calcium carbide, may be taken as the type. This is formed in the
electric furnace from a mixture of lime and carbon, and has the com-
position CaCj. When this is treated with water, calcium hydroxide
and acetylene are formed : —

CaCa -f 2 H,0 = Ca(OH), -f- C,H,.

Since the critical temperature of acetylene is 37®, and its critical
pressure 67 atmospheres, it can be readily liquefied. It is, however,
not preserved in the liquid condition in cylinders, like carbon dioxide,
on account of its explosive nature. It is generated as needed, by
allowing water in small quantity to come in contact with calcium

296

PROTCIPLES OF INOUGANTC CHEMISTRY

carbide when acetylene is desired. By regulating the flow of water
the rate of production of acetylene can be controlled. Acetylene
lamps are based upon this principle.

The amount of heat which is set free when acetylene ia bnmed is
very great indeed, being for a gram-molecular weight 310,000 calo-
ries. When acetylene is completely oxidized, the products are^ as
we would expectj carbon dioxide and water: —

2C»Ht + 50, = 2H,0 + 4CO|,

The Welsbach Light. — The Welsbach light differs from the or-
dinary gas-light in that solid substances are introduced into the
flame, which, when hot, have remarkable lighl^giving power. The
Welsbach light depends for its value entirely upon the "mantle,"
The mantle consists of a mix tare of thorium and cerium oxides* It
is prepared as follows : Fine cotton thread is woven into exactly the
form of the mantle. This is saturated with an aqueous solution of
a mixture of the nitrates of thorium and cerium. This mixture
contains 99 per cent of the thorium salt and one per cent of the cerium
salt The mantle is then dried and liighly heated to burn out all
organic matter, and to convert the cerium and thorium nitrates into
the oxides. It is then ready for use.

It is a remarkable fact that if the amount of cerium salt added
to the thorium is either increased or diminished appreciably, the
light-giving power of the Welsbach burner is greatly diminished.

The value of the burner is to be found in the power of these
oxides to convert heat energy in large quantity into light energy, so
that the final result is a conversion of more of the intrinsic energy of
the carbon compounds and other substances in the gas into light energy.

The Electric Light. — At first sight the relation between carbon
and the electric light may not appear to be very close, other than the
use of carbon as the source of energy to drive the dynamo which
generates the electrical energy. Indeed, the intrinsic energy of the
carbon, through the steam-engine and the dynamo, is converted into
electrical energy.

To obtain light energy from electrical energy, a resistance to the
passage of tlie electrical current is interposed. The current is usu-
ally passed between two carbon poles, which are heated to such a
high temperature that the carbon is partially volatilized. At this
temperature the highly heated cxirbon gives off an enormous amount
of light energy, and this is the source of the light in the electric are-
light In the incandescent light the carbon is heated white-hot, and
gives out light without undergoing any appreciable change.

CARBON 297

Measurement of the Belative Intensities of Light — We are famil-
iar with light of various degrees of intensity. Indeed, the intensity
of light varies from the brightness of the sun, to light which is so
feeble that it just produces a sensation when allowed to fall on our
retina. It is obviously desirable that some means should be avail-
able for measuring the relative intensities of light from different
sources. A number of instruments have been devised for this pur-
pose. These are known as photometers.

A very simple form of photometer has been devised by Bunsen, and
this will be briefly described. If a piece of paper is covered with oil
in one spot, and this spot observed in transmitted light, it will allow
more light to pass through than the remainder of the paper and will,
therefore, appear bright. If, on the other hand, it is observed in
reflected light, it will appear dark for the same reason, t.3. it trans-
mits more of the light and reflects less of it. If now a light of
standard intensity is placed upon one side of the spot, the light
whose intensity it is desired to compare with the standard is placed
upon the other side. The second light is moved towards or from
the spot until the spot disappears, i.e. until the spot has exactly the
same brightness as the remainder of the paper. When this condition
is reached the spot is illuminated on the two sides with exactly equal
intensity. Knowing the intensity of the standard light, its distance
from the screen, and the distance of the second light from the screen,
we have all the data necessary for calculating the relative intensity
of the second light. Intensity of illumination is usually expressed
in terms of candle-power. This means the amount of light given out
by a candle of a certain composition and certain dimensions burning
at a certain rate.

The science of photometry is especially useful in connection with
the light-giving power of coal-gas. This is frequently tested to see
whether it comes up to the desired standard.

CHAPTER XXIII

SILICON (At. Wt. =28.4)

The second member of the fourth group in the Periodic System
lis siUcon, This element la very widely distributed o?er the surface
*of the earth, and constitutes an important part of most rocks, Silicoa
occurs in great abundance as the dioxide, and forms an acid — silicic
,acid, whose salts make up many of our best-known rocks* Silicoa .
dioxide, or quartz, also occurs in huge masses, and is a constituent of
many rocks, especially granites, gneisses, etc, Silicon dioxide occura
in great abundance as sand, especially along the edges of large bodies
of water.

The Element Silicon. — Silicon is prepared from its compounds
by a numl>er of methods. One of these consists in heating silicon
tetrafiuoride with sodium. SiF4 + 4 Na — 4 NaF + Si. Another
method consists in heating sodium silicofluoride Ka^jSiFfl with metailio
sodium or potassium. The following reaction takes place : —

KagSiF, + 4K = 2NaF + 4KF + Si.

When this reaction is carried out some metallic zinc is added,
which melts and dissolves the silicon when formed. The zinc is then
dissolved in an acid and the crystallized silicon remains behind.

Silicon ia also formed when the dioxide is heated with a metal
like potassiumj sodium, magnesium, etc, : — ^

SiOj + 4 M ^ 2 MjO + SL

Silioon unites with fluoriue at ordinary temperatures forming the
tetrafluoride, SiF^, It combines with oxygen, chlorinej etc,, at elevated
temperatures.

Silicon exists both in the amorphons and crf/niftlUne condition*
The amorphous form^ obtained by the reduction of tlte dioxide or i
halide with metals, readily combines with oxygen, forming the
dioxide, also reacts with hydrofluoric acid, forming silicon tetraflu*
oride, and combines with a strong alkali^ forming a salt of silioia

^^^ 2 KOH + Si + H,0 ^ 2 H, + K,SiO^

298

SILICON 299

The potassium silicate formed is a salt of metasilicic acid, HjSiOa.

When amorphous silicon is heated to a high temperature it melts,
and on solidifying is crystalline. Crystallized silicon is best obtained
by dissolving molten silicon in molten zinc and allowing the mass to
cool. When the zinc is dissolved in acids the silicon remains as
grayish-black crystals, with a metallic lustre resembling graphite.

Crystallized silicon, which is the analogue of crystallized carbon,
has very different properties from amorphous silicon. It is much
less readily attacked by chemical reagents. It is not attacked by
oxygen at a white heat It is, however, attacked by fluorine at ordi-
nary temperatures, and by chlorine at elevated temperatures, forming
the fluoride or chloride of silicon. When highly heated with an
alkaline carbonate it forms the corresponding silicate.

Crystals of silicon are noted for their extreme hardness.

We have in these two varieties of silicon unquestionably the
analogues of amorphous and crystallized carbon.

Silicon Hydride or Hydrogen Silicide, SiHf. — Silicon forms a
compound with hydrogen, known as silicon hydride or hydrogen
silicide, containing one atom of silicon united with four atoms of
hydrogen. It is prepared by treating compounds of silicon with
aluminium or magnesium, with an acid.

The reaction in the case of magnesium hydride is represented
thus : — g. j^^^ 4- 4 HCl = 2 MgCl, -f SiH^.

Hydrogen silicide thus prepared is spontaneously inflammable
when it comes in contact with the air. The pure gas, however^ is
not inflammable at ordinary temperatures by mere contact with the
air, but ignites when slightly warmed. The fact that the gas takes
fire and combines with oxygen at ordinary temperatures is, therefore,
probably due to small quantities of some other substance, which is
present as an impurity.

As far as composition is concerned, silicon tetrahydride is analo-
gous to methane, — SiH4, — CH4. The former, however, is very
unstable, while the latter is quite stable. When silicon hydride is
burned in the presence of the air the products are silicon dioxide and

WJltPP • •^—

SiH4 + 208=Si02-f2HA

This is also analogous to methane, which yields on combustion
carbon dioxide and water.

Silicon also forms with hydrogen the compound SisHe, which is
spontaneously inflammable.

Silicon Dioxide, SiO^ — Silicon forms one compound with oxygen

800

PRINCIPLES OF INORGAOTC CHEMISTRY

— silicon dioxide. This is analogous to carbon dioxide. It does not

form the analogue of carbon monoxide,

Silicon dioxide occurs in nature in ^eat abundance. It is beau-
tifully crystalline in several varieties of quartz, sucb as aimthy^^
rock cn/stalf and tlie Uke, and with certain Imparities which give it
color it is of more or less value as gems^ such as £?pa^, jumper, onyx^
ugmUy etc.

It occura in great masses in less attractive forms, such as
quartz, sandjjlint, mmhtottet and the like, and is frequently the chief
constituent of large mountain ranges. When we consider the
abundance of the two formsj quartz and saudstoucj we can see the
importance of the element silicon in the iuorganic world, and from
a geological standpoint Silicon dioxiiie is also taken up by certain
plants, but its importance in the organic world is very small as com-
pared with the element carbon.

Silicon dioxide is very resistant to chemical reagents, and is not
attacked by aisids, with the exception of hydrofluoric, "When pow-
dered very iinely and fused with a caustic alkali, or an alkaline car-
bonate, it is transformed into a silicate : —

SiOg + 2 KOH = KjSiO, -h H^O.

This is a salt of metasilicic acid, having the composition H^SiOa,
The Acida of SUicoil ^ — Silicon combines with hydrogen and oxy-
gen, forming a number of acids which, however, can all be regarded
as derived from oue mother-substance. When an alkaline silicate
like that mentioned above is treated with an acid, the following
reaction probably takes place ; —

KaSiO, -h 2 HCl + H,0 - 2 KCl + H^SiO*.

This same compound is formed when silicon tetrafluoride is
treated with water : —

SiF, + 4 H,0 = 4 HF + H^SiO^.

The compound Il^SiO^, known as nomtcd mUcie mkl, or ortho-
silicic acidj is a white, gelatinous mass^ insoluble in water. To
separate the silicic acid from impnrities, such as potassium chloride,
the ordinary methods of washing are not sufficient. The silicic acidl
forms such a finely divided, jelly-like mass that it is scarcely pos-J
sible to dissolve out substances which are readily soluble in water^
because of the difficulty of securing good contact between the water j
and the substances, and the further difficulty of removing the solu-
tion when they are once dissolved.

SILICON 301

A new method of purification can be made use of in the case of
silicic acid. This is based upon the fact that substances like the
salts which easily form crystals, and many other substances, pass
readily through certain vegetable membranes, while other classes of
substances which do not crystallize, like silicic acid, do not pass
through such membranes. If a mixture of a salt like potassium
chloride and silicic acid is placed in a vessel whose bottom is closed
with vegetable parchment, and the vessel dipped into water, when
water is added to the mixture the salt will pass out through the
parchment and the silicic acid will remain behind. An apparatus
of this kind is known as a dialyzer, and the process as dialysis.
Substances which pass through such a membrane, since they gener-
ally form crystals, are known as crystalloids; while substances which
do not pass through such membranes are known as colloids. A large
number of substances belong to the colloids. These include starch,
albumen, and the finely divided metals, which will be considered
later.

Silicic acid, containing crystalloids as impurities, is allowed to
remain in the dialyzer for a time, and then the water in the outer
and inner vessel is removed and pure water added to both vessels.
This process is repeated a few times when all the crystalloids will
have passed through the parchment into the outer vessel, and have
thus been separated from the silicic acid.

Solutions of the colloids are not true solutions, but only pseudo-
solutions. This is shown by the fact that they do not lower the
freezing-point of the solvent, do not produce a rise in the boiling-
point of the solvent, and do not exert any osmotic pressure. Since
they do not exert osmotic pressure, they have no power to diffuse —
diffusion being caused by osmotic pressure. This is doubtless the
chief reason why the colloids do not pass through the parchment in
dialysis.

When normal silicic acid is heated it loses water and passes over
into metasilicic acid : —

H,Si04 = H,0 -f HjSiO,.

When it is further heated it loses more water and forms silicon
dioxide ; —

H^iO, = H,0 4-SiO^

Silicon has the power of combining with hydrogen and oxygen^
forming complex molecules which have acid properties. These are
known as polysilicic acids. They can all be regarded as derived
from the acid H4Si04, by removal of one or more molecules of water
from two or more molecules of the acid. Thus, by the removal of

302 PEINCrPLES OF INORGANIC CHEMISTEY

one molocule of water from two molecules of normal silicic acid
we have ; — ^ n^m\ = H,0 + H,Si A-

By removing two molecules of water : —

2 US^O^ ^ 2 H^O + H^SijO^
By removing three molecules of water : —

2H,SiO, = 3H,0 + H,SiaOB.

From three molecules of silicic acid we may, similarly, remove
one, twop three, molecules of water i —

3 H,SiO, = H^O + H,„Si Ai.

3 H,SiO, =: 2 H,0 + H«3i.0i^

3 H^SiO, - 3 H,0 + USW^

3 H^SiO* = 4 HjO + BfihO^

a H,8iO, = 5 IW + H^Si^O^

This series of acids, some of whose salts are known, suggest
the homologous compounds of carbon. The constant difference with
carbon is the group CHg. The constant difference with silicon is the
molecule of water H^.

Borne of the salts of the polysilicic acids are very important
substances, since they constitute many of the most abundant silicates.

The silicates are in general very stable, and with the exception
of the alkaline silicates, very insoluble substances, and hence are not
dissolved in appreciable quantities by the waters which come in
contact with them.

Conversion of Silicates into Carbonates, — N'otwithstanding the
great stability and insolubility of the silicates, they are being decom*
posed all over the surface of the earth by such a weak acid as
carbonic atiid. This mrbomUion ia taking place all over the surface
of the eaith, wiierever the carbon dioxide in the air and in the waters
comes in contact with silicates. This at first sight is very surprising*
How is it possible for such a weak acid as carbonic acid to displace
silicic acid from the very sta,hle silicates? This is es|jecially diffi-
cult to understand when we consider that carbonic acid is so easily
volatile and^ therefore, escapes from the field of chemical action.

The explanation is to be found in the effect of luftss on chemmd
wdivity^ as was pointed out by the German, Heinrich Rose, This is
one of the very best examples of mass action, as conditioning the
direction as well as the magnitude of chemical action.

The great amount of carbon dioxide in the air and in the water,

SILICON 803

acting slowly but continually for a long period of time, effects a
reaction which, in the laboratory, would be impossible.

The above process, known as the weathering of the rocks, is of
great geological and economical importance. By this means in part,
many of the most resistant rocks are decomposed and the surface
of the earth greatly changed in appearance. This process is of
tremendous economical importance in that the constituents of rocks
are made available for plants. The alkalies and other substances are
set free in the main as carbonates, and are either taken up by the
various plants, or are absorbed by the soil and retained until needed
by vegetation.

By absorption is meant the adhesion of solid matter to the surface
of the soil particles, and this has been shown to be a valuable principle
in connection with the fertilization of, and retention of, soluble mate-
rials in the soil.

We have had a number of examples of the effect of mass on
chemical activity. The law which governs this action has already
been formulated.

Compounds of Silicon with the Halogens. — It has already been
mentioned that silicon combines with chlorine directly at an ele-
vated temperature. The compound formed is silicon tetrachloride^
SiCl4, the analogue of carbon tetrachloride, CCI4, The analogy ex-
tends fai*ther, in that silicon can combine with hydrochloric acid,
forming the compound SiClaH, which is known as silicon chlorqform.
This is the silicon analogue of chloroform, which is trichlormethane
— CCI3H.

Silicon combines with bromine forming the tetrabromidey SiBr4,
and also silicon bromoformy SiBrgH, the analogue of bromoform —
CBraH.

With iodine silicon forms the tetraiodide, Sil4, and also silicon
iodoform, SilgH, the analogue of iodoform CI3H.

The compound of silicon and fluorine, SiF4, is of special interest,
since it is the compound formed when hydrofluoric acid acts on
glass. Silicon tetrqfluoride also decomposes with water, yielding an
acid of remarkable composition. Silicon tetrafluoride is formed by
the action of hydrofluoric acid on silicon dioxide. Since hydro-
fluoric acid is prepared most conveniently by the action of sulphuric
acid on calcium fluoride, silicon tetrafluoride is prepared by mixing
sand, calcium fluoride, and sulphuric acid.

When silicon tetrafluoride is treated with water the following
reaction takes place : —

3 SiF4 -f 4 HjO = H4Si04 + 2 H^iF«.

304 PRINCIPLES OF INORGANIC CHEMISTRY

The compound HsSiF^ is known as hydrofluosilicic acid. It is
readily soluble in water, having strongly acid properties. With
alkalies it forms salts of the composition M^SiFe. It is, therefore,
a dibasic acid, dissociating into —

H,H,SiFe.

When salts of hydrofluosilicic acid are heated they decompose into
the corresponding fluoride and silicon tetrafluoride : —

M,SiF«=2MF + SiF4.

When hydrofluosilicic acid is treated with an excess of an alkali,
it decomposes in the sense of the following equation : —

H^iFe + 6 MOH = 6 MF +2 H^O + Si(OH),.

The salts of hydrofluosilicic acid are generally soluble in water,
with the exception of certain salts of the alkalies and alkaline earth
metals. These will be considered later.

Compoand of Silicon with Carbon — Carbonmdnm. — From the
many analogies between silicon and carbon, we would not expect
these two elements to combine and form any very stable compound.
The facts are, however, quite different. Silicon and carbon form a
very stable compound, as is shown by the method of its preparation.
When finely powdered sand is mixed with carbon and sodium chlo-
ride, and the mixture subjected to the highest temperature of the
electric furnace (3500**), the following reaction takes place : —

3C-f*SiOj = 2CO-f SiC.

Carborundum is characterized by its extreme hardness, and is
useful technically on account of this property. It is extensively
used to cut glass, and in other connections where the diamond was
formerly employed. It is very resistant to chemical reagents, not
being attacked to any appreciable extent by any of the acids. When
fused with the strong alkalies it is decomposed. When powdered
and heated in a stream of oxygen, the carbon is burned out only
with great difficulty.

CHAPTER XXIV

GERMANIUM, TITANinM, ZIRCONIUM, CERIUM, THORIUM

Next to silicon, in group IV, comes germanium, which will be
briefly studied. Then come tin and lead, which will be taken
up much later under the metals. The four elements, titanium,
zirconium, cerium, and thorium, will be dealt with in the present
connection.

Oermanium (At. Wt. = 72.5). — Germanium is of interest in
that it was one of the elements predicted by Mendel^ff from the
Periodic System. It was discovered in 1886 by Clemens Winkler,
in the mineral argyrodite, which is the double sulphide of silver and
germanium, 4 Ag^.GeS,.

The element germanium is formed by reducing the oxide with
metallic magnesium, magnesium oxide and germanium resulting.
It forms two series of compounds, in one of which it is bivalent
and in the other quadrivalent. The more important of the bivalent
compounds of germanium are the hydroxide, Ge(OH)^ and the sul-
phide, GeS.

The tetravalent compounds of germanium resemble those of sili-
con. It forms the dioxide GeO„ the tetrahydrate Ge(0H)4, the
tetrachloride GeCl4, the tetrafluoride GeF4, the disulphide GeSj,
and so on. The analogy to silicon is further shown in the com-
pound KjGeFe, which is the analogue of K,SiFe.

Titanium (At. Wt = 48.1). — Titanium occurs in fairly large
quantities, and is widely distributed over the surface of the earth.
It occurs in the minerals tUanite, rutiley bauxUey etc. Titanium com-
bines with oxygen, forming the compounds TiO, Ti^Os, and TiO^.

It forms, with oxygen and hydrogen, titanic acid Ti(0H)4, which
is the analogue of silicic acid. It also loses water and forms metar
titanic acid, HgTiOs, the analogue of metasilicic acid.

Like carbon, titanium forms the tetrachloride TiC^ It also
forms the trichloride TiClj and the dichloride TiClj.

The analogy to silicon is shown in the compound potassium
titanofluoride K^TiFe, which is the analogue of potassium silico-

z ao6

306 PRINCIPLES OF INORGANIC CHEMISTRY

fluoride. Titanium also combines with carbon, forming titanium
carbide TiC, which is the analogue of silicon carbide, carborundum.

Zirconium (At. Wt. = 90.6). — The element zirconium occurs
chiefly as the silicate ZrSi04. This is the beautifully crystallized
mineral zircon. Zirconium acts as a tetravalent element, forming
the hydroxide Zr(0H)4, the chloride ZrCl4, the sulphate Zr(S04)»
and so on. Zirconium forms the dioxide ZrOg, also the analogous
carbide ZrC,. It also forms the compound HjZrFe — Hydrofluo-
zirconic acid, the analogue of hydrofluosilicic acid.

Cerium (At. Wt.= 140.26) . — Cerium is one of that group of rare
elements which occurs in monazite sand. The element is obtained
by electrolyzing the chloride. While cerium forms with oxygen the
dioxide CeOs, it acts in most of its corapoimds as a trivalent element.
Thus, it forms the chloride CeClj, the sulphate Ce2(S04)a, the nitrate
Ce(N03)8. The double nitrate of cerium and ammonium, 2 Ce(NO,)a,
3 NH4N0a -f 10 HjO, is a beautifully crystallized substance.

Cerium can, however, form compounds in which it acts as a
tetravalent element, the compound Ce(S04)8 being known. As
already mentioned, cerium is used in small quantity in preparing
the mantles of Welsbach burners.

Thorium (At. Wt. = 232.5). — Thorium is another of the rare
elements which occur in monazite sand. It also occurs as the
silicate, thorite, and in many other minerals such as gadoHnite^
samarskitef etc. It forms the hydroxide Th(0H)4.

As has already been mentioned, thorium is the chief constituent
of the Welsbach mantle. Thorium compounds have been shown to
be radioactive.

CHAPTER XXV

BORON (At. Wt. = 11.0)

We pass now to group III of the Periodic System, the first mem-
ber of which is boron. This is the only member of this group which
has distinctly acid-forming properties. We shall, therefore, take up
boron in the present connection, and the remaining members of the
group considerably later when we come to study the base-forming,
or metallic, elements.

Occurrence, Preparation, and Properties. — The borates, or salts
of boric acid, are the chief source of the element boron. We should
mention especially boracite, borocalcitey and borax.

Boron is prepared by the reduction of the trioxide of boron.
When borates are treated with a strong acid, boric acid, H,BOa, is
liberated. This loses water on heating, forming the trioxide : —

2HsBO, = 3HjO-f BjO,.

The oxide is reduced by potassium, magnesium, etc., at a high
temperature, the oxygen combining with the metal and setting free
the boron. In preparing boron the nitrogen of the air must be ex-
cluded by a layer of borax, since nitrogen combines with boron at
high temperatures. Boron forms beautiful crystals, which are char-
acterized by their great hardness. They seem to have about the
same hardness as the diamond. Crystallized boron combines with
oxygen only slowly, even at very high temperatures. It unites
with chlorine at elevated temperatures. It is not attacked at
ordinary temperatures either by acids or alkalies.

Amorphous boron is much less resistant to chemical reagents.
It is much more easily oxidized, and far more readily attacked by
acids and alkalies than the crystallized form.

It will be remembered that boron is one of the elements whose
specific heat does not conform to the law of Dulong and Petit. The
specific heat of boron as determined at ordinary temperature was too
small to accord with the law of Dulong and Petit. It was, however,
found that the specific heat of boron increases with the temperature,

307

308 PRINCIPLES OF INORGANIC CHEMISTRY

until a temperature of about 500^ is reached. The specific heat then
becomes constant and accords very well with the law.

Boron Trioxide, B^O^. — The compound of boron and oxygen,
boron trioxide, BjOj, is formed either by removing water from boric

^^^~ 2H3BO, = 3H,0 + BA,

or by burning boron in oxygen, when the two elements unite and

form the trioxide:- 2B + 30 = B/)^

Boron trioxide, as seen by the first method described for its
preparation, is an anhydride of boric acid, which we shall now
study.

Boric Acid, HaBO^. — Boric acid occurs in nature in the free con-
dition. It is volatile with water-vapor, and in the region of certain
hot springs, as in Tuscany, it is brought to the surface of the earth
by the escaping vapors.

Boric acid is soluble in water, forming beautiful, white crystals
when the aqueous solution is evaporated. It is easily recognized by
the fact that its alcoholic solution burns with a characteristic green
flame. If boric acid, or a borate treated with sulphuric acid, is
treated with a little alcohol and the alcohol ignited, the flame
appears green throughout if there is much boric acid present. If
only a small amount of boric acid is present the flame is green only
on the edges.

The salts of the normal boric acid, HsBOj, do not exist. When
boric acid is heated, however, it loses water and passes over into
another acid, whose salts are well known. The first product of the
dehydration of boric acid is metaboric acid — HBO,: —

H,B03 = H20-|-HBOj,

Boric acid can, however, lose water in a different manner and
form an acid whose salts are well known : —

4 HsBO^ = 5 HsO + H,B A.

The acid HjB^O- is known as tetraboric acid, and its sodium salt,
Na2B407, is ordinary borax. Boi-ax melts easily, forming a colorless
liquid. This liquid has the power of dissolving certain metal oxides
and forming with them, when cold, glass-like masses which have char-
acteristic colors. The borax bead is, consequently, of importance in
blowpipe analysis for the detection of metals.

Boron Kitride, BH. — Boron combines with nitrogen, forming the
compound BN. Amorphous boron, when heated to a high tempera-

BORON 809

ture in the presence of nitrogen, combines with it and forms the
nitride. Boron nitride is an amorphous solid, very resistant to
chemical reagents. Boron nitride, when heated in the air, becomes
phosphorescent; with water-vapor it decomposes into ammonia and
metaboric acid. The existence of this compound makes it necessary,
in preparing boron, to protect the element, when formed at the high
temperature, from the air, otherwise the boron will combine in part
with nitrogen, and the result will be a mixture of boron and boron
nitride.

Compoimds of Boron with Other Elements. — Boron combines with
a number of the elements. With chlorine it forms the trichloride
BClg, with bromine the tribromide BBr,, with iodine the triodide BI,,
and with fluorine the trifluoride BFj. Boron also combines with sul-
phur, forming the trisulphide B,Ss.

Boron forms a salt with phosphoric acid having the composition
BPO4.

In the above compounds boron acts as a basic element.

Summary, — We have studied thus far oxygen, hydrogen, and the
halogens or members of group VII, in the Periodic System. The
analogues of oxygen ; sulphur, selenium and tellurium, in group VI,
were then taken up. The remaining members of group VI are so
distinctly metallic that they will be studied with the metals. The
entire group V was then studied, and all of group IV, with the
exception of the important metals, tin and lead.

We have begun the study of group III with the first member,
boron, which is distinctively an acid-forming element The remain-
ing members of this group, however, are so distinctively base-form-
ing or metallic that they will be studied with the metals.

Having completed our study of the non-metals, or metalloids, as
they are termed, we shall now turn to the metals, and with these
we shall begin with group I — the alkalies.

CHAPTER XXVI

THE MTTTAIiS

The metals have certain properties in common which distingTiish
them from the other elements. With one exception, they are all
solids at oi*diiiary temperatureB. They are good conductors of heat,
and for the most part good conductors of electricity. Their power
to conduct electricity^ however, varies considerably from metal to
metaL Some of the metals^ like sotUum^ potassium, etc,, arc very
active chemically, while othei-a, like platinumj gold, etc., are very
resistant to chemical reagents.

The chemistry of the metals is in general much simpler than that
of the metalloids. The metals form ions charged with positive elec-
ti-icity, — cations, — ^and these combine with the anions of acids,
forming salts. The cations are generally very much simpler than
the anions, consisting usually of single metal atoms charged ynth.
positive electricity. There are, however, exceptions to this general
statement; a metal may be in combination with other substances,
forming part of an aniom

Since when metals react chemically they pass into solution, ue,
into the ionic state, the study of the chemistry of the metals iB
largely the study of the ions which they form. Indeed, we have
excellent reason for believing that in order that the metals should
react chemically they must be in the ionic state. If tliis Ije true^
the study of the chemistry of the metals is in reality a study of them
in the ionic condition.

We shall now take up the metals one by one, and see what ar©
their peculiarities and the most interesting reactions into which they
enter. We shall not, however, take them up at random, since some
of the elements are very closely allied in their chemical properties,
while others show few and remote relationships. Here, agaiUj we
are greatly aided by the Periotlic System. In this system the ele-
ments which are allied chemically fall into the same groups. While
we shall be guided by this system, we sliall not hesitate to depart
from it in the case of the metals, where relations can be better seen

SIO

THE METALS 311

by doing so, as we have already departed from it in the case of the
metalloids.

We shall take up first the alkali metals, consisting of lithium,
sodium, potassiunr, rubidium, and caesium.

Next in order come the metals of group II. These fall, with
respect to their relationships, into two divisions; calcium, stron-
tium, and barium, on the one hand, and beryllium or glucinum,
magnesium, zinc, cadmium, and mercury on the other.

When we pass to group III we find that boron has already been
studied with the metalloids. The first metal in this group is alu-
minium. In the same group are the rare elements, scandium,
gallium, yttrium, indium, lanthanum, ytterbium, thallium, and
samarium.

Passing to iron, we have in this same group nickel, cobalt, man-
ganese, chromium, and the rarer elements, molybdenum, tungsten,
and uranium.

Next are taken up copper, silver, and gold, and then lead and
tin, the former appearing in group I, the latter in group IV.

Finally, among the noble metals, we have rhodium, ruthenium,
palladium, osmium, iridium, and platinum.

CHAPTER XXVII

THB AZ.KALI BfETALS

LITfflUM, SODIUM, POTASSIUM, RUBIDIUM, AND C.£SIUM

SODIUM (At Wt = 23.05)

The natural order in which to take up the alkali metals would be
to start with the one with the lowest atomic weight, lithium, and
proceed in the order of increasing atomic weights to caesium.
There are other reasons, however, why this order should not be
adopted.

Lithium is a comparatively rare substance occurring only in
relatively small quantities.

It is far better, in order to become acquainted with this group, to
take up an element which* occurs in large quantity, and which can be
readily obtained and worked with in the laboratory. Such an element
is sodium, the second of the alkalies in the order of increasing
atomic weight.

Occurrence of the Element Sodium. — The element sodium is very
widely distributed and occurs in combination with other elements in
many places in large quantities. On account of its great chemical
activity it does not occur in nature in the free condition. Nearly all
of the salts of sodium are soluble in water. We should, therefore,
expect to find most of the sodium compounds dissolved in the waters
of the sea, and such is the fact When the rocks undergo weather- .
ing and set the sodium compounds free, these dissolve readily in
water and are swept down to the sea. In this way compounds of
sodium, and especially sodium chloride, have been accumulating for
ages in the sea, and this is in part the explanation of the saltiness of
sea-water. Since practically all of the simple salts of sodium readily
dissolve in water, we do not find an accumulation of these salts in
regions where there is appreciable rainfall. In certain arid regions,
however, one of the most soluble salts of sodium exists in large beds.
In Chili large beds of sodium nitrate are found which, from their
analogy to potassium nitrate or ordinary saltpetre, are known as Chili
saltpetre.

813

THE ALKALI METALS 818

Sodium salts exist in great abundance in certain regions where
the waters of the sea have evaporated. In the great salt beds of the
earth such as those at Stassf urt, the chloride and other compounds of
sodium occur. One compound of sodium which has recently come
into prominence in connection with the manufacture of aluminium
should be mentioned. This is the double fluoride of sodium and
aluminium, Na^AlF,, occurring in Greenland and known as cryolUe.
Sodium occurs in small quantities practically everywhere and in
everything. We have a very sensitive means in the spectroscope of
detecting the presence of minute quantities of sodium. When almost
any substance is examined in the spectroscope it shows the presence
of traces of sodium. Indeed, the atmospheric air always contains
sodium. To obtain any substance free from sodium requires the
very greatest precautions. The universal presence of sodium seems
to be due to its existence in the atmosphere. The chloride is taken up
with the water- vapor over the sea, and distributed in minute quantity
everywhere.

Preparation of Sodium. — The preparation of the element sodium
is of special historical interest.

The compound which we know to-day as sodium hydroxide was
supposed for a long time to be an element. When Sir Humphry
Davy constructed his enormous voltaic battery in connection with
the Koyal Institution in London, he tried the action of the current
upon a large number of substances, and among these upon fused
sodium hydroxide. The result is well known. A metallic substance
separated at the cathode, which rose to the surface of the molten
hydroxide and took fire spontaneously on coming in contact with the
air. The compound nature of sodium hydroxide and the elementary
nature of sodium were thus proved beyond question.

Sodium was prepared for a long time by the reduction of the

oxide or hydroxide by means of metallic magnesium, or by highly

heated carbon : — xt r^ . nr nr,^ . xr
NaO-f Mg := MgO-f Na.

All of these reduction methods are now abandoned when it is de-
sired to prepare sodium on a large scale. The electrolytic method is
used entirely. Considerable sodium has been prepared by the elec-
trolysis of the fused chloride, but these methods involving the use of
the chloride are more difficult to carry out than the method employ-
ing the fused hydroxide of sodium.

Sodium prepared by electrolysis of the fused hydroxide is not an
expensive substance, the price having been reduced immensely by
the application of the electrolytic process.

314 PRINCIPLES OF INORGANIC CHEMISTRY

Propertiei of Ketallio Sodinxo. — Sodium is a 8oft solid, which,
when freshly cut with a knife, has a metallic lustra and a steel-gray
color. The surface becomes quickly tarnished, due to the rapidity
with which it takes up oxygen from the air or from moisture, form-
ing the oxide or hydroxide. Sodium is a very active substance
chemically. It combines readily with moist oxygen, but very slowly,
indeed, with dry oxygen. When heated in the presence of oxygen
it forms the peroxide NaO. In the presence of water the following
reaction takes place : —

2 Na+2 H,0 = 2 NaOH + Hj,

The hydrogen, which is liberated when apiece of sodium is thrown
upon water, does not take fire if the sodium is allowed to move about
over the surface of the water. If the sodium is held in one place, as
by throwing it upon a piece of filter-paper upon the water, enough
heat is produced to ignite the hydrogen.

With potassium, sodium forms alloys^ which are liquid at ordinary
temperatures. When one part of sodium is fused with four or five
parts of potassium the alloy formed is a liquid with metallic appear-
ance, resembling in some respects the am^gams or solutions of the
metals in mercury.

Sodium, in the presence of water, forms the hydroxide NaOH, as
we have just seen. This, we shall learn, is one of the very strongest
bases, and as we would expect combines with all acids. In the light
of these facts it is most remarkable that pei'fectly dry sodium does not
react with perfectly dry sulphuric acid. When sodium which has been
dried with the very greatest precaution is plunged into sulphuric
acid from which every trace of water has been removed, it remains
suspended in the acid without the least sign of chemical activity.
In such experiments unusual precautions must, of course, be taken
to remove the last traces of moisture. When there is any moisture
present the sodium forms with the water sodium hydroxide, which
dissociates, yielding hydroxyl ions, which would then combine with
the hydrogen ions resulting from the action of the moisture on the
acid. The above is one of the most remarkable facts in the whole
field of chemistry, if we try to interpret it in any other light than that
furnished by the new physical chemistry. In terms of the theory of
electrolytic dissociation and catalysis these facts are just what
would be expected, and could have been predicted before tliey were
discovered.

THE ALKALI METALS 815

COMPOUNDS OF SODIUM WITH OXYGEN AND HYDROGEN

Sodinm Hydride, KaH. — The hydride of sodium, NaH, is formed
when sodium is heated to 300^ in an atmosphere of hydrogen.

Sodium Peroxide, KaO. — ^ As far as is known with certainty sodium
forms only one compound with oxygen. This is the peroxide NaO.
It is obtained when sodium is heated in the atmosphere to about
300® to 350**. It is a light-yellow powder, and dissolves readily
in water, forming sodium hydroxide and hydrogen dioxide. The
reaction would be represented thus: —

2NaO + 2H,0 = 2NaOH+H,0^

If the temperature is not kept low, a certain amount of oxygen is
evolved, and the reaction would be represented thus : —

4Na0 4-2H,0 = 4NaOH4-0^

Sodium Hydroxide, KaOH. — We have just seen that one method
of preparing sodium hydroxide is to treat the oxide with water.
Another method, and perhaps the best for preparing a little sodium
hydroxide in very pure condition, is to allow water to act on metallic
sodium : —

2Na4-2H,0 = 2NaOH + Hs.

On account of the violence of this reaction it must be carried out
with certain precautions. If water were allowed to flow directly
upon the metal, the reaction would proceed with such an enormous
evolution of heat that an explosion would very probably result.
The best way to carry out the reaction is to place a piece of pure
sodium in a porcelain dish, and float the dish upon water in a larger
vessel — the whole being covered with a bell-jar filled with air from
which all carbon dioxide had been removed. The water-vapor
comes in contact with the sodium, and the reaction proceeds slowly
and without any indication of an explosion or spattering of the
alkali. Sodium hydroxide can also be prepared in very pure condi-
tion, by treating sodium carbonate with a solution of the hydroxide
of any metal which forms an insoluble carbonate; or by treating
sodium sulphate with the hydroxide of any metal which forms an
insoluble sulphate. If a solution of sodium carbonate is treated with
lime water, we have : —

NajCOa + Ca(OH), = CaCOs + 2 NaOH.

SW FRINXIPLES OF INORGANIC CHEMISTRY

If ;i $o>Iution of sodium sulphate is treated with a solution of
hjir:um b vdroxide, we have : —

Na^04 -f Ba(OH), = BaSO^ + 2 NaOH.

Sodium hydroxide is one of the strongest bases known, as is
seen by the following conductivities: —

r

Mrl8°

a

1

140

79.8%

10

170

90.4

100

187

99.5

it^) 500

188

100.0

AVhen brought into the presence of water it dissociates thus : —
NaOH = Na,0~H.

The basic nature of sodium hydroxide is due to the presence of
the hydroxyl ion, as has already been stated. When sodium hydrox-
ide is treated with an acid they react as follows : —

Na, OH 4- H, CI = Na, CI -f H^O.

This is the typical reaction between an acid and a base, an example
of which we have already met with in the case of ammonium
hydroxide. The cation of the base and the anion of the acid
remain in the same condition after neutralization as before. The
anion of the base, hydroxyl, and the cation of the acid, hydrogen,
unite and form water, and this is all that takes place in the process
of neutralization. To obtain the salt of sodium it is only necessary
to evaporate the solution after neutralization. In the above exam-
ple, when the water is removed, the sodium and chlorine ions unite
and form the salt sodium chloride.

The Chemistry of Sodium the Chemistry of the Sodium Ion. — The
chemistry of the element sodium is not the chemistry of tlie atom or
molecule of soilium, since there is every reason for believing that
these are prat^tieally inert. Perfectly dry sodium contains an abun-
dance of atoms or molecules, and yet will not react chemically. We
have already seen that dry sodium, free from every trace of moisture,
will not act ujxm perfectly dry sulphuric acid. Dry sodium will not
act on dry chlorine gi\s, but will remain molten in contact with the
gas without having even its surface tarnished. There is good reason
for believing that dry sodium will not act upon dry oxygen, and thus

THE ALKALI METALS 317

it goes through the list of chemical reactions which are characteris-
tic of this element.

It may be said that the presence of traces of moisture acts cata-
lytically, effecting the action without taking part in it. This is,
of course, a mere assumption and does not explain anything. It
seems far more probable that the presence of even traces of water
causes a slight dissociation of the substances, and as soon as the
ions thus formed are used up more ions are formed, and this may
continue until the reaction proceeds practically to the end.

On the other hand, wherever we have sodium ions present we
have the reactions which are characteristic of this element. Some
of these reactions, together with their products, we shall now study.

Compounds of Sodium with the Halogens. — The strong base, so-
dium hydroxide, combines with the strong halogen acids, as we would
expect, forming salts. Or, more accurately stated, the cation sodium
combines with the halogen anion, forming the sodium salt of the
halogen.

Sodinm Chloride is found in great abundance in sea-water, which,
on the average, contains about 2.7 per cent of the salt. The amount
of sodium chloride in sea-water, however, varies greatly from one
locality to another. In tropical regions, where the evaporation of
the water is relatively rapid, the concentration may be as much as 3.6
to 3.8 per cent. Where large bodies of fresh water pour into the
sea the percentage of sodium chloride may be reduced below unity.
In certain isolated bodies of water, as the Dead Sea, the amount of
sodium chloride may be more than 20 per cent. It also occurs in
the solid form in a number of the great salt-deposits, and especially
in those of Salzburg, Germany. This salt is also obtained from sea-
water by evaporation. The water is allowed to flow into shallow
pools, and be evaporated by the heat of the sun. After this has been
repeated a sufficient number of times the salt is removed and purified.

Sodium chloride is also obtained directly in the solid form from
many of the great salt beds. Sodium chloride crystallizes in char-
acteristic, hopper-shaped cubes. These cubes are not completely filled
out, but are hollow in the centre. It melts at about 780**, decrepitat-
ing or flying to pieces when heated. The salt is nearly as soluble
in cold water as in hot, one part of water dissolving 0.36 parts of
sodium chloride. Sodium chloride is present in considerable quan-
tity in the animal body, and is especially sought for by herbivorous
animals.

On account of its abundance and cheapness, sodium chloride is
of great importance as a source of both chlorine and sodium. When

318

FHINCIFLES OF IKOHGANIC CHEMISTRY

electrolyzed, either in the molten condition or in concentrated solu-
tion, chlorine is set free at the anode. When treated with a strong
non-volatile acid such as sulphuric, hydrochloric acid is evolved. It
is readily transformed into otlier compouuJa of sodiujii, including
the hydroxide, and is thus the source of such comi>ounds as well as
of the elejuent itself.

Purification of Sodium Chloride. — Since sodium chloride is just
ahout as soluble in euld us in hot water, the ordinary method of puri-
fication, based upuu fractioual crystallization, cannot be verj^ success- '
fully applied in this case. A method of puriheation which inFolves
a general principle can, however, be applied.

If to a satitrated solution of sodium chloride either sodium ions or
cMormB tans are added, some of the sodium chloride will be precipi-
tated. The most convenient method of adding chlorine ions to a
solution without increasing the amount of the solvent present^ is to
pass into the solution hydrochloric acid gas. If hydrot*hloric acid
gas is passed into a saturated solution of sodium chloride, some of
the latter compound is precipitated, and the amount precipitated
depends upon the amount of acid which is run into the solution.

We can test the first part of the statement, that an addition of
sodium ions will cause a precipitation of the sodium chloride with
which the soUitiou is saturated^ by adding to such a solution some
solid, sotHum salt, such as the nitrate. The sodium nitrate will dis-
solve in the saturated solution of the chloride, and dissociate into

sodium ions and the nitro-ion NO^ A part of the swUum chloride
in the saturated solution will be precipitated in solid form.

There is a general principle involved here, deduced by the Ger-
man physical chemist, Nernst, from the law of the action of mass,
with which we have already become familiar. In any saturated
solution the produet of the number of cattojis timeii the nnmher of
^niojis is a mnMant This condition always obtains for a satumted
solution, and may be known as the Hw ofmtnralion.

If to a saturated solution of sodium chloride we add either com-
mon ion, sodium or chlorine, the solution will still be saturated withj
sodium and chlorine ions, and only saturated.

In order that the relation —

cations x anions = constant

should obtain, if we increase the nuiDber of cations, the number of
anions present must diminish ; or if we increase the uuml>er of anions,
the number of catious present must decrease. The way in which
this can occur is for a certain number of cations to combine with an

THE ALKALI METALS 319

equal number of anions, and form molecules of the salt. The solu-
tion is, however, saturated with respect to the salt, and as quickly as
any of the salt is formed by a combination of its ions it is precipi-
tated. It is obvious from the above equation that the larger the
excess of cations or anions added to the solution, the greater the
amount of the salt which will be precipitated.

This principle is taken up here because it is an excellent means
of purifying chlorides. Hydrochloric acid being volatile, it can
readily be conducted into a saturated solution of any chloride, when
some of the chloride with which the solution is saturated will be

precipitated, and only this substance. If the nitro-ions NOj could be

added to a saturated solution of a nitrate, or the sulphuric ions SO4
to a saturated solution of a sulphate, we should have a part of the
original nitrate or sulphate precipitated.

The physical chemical importance of this law of saturation is
very great indeed, since it has led to a "method of measuring electro-
lytic dissociation, which, however, it would lead us too far to discuss.

Sqdium forms with chlorine a subchloride, NajCl. This is
obtained by heating sodium chloride with metallic sodium at a
high temperature. It is deep-blue in color.

Sodium Bromide (KaBr) and Sodium Iodide (Kal). — These salts
resemble sodium chloride so closely that a detailed study of them
is not necessary. Certain phenomena connected with their forma-
tion from aqueous solution and with their solubility are of interest.

If these salts are crystallized from hot, aqueous solutions, they
come down in the anhydrous condition. If the temperature at
which they crystallize is below 30®, they come down with two
molecules of water of crystallization. If the salts containing water
of ciystallization are heated, the bromide to oO®, the iodide to 67®,
they form the corresponding anhydrous salts and saturated solu-
tions of these salts.

If we study the solubility of these two salts with rise in
temperature and plot the results as curves, the abscissas being
temperatures and the ordinates parts of salts in oue hundred parts
of water, the curves would have the form shown in Fig. 33.

The curves show that the salt with water of crystallization
increases in solubility with rise in temperature until the transition
point, where the two curves intersect, is reached. After this tem-
perature is passed the solubility of sodium bromide remains nearly
constant as the temperature rises, while the solubility of sodium
iodide increases with rise in temperature, but much more slowly

820

PRINCIPLES OF INORGANIC CHEMISTRY

than the solubility of the iodide with water of crystallization. The
curves are all plotted beyond the transition-points as dotted lines.
This means that we may have either phase extending into the region
of the other in a metastable condition.

020

280

240

200

160

120

80

40

f-

"'TirT

J

/

/

^

^

rt.\.^

K^

^

^^

-2^-'

^

/

NaBr

-^

^

X*^

— '

•

--

-32-

I^nTcl

—

^

-20'*

0**

20-

40** 60*
FlO. 33.

80'

100*

120*

140*

Ostwald has shown that the merest trace of the phase which is
stable under the conditions, is sufficient to cause the metastable to
pass over into the stable condition.

Sodium Hypochlorite (KaOCl), Chlorate (KaClOs), and Bromate
(VaBrOs). — Sodium forms salts with the oxygen acids of the halo-
gens, which are well-defined, stable substances, but for the most part
are without special chemical interest.

The Jij/pochlorite is used now to some extent as a disinfectant,
and in bleaching.

The chlorate is formed by the action of chlorine on sodium
hydroxide : —

6 NaOH -f 3 Clj = 6 NaCl -f NaClO, -f 3 H,0.

Sodium chlorate, unlike potassium chlorate, is quite soluble in
water, and is, therefore, much more difficult than potassium chlorate
to separate from the corresponding chloride. It crystallizes in

THE ALKALI METALS 821

cubes, and has one property of more than the average interest
When a beam of polarized light is passed through the crystal, the
plane of the beam is turned around an axis. As we say, it has the
power to rotate the plane of polarization. There are many sub-
stances known which have this power, but they are mainly in the
field of organic chemistry.

Sodium bromate is prepared in a manner which is analogous to
the preparation of the chlorate, ue. by the action of bromine on
sodium hydroxide.

Sodium Triazoate (KalTs) and Sodium Amide (KaHHs).— The so-
dium salt of triazoic or hydrazoic acid is well known. It is formed by
the action of the hydroxide or carbonate of sodium upon hydrazoic
or triazoic acid. The salt crystallizes in cubes and is remarkable
for its composition, consisting of a sodium atom in combination
with three nitrogen atoms. As we would expect, such a compound
is imstable, and in the dry condition easily explodes.

Sodium amide, NaNH2, is formed when perfectly dry ammonia
gas is passed over heated sodium. The action is as follows : —

2NH3 + 2Na = 2NHjNa + H^

sodium replacing one hydrogen atom in ammonia and combining
with the residue NHj. When sodium amide is treated with nitrous
oxide the sodium salt of hydrazoic acid is formed : —

2 NaNH, + N,0 = NaOH + NH, + NaN,.

When the sodium salt is treated with a strong acid, triazoic acid
is formed, and this is the simplest means of preparing this substance.

Sodium Kitrate, KalTOs. — Sodium nitrate is called Chili saltpetre
because it is found in a certain, rainless district between Chili and
Peru, and since it is the sodium analogue of potassium nitrate or
ordinary saltpetre. It is extremely soluble in water, and, therefore,
could not exist in the solid condition in regions where there is
appreciable rainfall. To give an idea as to its extreme solubility a
few data are added. At the following temperatures, one part of
water dissolves so many parts of sodium nitrate : —

TSMPKBATUBB

NaXOs, Pabtc dimolybd bt onb Pabt
Watbb

0°
80*»

0.78
1.02
1.60
2.00

322

PRINCIPLES OF INORGANIC CHEMISTRY

Sodium nitrate, being so readily soluble m water, forms in aqueous
solution, a solution of NO^ ions. These aie taken up by the plants
ami, if not too concentrated, ai'e among the most valuable fertilizing
agents. Considerable quantities of sodium nitmte are added directly
to the soils as a fertiliiter, where quick results are desired. On ac^
count of its great solubility, it is quickly accessible to plants^ and for
such products as are common to the garden it is one of the %'ery
best fertilizers.

ScHiium nitrate cannot be used in making gun-powder in the place
of ivotassium nitrate^ since it absorbs water from the air, or iB
deliqitesfeeni, as we say. It is, however, extensively used in the prep*
aration of potassium nitrate. It is also extensively used in the
preparation of ammonium nitrate, free nitric acid, and of sodium
nitrite.

Sodium Nitrite^ TSIbJUO..— The nitrites can he obtained in general
by heatiog the nitrates. Unleaa tlie heating is very carefully done,
and even then, there is considerable decomposition of the nitrite.
The best method of preparing sodium nitrite is by fusing the nitrate
with some mild reducing agent such as metallic lead : —

KaXO, -h Pb = PbO -h NaNO^

Sodium nitrite is extensively used in the preparation of artificial
dyestuffs. When treated with an acid, sodium nitrite breaks down
thus : —

2 ^aNO, + H,SO, = Na^SO, + 2 HNO^

The nitrous acid, however, undergoes decomposition : —

2HK0s=H,0+N^s,

and the gas KjO^ is useful in effecting certain reactions in organic
chemistry.

Sodium Hydrosulphide (ITaHS) and Sodium Sulphides (ITa^ to
FaJBj). — Sodium liydrosidi>liide i*^ formed wlien a solution of sodium
hydroxide is saturated with hydrogen sulphide gas: —

NaOH + H,S =H,0 + NaSH.

If to a solution of sodium hydrosulphide an equivalent of sodium
hydroxide is added, the sulphide of sodium is formed :^ —

KaHS + XaOH=Ka^ + HA

There are a number of poly sulphides of sodium varying in composi-
tion from Xa^Sa to NaaS*, Tliese are prepared by fusing sulphur

THE ALKALI METALS 323

with sodium carbonate. When treated with an acid they liberate
hydrogen sulphide and free sulphur.

Na2S3 + 2HCl = 2NaCl + H2S + 2S.

Sodium Sulphite, SaJSOs. 7 HsO.-r— When sulphur dioxide is con-
ducted into a solution of sodium hydroxide, sodium sulphite is
formed ; —

2 NaOH + SO2 = Na^SO, + H,0.

Sulphurous acid also forms the acid salt NaHSOg. Sodium sulphite
is oxidized to some extent to the sulphate by the oxygen of the air.
It is hydrolytically dissociated by water, as is shown by the alkaline
reaction of its aqueous solution : —

Na^Og + HjO = Na, OH + Na, HSO3.

The hydroxy 1 ion gives its characteristic alkaline reaction.

Sodium Sulphate, ira^04 . IOH2O. — Sodium sulphate, called from
its discoverer, Glauber's salt, exists in certain mineral waters as those
of Carlsbad, and occurs as the mineral thenardite. It is formed in a
large number of reactions. When sulphuric acid is neutralized with
sodium hydroxide we have : —

2 NaOH -f H,S04 = 2 H,0 4- Na,S04.

When sodium chloride is treated with sulphuric acid, sodium sul-
phate is formed, not because sulphuric acid is as strong as hydro-
chloric, but because the latter is volatile : —

2 NaCl 4- H2SO4 = 2 HCl 4- Na,S04.

A similar reaction takes place with sodium nitrate, nitric acid being
volatile ; —

2 NaNOa 4- H2SO4 = 2 HNOj 4- Na,S04.

Similarly, with sodium carbonate : —

Na^CO., 4- ^SO, = H2O 4- CO, 4- Na2S04,

and, in general, whenever a salt of a volatile acid is treated with
sulphuric acid and the temperature raised, the volatile acid escapes
and the sulphate remains behind.

Sulphates can also be formed by double decomposition or metath-
esis. Thus, when a solution of sodium carbonate is treated with a
solution of a sulphate of a metal whose carbonate is insoluble, an
exchange of ions takes place. Take as an example sodium carbon-
ate and zinc sulphate: —

Na, Na, CO3 4- Zn, SO4 =Na, Na, SO4 4- ZnCO,.

824

PRINCIPLES OF INORGANIC CHEMISTRY

The zinc carbonate is insoluble and can be filtered off, the sodium
sulphate remaining in the solution in the ionic condition. When
the solution is evaporated the ions combine and sodium sulphate is
obtained. When sodium sulphate with ten molecules of water of
crystallization is exposed to the air, it loses part of its water at ordi-
nary temperatures. Such salts are termed efflorescent.

Sodium sulphate is used as a purgative. It is a stage, as we shall
soon see, in the manufacture of sodium carbonate. When mixed
with concentrated hydrochloric acid it forms a good refrigerating
agent.

Solutions of sodium sulphate in water present a number of points
of interest. The facts are these : If sodium sulphate is allowed to

crystallize from its
solution above 33**, the
anhydrous salt Na,S04
separates. The solu-
bility of the anhydrous
salt is shown in curve
1, Fig. 34, the solubility
decreasing with rise in
temperature. This an-
hydrous salt can exist
below 32% if there is
not more than 0.000001
milligram of the hy-
drated salt present.
The solubility of this
salt has been studied
considerably below 32®, and the results are shown in the dotted
extension of curve 1. If there is present even the smallest trace of
the salt with ten molecules of water, the anhydrous salt cannot exist
below 32^

The solubility of ordinary Glauber's salt with ten molecules of
water of crystallization, has been studied at different temperatures,
and the results are plotted in curve 2. Unlike the anhydrous salt
the solubility of the hydrated salt decreases with decrease in tem-
perature, and this very rapidly. At 32® the two curves intersect,
the point of intersection representing equilibrium between the
anhydrous salt, the salt with ten molecules of water of crystalliza-
tion, and the saturated solution.

If a supersaturated solution of Glauber's salt is cooled to 5®,
another solid phase separates, having the composition Na2S04. 7 HjO.

T«imp«rature

Fia. S4.

THE ALKALI METALS 825

The solubility of this salt has been studied and the results are
plotted in curve 3 ; its solubility is greater than that of the Glauber's
salt.

Acid Sodium Sulphate (JStiHBO^) and Sodium Pyrosulphate
(JShSiOj), — Acid sodium sulphate is formed by the action of sul-
phuric acid on salts of sodium with volatile acids : —

NaCl 4- H^04 = NaHS04 4- HCl,
NaNOs 4- H^04 = NaHS04 4- HNO,.

Also by the action of sulphuric acid on the neutral sulphate : —

NajiSO* + H,S04 = 2 NaHS04.

When carefully heated in a vacuum to 300^ it loses water and
forms the pyrosulphate : —

2 NaHSO* = HjO + Na^ A-

When the pyrosulphate is heated still higher, and especially when
in contact with the oxides of certain metals, the following decom-
position takes place : —

Na^ A = Na,S04 4- SO3.

Sulphur trioxide is a powerful reagent, attacking most substances
with which it comes in contact The acid sulphates are, therefore,
useful as reagents, especially in the fused condition.

Sodium Thiosulphate, Na^sOs.SHsO. — This salt is frequently
referred to as sodium hyposulphite, or in commerce simply as
"hypo." It is prepared by dissolving sulphur in a solution of
sodium sulphite : —

Na,SOs+S = Na^A.

Solutions of this salt dissolve silver chloride and bromide, and it
is, therefore, used to remove these substances from the photographic
plate after the plate has been exposed to the light. If the unchanged
portions of these salts were not removed, the plate would still be
sensitive to light, and a slight exposure would ruin the photograph.
It is known in photography as a "fixing" agent. As has already
been mentioned this salt is used to remove the last traces of chlorine
from fabric which has been bleached by chlorine. In this capacity
it is known as "antichlor." The products of this reaction are
hydrochloric acid, sulphuric acid, sodium chloride, and sodium
sulphate : —

NajS A + 4 CI, + 5 HjO = 7 HCl + NaCl + H^04 + NaHS04.

PRINCIPLES OF INORGANIC CHEMISTRY

Although bromine reacts with the thiosulpbate in a similar man-
ner, iodine behaves very differently. The thiosulphate is converted
by iodine into the tetrathionate, and sodinm iodide is formed ; —

2 Na^S A + 21 = Na,S,0, + 2 Nal.

A standard Bolntion of sodinm Ibiosulphate, containing a known
amonnt of the salt in a giveii volume of the solutiou, is used to
detennine the amount of free iodine present under any given condi-
tions. The iodine is transformed into sodium iodide, which is color-
less. It is, therefore, a very simple matter to determine when the
iodine is all transformed into the iodide.

Sodiiun Caxbonate, N&^0;j.lO H.0, — The salt with ten 'molecules
of water cryt^tallizes from a solution allowed to cool on the air. If
a hot, concentrated solution is allowed to cool in such a way as
to be protected from any sodium carlwrHiate which may Ije in the air,
the salt NajCOa-T IliO separates* There are two varieties of the
geptahydrate which have diifer«ut crystalline formsj and which differ
considerably in their solubility in water.

If the saturated solution of sodium carbonate is boiled to crystal-
lization, the monohydrate iNajCOg.HaO separates. The salt with
ten molecules of water passes into the salt with one molecule
of water at 34°, This is the transition piunt between these two
phases. The monohydrate is less soluble the higher the tempera-
ture, and in this respect is analogous to anhydrous sodium sulphate.

When any one of the hydrates of sodium carbonate is heated to a
sufficiently high temperature, it gives np water step by step, and
passes over into the anhydrous salt. ludeed, the salt with ten mole-
cules of water loses a part of its water at ordinary temperatures —
is effloresceiiL

Sodium carbonate was formerly obtained from the aahes of sear
plants, but on account of its great importance, especially in connec-
tion with the manufacture of glass and soap, methods have been
devised for manufacturing it

The Le Blanc M<sthod of preparing sodium carbtmate was used
almost exclusively until quite recently, when another method was
devised which bids fair to supplant it In the Le Blanc method the
sofUum chloride is converted into the sulphate by means of sulphuric
acid. The sulphate is reduced to the sulphide by means of highly
heated carbon. The third and last process is to heat sodium sulphidet^
with calcium carbonate, wlien, at a sufficiently high temperature,
calcium sulphide and sodium carbonate are formed* The reactions
expressing these three tranaformatious are : —

THE ALKALI METALS S27

I. 2 NaCl 4- H,S04 = 2 HCl 4- Na,S04,

II. Na,S04 + 4 C = 4 CO 4- Na^,

III. Naj^ 4- CaCOa = CaS 4- Na^COs.

It is not difficult to separate the sodium carbonate from the cal-
cium sulphide, since the latter is difficultly soluble in water, while
sodium carbonate is readily soluble.

The sodium carbonate thus obtained is impure and is known as
soda ash or crude soda. It is purified by crystallization from water.

The sulphur is obtained from the calcium sulphide by means of
carbon dioxide. When this is conducted into the moist, calcium
sulphide, calcium carbonate and hydrogen sulphide are formed. The
hydrogen sulphide is then burned by means of the oxygen of the air
and there is formed sulphur, or sulphur dioxide, which can readily
be transformed into the trioxide, and this with water into sulphuric
acid.

The Solvay or Ammonia Process is based upon the fact that acid
sodium carbonate is much less soluble in water than acid ammonium
carbonate. Ammonia and sodium chloride are dissolved in water,
and carbon dioxide passed into the mixture. Under these conditions
acid sodium carbonate, on account of its small solubility, separates
from the solution. The reactions may be represented thus : —

NH, 4- HjO 4- CO, = NH4HCO3,
NH4HCO3 4- NaCl = NH4CI 4- HNaCOs.

The acid sodium carbonate when heated forms the normal car-
bonate, carbon dioxide, and water : —

2 HNaCOs = H2O 4- CO, 4- NajCOs.

The carbon dioxide is conducted into more ammonia in the pres-
ence of sodium chloride, and the process is thus a continuous one.

Acid Sodium Carbonates, ITaHCOs and ira3H(C08)2, 2 H2O. —Primary
sodium carbonate, or acid sodium carbonate, or the " bicarbonate of
soda," is formed by the action of carbon dioxide on the normal
carbonate * ^— ~

NajCOa 4- CO, 4- H,0 = 2 NaHCOa.

It is also formed, as we have just seen, in the preparation of nor-
mal sodium carbonate by the ammonia process. When acid ammo-
nium carbonate, formed by the action of carbon dioxide on ammonia,
is treated with sodium chloride, acid sodium carbonate is formed : —

NH4HCO3 4- NaCI = NaHCOa 4- NH4CL

S28

PRINCIPLES OF mORGiJJIC CHEMISTRY

When primary sodium carbonate is boiled with water, it loses cap- 1
bon dioxidt! and dissolves as the normaJ carbonate. When the solu-
tion is evaporated quickly the sesquicarbouate Na^COa, NaHCOa. 2 H^O
separates.

Hydrolyeis of the Carbonates. — The aqueous solution of normal
sodium carbonate has a strongly alkaline reaction. Indeed, the pri-
mary or acid carbonate has a weakly alkaline reaction. This is duo
to the presence of hydroxy 1 ions in the aqueous solutions of the car-
bonates. The carbonates are salts of the very weak, carbonic acid,
and like all salts of weak acids are hydrolyzed to a greater or less
extent by water : —

Ka,CO« + HjO = Ka, OH + Na, HCOa-

Even the acid carbonate is hydrolyzed to a sufficient extent to show
an alkaline reaction.

All carbonates which are soluble in water are hydrolyzed to a
snffieient extent to show an alkaline reaction.

The Fhoiphates of Sodium. — Since phosphoric acid is tribaaic,
there are three sodium salts of this acid possible, and all are known.
The gecondary sodium phosphutef KaaHPOi, is by far the best known,
and is always meant when the term sodium p^hosphate is used with-
out qualification. It contains twelve molecules of water of crystalli-
sation. It is a beautifully crystalline compound, containing, as ordi-
narily formed, twelve molecules of water of crystallization. When
crystallised above 35^^ it comes down with seven molecules of water.

Its aqueous solution is slightly alkaline, due to the fact that
the phosphoric acid is a weak acid, and it is slightly hydrolyzed by
water : —

Ka^HPO^ -h HjO = k\ OH + Na, H,f 0^,

When carbon dioxide is conducted into a solution of disodium
phosphate, we have a liquid with both acid and alkaline reaction, —
it colors blue Htm us red, and red litmns blue. Such reactions af6
known as ampholenc.

When dismlium phosphate is treated with one equivalent of phoa-
phoric acid, the monomdimn phosphate HsKaPO* is formed: —

Na,HFO, + H,PO* = 2 Nall.PO^.

The monosodium phosphate crystallizes in two forms, each cou-
taining one molecule of water of crystallization.

If, on the other hand, one equivalent of sodium hydroxide is added

THE ALKALI METALS 329

to one equivalent of disodium phosphate, the trisodium phosphate is
formed : —

Na,HP04 + NaOH = Na8P04 + H,0.

The trisodium phosphate crystallizes with twelve molecules of
water, Na8P04 . 12 HjO. When this salt is dissolved in water the
solutiqn shows a strongly alkaline reaction, due to the marked hydrol-
ysis of this salt by water : —

NajPO^ + HjO = Na, OH -f Na, Na, HPO4.

When a primary phosphate is heated it loses water, forming a
metaphospJicUe : —

NaH^04 = HjO + NaPO^

When disodium phosphate is heated, a jn/ropfiosphate is formed : —

2 HNajPO* ^ H,0 4- Na4P207.

This crystallizes with ten molecules of water : —

Na«PA.10H,O.

The metaphosphate of sodium is used like borax in qualitative
analysis. It readily fuses, forming a clear liquid, and this liquid
dissolves the oxides of many metals, forming glass-like masses,
which have characteristic colors. By means of this reaction, small
quantities of many metals can be detected. The metaphosphate is
fused in a Bunsen burner, or a blowpipe, in the loop of a platinum
wire, and a small part of the substance to be analyzed is added to the
fused metaphosphate. The color of the bead when hot and when
cold is observed, and the metal thus identified.

Sodium Ammonium Phoftphate, NaNH4HP04.4H20. — The double
phosphate of sodium and ammonium is also of importance in analysis.
It is really a triple phosphate of sodium, ammonium, and hydrogen.
It is formed by bringing together, in solution, disodium phosphate
and ammonium chloride : —

Na,HP04 4- NH4CI = NaCl + NH4NaHP04.

When heated it decomposes into ammonia, water, and sodium
metaphosphate : —

NH4NaHP04 = NH, + H^O + NaPOj,.

This salt is known as microcoamic salt, and is generally used as
the means of preparing sodium metaphosphate. When a small piece of
microcosmic salt is heated in the loop of a platinum wire, the above

830 PRINCIPLES OF INORGANIC CHEMISTRY

decomposition takes place, and sodium metaphospbate results. This
dissolves metal oxides as already described.

Sodium Borate or Tetraborate, irasB407 . 10 Rfi. — Borax is not the
salt of normal boric acid, HgBOj, but of a polyboric acid derived from
the normal acid by loss of water. When four molecules of boric
acid lose five molecules of water, the acid from which borax is
formed results: —

4 H3BO, = 5 H,0 4- H,BA.

Borax, or sodium tetraborate, is found in a few localities in
certain lakes, notably in the western part of the United States and
in Asia. It is formed when boric acid is neutralized with sodium
carbonate : —

4 H3BO3 4- Na^COa = Na^B A 4- 6 H,0 4- CO^

When borax crystallizes from aqueous solution below 56^, it
comes down with ten molecules of water of crystallization —
NajBA . 10 HjO. AVhen it crystallizes from a solution above 56"^,
it contains only five molecules of water, Na2B407. 5 HjO. This tem-
|)erature (56^) is then the transition temperature between the two
hydrates of borax. The former, from its prismatic form, is known
as "prismatic" borax, the latter as "octahedral" borax. When
borax is heated to a higher temperature, it gives off all of its water
and then fuses to a clear liquid. This liquid, as we have already
stated, has the power to dissolve oxides of certain metals, and form
with them vitreous masses with characteristic colors. The " borax
bead," like the microcosmic bead, is, therefore, very useful in quali-
tative chemistry.

On account of this property of dissolving metal oxides borax is
frequently used as a flux. It is used for the same reason to clean
two metal surfaces which it is desirable to solder together. These
usually become covered with a layer of oxide when the metal is
heated, and the solder, or easily fusible alloy, will not adhere to the
surfaces while the oxide is present. When a little borax is added, it
removes at the elevated temperature the oxides already formed, and
at the same time protects the hot metal surfaces from the oxygen of
the air. Borax is used only when hard solder, which is an alloy of
zinc, copper, and silver, is employed. When ordinary soft solder is
used, it is better to moisten the surfaces with a solution of zinc
chloride to which a little hydrochloric acid has been added.

Sodium Silicate, Wa^SiOs. — It has already been mentioned that the
silicates of the alkalies are soluble in water. When finely powdered

THE ALKALI METALS 831

sand is fused with sodium hydroxide or "sodium carbouate, sodium
silicate is formed : —

SiO, + 2 NaOH = Na^SiOj 4- HA

Sodium silicate is known as sodium vxUerglaas. It is readily
soluble in water, depositing a vitreous coating when the water is
allowed to evaporate.

Sodium Pyroantimoniate, NajHsSbA* — This salt is of special
interest, because it is one of the few salts of sodium which is diffi-
cultly soluble in water. It requires about 350 parts of water to
dissolve this salt. When a solution of potassium pyroantimoniate
is added to a solution of a sodium salt, sodium pyroantimoniate is
precipitated: —

K,H,SbA + 2 NaCl = 2 KCl + Na^H^Sb A-

The sodium salt of sulphantimonic acid — NajSbSi — known as
ScJdippe's salty has already been referred to.

Sodium Acetate, CHsCOONa.SHA — One or two compounds of
sodium with organic acids will be referred to. Sodium acetate is
formed by neutralizing acetic acid with sodium hydroxide, and
evaporating the solution to crystallization. Sodium acetate is very
soluble in water, one part of salt dissolving at 50** in 1.7 parts of
water.

Sodium acetate is used extensively in analysis. When an acid is
added to sodium acetate, i.e, when free hydrogen ions are added, acetic
acid is formed. Acetic acid, however, being a very weak acid, is only
slightly dissociated in aqueous solution. When free hydrogen ions
are added to sodium acetate the following reaction takes place : —

H, CI 4- Na, CHaCOO = Na, CI -f CH3COOP.

The hydrogen ions are thus removed from the field of action, and
are prevented from dissolving substances which are soluble in water
containing a large number of these ions.

Sodinm Cyanide, NaCN. — The sodium compound of hydrocyanic
acid presents interesting solubility relations. When it crystallizes
from a solution whose temperature is above 33®, the anhydrous salt
separates. Below this temperature the salt crystallizes with one or
more molecules of water, depending upon conditions.

Spectrum of Sodium. — Sodium is readily recognized by means of
the color which it imparts to the flame. If a platinum wire contain-
ing a sodium salt is introduced into the colorless flame of a Bunsen
burner, the flame becomes immediately colored bright yellow. If

882 PRINCIPLES OF INORGANIC CHEMISTRY

such a flame is examined by means of the spectroscope it will be
found to contain an intensely bright line in the yellow. This is
known as the sodium line, or in spectroscopy as the D line.

The almost universal presence of sodium is shown by means of
the si^ectroscope. The D line appears under almost all conditions^
unless very special precautions are taken to exclude it. Whenever
any flame is examined by the spectroscope under ordinary conditions
the I) line appears, and is used as a standlurd with which to compare
other spectroscopic lines.

CHAPTER XXVIII

POTASSIUM (At. Wt.= 39.15)

Oocnrrence and Preparation. — Potassium like sodium does not
occur in nature in the free condition, and for the same reason, viz.
the great chemical activity of the substance. Although the salts of
potassium, like those of sodium, are soluble in water, they are not
carried down to the sea in anything like the same relative quantity.

This is due to the fact that plants have the power of taking up
potassium ions in large quantities and building them up in their
tissues. They exercise selective absorption for potassium ions,
allowing the sodium to be carried on by the same waters from which
they remove the potassium. When such plants are burned the
potassium salts remain behind in the ashes. Potassium occurs
in large quantity in the ashes of certain kinds of wood, as is well
known, and can be easily leached out of the ashes by means of
water which is allowed to trickle through them. The lye thus
obtained contains a large amount of potassium hydroxide.

Potassium also occurs in many of the more common rocks and
minerals in the form of silicates. Ordinary feldspar is the double
silicate of potassium and aluminium. When these are decom-
posed by weathering the potassium salts are set free and become
available for plants.

Potassium salts also occur in the great salt beds, especially in
those of Stassfurt, in Germany. The chloride is known as sylvite, the
nitrate as saltpetre, and the sulphate, when in combination with other
metallic sulphates, as cdums, CamaUite contains also magnesium.

The preparation of the element potassium is of the same histori-
cal interest as the preparation of sodium. Potassium hydroxide
like sodium hydroxide was regarded as elementary until Sir
Humphry Davy, in 1807, electrolyzed fused protassium hydroxide.
This compound was decomposed by the current, yielding metallic
potassium at the cathode and oxygen at the anode. The metal rose
• to the surface of the fused hydroxide, and took fire spontaneously on
coming in contact with the air.

S;U PRIXCTrLE? OF TXOEGAXIC CHEMISTRY

Potassium w«s nrvi jiJWiMVNi by reducing potassium carbonate
*itho*rhon:- ^^-„^^3o = 2K + 3CO.

PotR.<«»ium WR8 prii^fiM^ l*^?r by reducing the sulphide or
hydroxide with higli'y hiiatieiii metals, such as magnesium, alumin-
ium, iron, etc. : — v- - ^^ t- o . t-

rKi^J^2>Ig=2K4-2MgO-fH^

All of th««* ww:>w%5* 5«iyt^ now been abandoned in favor of the
<»lectwlvti<-. XjiMJi^Uu jv«*»ium is now prepared by electrolyzing
tlw' <^blAndc^. Ar KyArAX^A^ The high fusing-point of the chloride
i$ obv^^tioimhte, >*>»^ -*« nV«» high temperatures metallic potassium
ikh:s oi. Tv^ttissiiiu. v^Wl/^Jt^&^ f\>rming the subchloride. To lower the
♦;<^mni*r»Nw ni "^'hwiV; 5fv»fly^um chloride can be kept in the molten
,<yw»i?f>M». '1 ^ MiX'^ ^^'^ vtUoium chloride.

)if^l{««i4tf|«; v|^ I^MMtiMU — Potassium is characterized by its great
<»V«»>w. ft^*J^^'^^•. ^i^ ^^^ "ior® active than sodium. When a
^?nwl t*^^V" <v ^^ittf>;^i.J^^ i* thrown upon water it decomposes it in
fhf >*»t»»M »w»»»i»Vi. -W^ ^Hii;um» yielding potassium hydroxide and set-

J«,V>4-2K = 2KOH + H^

^^v' ^Vv»i* i'-* ^<^ v\w*> Kxf potassium is, however, so vigorous that
,v\v» ^'>^ ^"W uic^ U *U\>wtHl to move around over the surface of
iV >* <-v^* ciKs^ii hv\U b J^»^^erated to ignite the hydrogen.

^\sxvw-:u vwuiNucH with the oxygen of the air with the greatest
.^>^i-;vvv>^ .Cixi^ irhcivtVrv% cannot be kept in contact with the air.
\ Ax vN\a;i:u II ..s i>ix\s*?rv\\l under petroleum.

vS; .ix\v»*;ui v»t us v.vmbiuiuic so readily with oxygen it is an
,\xv,.v;*. iwiiuiJij; ^^oat, s^ettiiig metals such as aluminium free from

.•VN,uw..N.rt <iiH*< 'i^^ coml>iH^ H'ith dry oxygen, but combines with
ths' >;i\\*u\Ni ^iK^»i ii ^^^' mert^t traiH^ of moisture is admitted. This
»\ .vJUN^.'i ov4;a^>!i» ^»i' th^ ^vuiderful influence exerted by water on

i\s,i,H\ii::ii iN ^'t ^^^ "^ami* v;enenil appearance as sodium. A fresh
NUiia.v VwN * Isii^i-'^J^'^'i color, but isH very quickly tarnished by con-
1.4, i \\i;U '.iu^ kn. It uiolts at tV2^.5, and lx)ils at 670^

\'i\s^ wXkw^^-^kv \\ci>;ht of motallio }H>tassium dissolved in nier-
xu»\, ♦> .lv;vMn'Mi\l l»\ the loworin^r of the freezinp:-j)oiut of the
wu.\*. \. 'J» i»Ku*iu\illN identical with the atomic weight.

POTASSIUM 835

Potassium Hydride, XH. — When potassium is heated in an atmos-
phere of hydrogen, the two elements combine, forming the hydride
of potassium. The combination takes place rapidly at 350**. If
potassium hydride is heated a little above 400**, it dissociates into
potassium and hydrogen.

Potassium Peroxide, XOf — The only compound which potassium
is known to form with oxygen is the dioxide. It is obtained by
heating potassium in a current of dry oxygen. It is an orange-
colored powder, melting at 280**. In contact with water it forms
potassium hydroxide and hydrogen dioxide; —

2 KO,4-2 H20 = 2 KOH -f HA 4- 0,.

Potassium Hydroxide, KOH. — When metallic potassium is thrown
upon water, potassium hydi'oxide is formed, as we have seen.

It is also formed when the peroxide is dissolved in water.

It can be readily formed by treating potassium carbonate with
the hydroxide of a metal whose carbonate is insoluble, say calcium
hydroxide : —

KjCOa + Ca(OH), = CaCOs + 2 KOH.

The calcium carbonate being insoluble, is iiltered off, while potassium
hydroxide remains in solution.

It is readily obtained by treating a solution of potassium sulphate
with a solution of the hydroxide of a metal whose sulphate is insolu-
ble, such as barium hydroxide : — •

K2SO4 + Ba(0H)8 = BaS04 -f 2 KOH.

None of these methods are used extensively where it is desired
to prepare potassium hydroxide upon the large scale. They have all
been supplanted by the electrolytic method.

When a concentrated aqueous solution of potassium chloride is
electrohjzedy the potassium ions move with the current to the cathode,
but do not separate upon it. They find around the cathode a few
hydrogen ions from the slightly dissociated water, and these, holding
their charge less firmly than the potassium ions, give it up to the
cathode and escape as hydrogen gas. The hydroxyl ions from the
dissociated water, corresponding to the hydrogen ions which have
escaped, remain in solution around the cathode, and with the potas-
sium ions form potassium hydroxide.

The chlorine anions, having moved over to the anode, holding
their charge less firmly than the hydroxyl ions from the disso-
ciated water, give it up to the anode and escape as ordinary gaseous

886

PRINCIPLES OF INORGANIC CHEMISTRY

chlorine. The reaction which takes place as the result of the deoom-
posing action of the current may be represented thus : —

2 K , 2 CI 4- 2 H, 2 0 H = 2 k , 2 d H 4- H , + Cl^

If mercury is used as the cathode, the potassium dissolves in the
mercury and forms potassium amalgam, which, when treated with
water, forms potassium hydroxide and liberates hydrogen.

Potassium hydroxide dissolves very readily in water, forming
caustic potash, and the solution is one of the strongest bases known.
It dissociates completely into potassium and hydroxyl ions, —

KOH = k, OH,

and this at no very great dilution. When a dilution of about 1000
litres is reached, the potassium hydroxide is completely dissociated
into its ions. The great dissociation of this base is shown by the
following very high conductivities, which are given for several
dilutions : —

r

Mr (18°)

a

1

10

100
M» 600

171.8
198.6
212.0
214.0

80.3 per cent

92.8 per cent

99.1 percent

100.0 per cent

Potassium hydroxide, on account of its solubility, readily pre-
cipitates the hydroxides of the heavy metals from aqueous solutions
of their salts : —

Ag, NO3 4- K, OH = AgOH -f K, NO,,

Cd, CI, CI -f OH, k 4- 0~H, k = Cd(OH), + K, 01 + K, CI,

¥e] CI, (II, Cl + Na, o"H-|-Na, 0~H +Na, 0~H =

re(0H)8 4- CI, Na 4- CI, Na 4- CI, N a

Potassium hydroxide, on account of its strongly basic nature, acts
vigorously upon organic matter, decomposing it into simpler sub-
stances. When brought in contact with the skin it disintegrates the
organic matter and partly dissolves it.

The white, hard, solid potassium hydroxide melts at a red heat,
and when in contact with the air at ordinary temperatures, absorbs
carbon dioxide from it and forms the carbonate.* Potassium hydrox-

1 Silver hydroxide, however, breaks down into Bil?er oxide and water.

POTASSIUM 837

ide which has stood in contact with the air for any length of time,
is, therefore, always contaminated with potassium carbonate. To
free it from the carbonate it is dissolved in alcohol in which the car-
bonate is insoluble. When the alcoholic solution of the hydroxide
is filtered to remove the carbonate, and evaporated away from all
traces of carbon dioxide, the pure hydroxide is obtained and is known
as potassium hydroxide by alcohol.

Compoundfl of Potasiinm with the Halogens. — Potassium com-
bines with all of the four halogens, forming beautifully crystalline
and stable compounds.

Potassinm Chloride, KCl, occurs in nature in combination with
magnesium chloride as the mineral camalUte, RMgClg . 6 H^O,
and in other combinations. When a hot solution of this salt
crystallizes, the double salt decomposes, potassium chloride sepa-
rating out. Potassium chloride when it occurs in the pure con-
dition is known as sylvine. It is a beautifully white substance,
crystallizing in cubes, which are readily dissolved by water. It
can be easily purified by crystallization, since it is far more
soluble in hot water than in cold. At 0^ one part of water dis-
solves 0.28 part of the salt, while at 100® one part of water
dissolves 57 parts of the salt. Potassium chloride is a type of a salt
of a strong acid with a strong base. We have seen that a strong
acid and a strong base mean those which are greatly dissociated.

The compound formed by the union of the cation of the base K, with

the anion of the acid CI, is among the most strongly dissociated
substances known. At a dilution of 400 to 500 litres potassium chlo-

ride is completely dissociated into its ions K and CI. A dilute
solution of potassium chloride is, therefore, a solution of potassium
and chlorine ions and nothing else, there being no molecules in the
solution. All the properties of such solutions, both chemical and
physical, are the properties of chlorine ions and potassium ions, since
these only are present. A study of the chemical and physical prop-
erties of such a solution proves this to be the fact. Such a solution
shows in general the chemical reactions of potassium ions, and of
chlorine ions, one of the most characteristic being the union with
the silver ion forming silver chloride.

The physical properties of such solutions are always additive, as we
say, i.e. tlie sum of two constants, one depending upon the cation, the
other upon the anion. In this class belong the density, power to re-
fract light, surface-tension, heat expansion, lowering of freezing-point,
lowering of vapor-tension, and in general all physical properties.

838

PRU^XIPLES OF INORGANIC CHEMISTRY

Without the aid of the theory of electrolytic dissociation it has
been impossible to interpret aueh facts; they can not only be inter-
pre ted by means of this theory, but are a necessary con sequence of it.

Potassium chloride melta at 770^ and passes into vapor at a white
heat. When potasBiura chloride is fused with metallic potassinm
the subchforifle K.jUl is formed.

Potaasium Bromide, KBr, is formed by the action of bromine on
caustic potash : —

6 KOH + 6 Br = 5 KEr + KBrOj, + 3 HA

also by the action of caustic i>otash upon ferroiia bromide :^ —
FeBrj + 2 KOH = Fe (OH), + 2 KBr.

It is a beautifully crystalline solid, melting at 71o°, and very soluble
in water, one part dissolving in one part of water at 100°,

Its aqueous solution is completely dissociated at moderate dilu-
tion, yielding potassium and bromine ions. This salt furnishes im
with one of the most convenient means of obtaining bromine ions ia
solution at any desired conceutmtian,

Potasiinm Iodide, KI, is formed by treating ferrous iodide with i
caustic potaaL The ferrous iodide is prepared by the action of
iodine on iron in the presence of water. Ferrous iodide has the
power to take np more iodine and form F'eala, When this compound
is treated with potassium hydroxide or potassium carlx>nate, ferrous
hydroxide is precipitated and potassium iodide reinains iu solution.
It crystallizes froju the solution on evaporation, in the form of beau-
tifully white cubes. These melt at 623"^, and are more soluble in
water than evt^u potassium bromide. One part of water at 0* dia-
solves 1,27 parts of potassium iodide.

The aqueous solution of potassium iodide, like the bromide and '
chloride, is completely dissociated at moderate dilution into potaa-
sium and iodine ions* By dissolving this compound iu water we can
easily prepare a concentrated solution of iodine ions,

PotasalEtn Fluoride, KF, 2 H^O, is formed by the aetion of hydro- i
fluoric aeid on potassivnu hydroxide. Like the remaining halogen J
comfjounds of potassium it is a white solid; unlike them, howeverpJ
it crystallizes with two molecules of water. When potassium fiuo-l
ride ia treated with an equivalent of hydrofluoric acid it forms tha]
compound KHFj, which, together with other facts, points to the diba* j
sic nature of hydrofluoric acid.

Hydrofluoric acid, then, probably has the composition HaF^ |
potassium fluoride the composition K^F^ and the acid salt, KHF»

POTASSIUM 339

FotaBsium Chlorate, XCIO3. — Potassium combines with the oxygen
acids of chlorine, forming well-defined salts. A few of these are of
sufficient importance to merit special consideration. Potassium
chlorate is prepared, as we have seen, by the action of chlorine on
caustic potash ; —

6 KOH + 3 CI2 = 5 KCl + KCIO3 + 3 UjO.

It is separated from the chloride by its solubility in water being
much less than that of the chloride. It is also formed by the action
of potassium chloride on calcium chlorate. When chlorine obtained
electrolytically is conducted into lime calcium hypochlorite is formed.
When a solution of calcium hypochlorite is boiled it passes into the
chlorate and chloride : —

3 Ca(0Cl)2 = CaCClOa), + 2 CaClj.

When the solution of calcium chlorate is treated with a solution
of potassium chloride the following reaction takes place : —

CaCClOa), + 2 KCl = CaClj + 2 KClOa.

Potassium chlorate, on account of its smaller solubility, is formed
and can be readily obtained from the solution in beautiful white
plates.

When potassium chlorate is heated it decomposes in the sense of
the following equation : —

2 KClOs = KCl + KCIO4 + 0»

potassium chloride and i)erchlorate being formed. When the per-
chlorate is heated to a still higher temperature it breaks down into
the chloride and oxygen. Potassium chlorate is useful chiefly
because of the large amount of oxygen which it contains and which
it can readily give up. It is, therefore, an excellent oxidizing agent,
and as such is useful in chemistry. Its oxidizing power is due to the

ease with which the chloric ion CIO3 passes into the chlorine ion CI,
liberating three oxygen atoms. This decomposition takes place with
the evolution of a large amount of heat, which explains the violent
nature of such reactions as the following : —

When potassium chlorate is powdered with a small piece of
sulphur, an explosion occurs, which is violent if an appreciable
quantity of sulphur is used. A violent explosion results if potassium
chlorate is brought together with phosphorus. With antimony
sulphide an explosive mixture is also formed.

Potassium chlorate is extensively used in the preparation of
matches. The so-called safety matches are made of a mixture of

840

PRINCIPLES OF INORGANIC ClIKMISTKr

potassium chlorate and sulphide of antimony. When these are
rubbed upon a surface covered with red phosphorus, a miniature
explosion results and the whole mass is ignited. When rubbed
upon an ordinary object such matches do not take fire*

When treated with concentrated sulphuric acid potassium chlorate
liberates oxides of chlorine, which are unstable and frequently explode
with great violence. The instability of these coin pounds accounts
for the reaction when a mixture of sugar and j>otassium chlorate m
treated with snlphunc acid* They break down, and in doing so
liberate enough heat to ignite the cane-sugar,

Potasfiinm Perchlorate, £010^, is prepared, as we have already
seen, by heating potassium chlorate carefully until the molten mass
has resolidified and the first evolution of oxygen has practicaily
ceased. The salt is of interest because it contains more oxygen
than iwtassium chlorate, and is still a much more stable compound,
giving up its oxygen only at a considerably higher temperature.
This undoubtedly has to do in some way with the manner in which
the oxygen is combined in the compound — with the constitution
of the compound. It cannot be explained if we regard the molecule
as simply a material system, and disregard the way in which the
system is made up. It may have to do also with the arrangement
of the atoms in space ^ — with the stereochemistry of the molecule.
By adding one oxygen atom to potassium chlorate the geometrical
configuration of the molecule may be so changed as to form a more
stable system. In reference to these matters we at present, however,
know nothing, and but little is gained by speculation. Potassium
perchlorate is of importance on account of its small solubility in
water, one gram of the salt dissolving in 143 grams of water at 0*
If alcohol is added to the water the solubility of potassium perch lo»
rate is still further greatly diminished. This is one of the few
difficultly soluble salts which |)otassium forms with acids, and is,
therefore, used to defect the presence of p&tcufslum in a mhtlmi of
potaamum ions.

f otaiiium Hydrazoate, KH^^, and Potassium Amide, ENH^. — The
potassium salt of hydrazoic or triazoic acid is formed when a solu-
tion of potassium hydroxide is neutralized with liydmzoic acid. It
resembles in its appearance potassium chloride, and in its properties
the sodium salt of this acid.

Potassium amide is formed by conducting carefully dried am*
monia over metallic potassium. One hydrogen atom of the ammonia
is set free and potassium takes its place : —

POTASSIUM 841

When potassium amide is treated with nitrous oxide, potassium
hydrazoate is formed : —

2 KNHs + N,0 = KOH + NH, + KN,.

This reaction is strictly analogous to that which we have already
studied under the element sodium.

Fotassinm Nitrate, KSO^, — Potassium nitrate or saltpetre is one
of the most important salts of potassium. It is very soluble in water
and, therefore, does not occur in the solid state in any considerable
quantity in regions where there is an abundant rainfall. It is, how-
ever, leached out of the soil in certain regions in the East Indies
during the rainy season, and deposited when the rain has ceased.
This is known as the " India crude saltpetre."

Potassium nitrate is formed in large quantity in the saltpetreplan-
tations. Refuse animal matter which contains nitrogen is mixed
with potassium or calcium carbonate, or with earth or wood-ashes
containing these substances, and exposed to the action of the " nitri-
fying ferment " or " saltpetre bacteria " in the soil. The oxygen of
the air, through the agency of these bacteria, oxidizes the ammonia
formed from the decomposing organic matter, to nitric acid, which
then combines with potassium or calcium hydroxide or carbonate and
forms the corresponding nitrate. After the action has" continued for
several years the whole mass is treated with water, which dissolves
all of the nitrates, including in addition to potassium especially those
of calcium and magnesium. The mixture of nitrates is treated with
the product of the leaching of wood-ashes, Le. with a solution of
potassium carbonate. Calcium and magnesium carbonates are pre-
cipitated, and potassium nitrate remains behind in solution.

Frequently, nitrates are formed around stables and other places
where organic matter is undergoing decomposition, and in rainless
regions this forms incrustations which are dissolved in water and
converted into saltpetre.

Potassium nitrate is made to-day chiefly from sodium nitrate or
Chili saltpetre. AVhen a solution of sodium nitrate is mixed with a
solution of potassium chloride, the following reaction takes place : —

Na, NO3 + K, CI = NaCl + K, NO,.

The sodium chloride is deposited at higher temperatures. When
the solution is allowed to cool down, potassium nitrate is deposited.
The reason why this reaction takes place is found in the relative
solubilities of the four salts, potassium and sodium chlorides and
potassium and sodium nitrates.

342

PRINCIPLES OF INORGANIC CHEMISTRY

These relations are shown in Fig. 35. At the higher temperatures
the solubility of sodium chloride is less than that of potassium chlo-
ride, and much less than that of potassium nitrate. It, therefore,
separates when the solution is concentrated at the higher tempera-
ture. Further, the solubility of sodium chloride is as great at 0® as
at 100°. When a solution saturated with sodium chloride at 100®
is allowed to cool down to 0®, it will, therefore, be only saturated at
the lower temperature and will not dejwsit any of the salt. On
the other hand, potassium nitrate is many times as soluble at 100® as

NOt

NaNOt

100 c.

at 0®. WTien a solution of potassium nitrate which is far from satu-
rated at 100® is cooled down it will, therefore, deposit crystals long
before zero degrees is reached.

These solubility curves explain more at a glance concerning the
reason why the above transformation takes place, than could be done
by pages of description. When the solution is evaporated sodium
chloride separates from it while hot. This is removed and the solu-
tion allowed to cool, when potassium nitrate crystallizes out. This
process is repeated a few times when the transformation indicated by
the above equation is practically complete.

POTASSIUM 343

Potassium nitrate readily gives up a part of its oxygen and is,
therefore, an excellent oxidizing agent. It is, therefore, used where
rapid oxidization is desired, such as in fireworks and especially in
gunpowder. Gunpowder is a mixture of potassium nitrate, sulphur,
and carbon, in such proportions as to secure complete combustion.
The nitrate gives off oxygen, which combines with the carbon forming
carbon dioxide; the nitrogen escapes as such, and the potassium
remains behind in the form of sulphide or sulphate. The equation
which is usually written to express the decomposition of gunpowder

^^''~ 2KN03-f3C + S = KjS-f3CO, + N2.

This is an idealized equation, the reaction which takes place
being far more complex than it would indicate. In addition to the
above products, when gunpowder decomposes there are formed potas-
sium sulphate, potassium sulphide, potassium carbonate, and carbon
monoxide.

Gunpowder is prepared by mixing the three constituents in the
following proportions : —

KNO3 75 per cent

S 12 per cent

C 13 per cent

This corresponds almost exactly to three molecules of saltpetre,
three atoms of carbon, and one atom of sulphur, and is the chief
reason for writing the above very simple equation.

The reactions which take place are, as already stated, far more
complicated.

When gunpowder is ignited, the gases liberated occupy several
hundred times the volume of the powder; or if they are forced to
occupy the same volume as the original powder, the pressure exerted
is several hundred atmospheres. This is the principle made use of
in employing explosives to drive missiles with a high velocity. The
gunpowder is exploded in a metal tube closed on all sides and open
at one end. The ball is placed tightly upon the powder, so that
when the latter explodes the gases are liberated in practically a
closed space. An enormous pressure is thus generated, which drives
the ball out of the end of the gun with a high velocity. In calcu-
lating the force produced by an explosion of gunpowder, we should
always take into account the further fact that the gases are greatly
heated by the heat energy produced as the result of the reaction,
and, therefore, in approximate accordance with the law of Gay-
Lussac, exert a still greater pressure.

344 PRINCIPLES OF INORGANIC CHEMISTRY

Potassium Nitrite, KSOi.Kfi. — When potassium nitrate is care-
fully heated it melts at about 338®. When heated higher it loses
oxygen and forms potassium nitrite : —

2KN08 = 2KN08 + Oj.

This is the method by which oxygen was first prepared by its
discoverer, the great Swede, Scheele. Potassium nitrite is best pre-
pared by heating the nitrate with a mild reducing agent, such as
metallic lead:- rNO, + Pb = KNO, + PbO.

Under these conditions there is less decomposition of the com-
pounds, and, altogether, the reaction is a much smoother one. It is
also formed by neutralizing a solution of nitrous acid with potas-
sium hydroxide. When potassium nitrite is treated with a strong
acid nitrous acid is liberated. It can, therefore, be used as a means
of preparing nitrous acid. The potassium salt of nitrous acid, like
the salts of weak acids in general, is hydrolyzed to some extent by
water: — 4. - + -

K, NO2 + H,0 =K, OH -f HNOjT.

Nitrous acid, HNOj, being a weak acid, is only slightly dissociated,
and, therefore, there are many more hydroxyl ions in the above solu-
tion than hydrogen ions, and the solution reacts alkaline.

Compounds of Potassium with Sulphur. — The compounds of
potassium with sulphur do not present many points of difference
from those of sodium. Potdssium hydrosidjyhide, KSH, iH,0, is
formed when hydrogen sulphide is conducted into a solution of
potassium hydroxide: —

KOH + H,S = HjO + KSH.

This being a salt of a weak acid, is hydrolyzed by water,
showing an alkaline reaction: —

K, SH + H,0 = K, Oll + H, HS.

Since hydrogen sulphide is a weak acid, there are only a few hydro-
gen ions in the solution, and the hydroxyl ions show basic or alkaline
reactions.

When an equivalent of potassium hydroxide is added to potas-
sium hydrosulphide, the normal sidphide KjS.oHjO is formed: —

KSH -f KOH = K,S -f H2O.

This is also strongly hydrolyzed by water : —

K^ + H,0 = K, 6h + HS, K.

POTASSIUM 345

When sulphur is added to a hot solution of potassium sulphide,
it dissolves and forms polysulphidea varying in composition from
■K2S3 to 1^305

Compounds of Potassium with Sulphur and Oxygen. — When sul-
phur dioxide is passed into a solution of potassium hydroxide or
carbonate, potassium sulphite^ KsSOs, is formed. If the gas is passed
through the solution until the solution will take up no more of it,
the acid sulphite KHSOg results.

The potassium salt of persulphuric acid, KSO4 or KjSjOa, is a
well-crystallized substance. It is obtained by electrolyzing acid
potassium sulphate, and separates from the solution on account of it
being difficultly soluble.

Potassium Sulphate, JifiO^, occurs in nature in combination with
magnesium sulphate and magnesium chloride as Jcainite. This
occurs in a number of the salt-beds, but especially in those at Stass-
furt and other places in Germany. Kainite has the composition
KjS04. MgS04. MgClj. 6 HjO. The double sulphate of potassium and
magnesium is treated with chloride of potassium, when magnesium
chloride and potassium sulphate result. This salt is used to pre-
pare potassium carbonate. It is also used in the preparation of the
double sulphate of potassium and aluminium, or ordinary potassium
alum. It is extensively used as kainite to enrich the soil in potas-
sium ions, which are so much needed by many plants. Potassium
sulphate is not very soluble in water, one part of water at 0° dis-
solving only 0.085 part of the salt

When normal potassium sulphate is treated with an equivalent
of sulphuric acid the acidy or primary sulphaJtey is formed : —

K,S04 -f H,S04 = 2 KHSO4.

While the normal sulphate of potassium is only slightly soluble
in water, the acid sulphate is very soluble, one part of water dis-
solving 0.33 parts of the salt at 0^ Acid potassium sulphate occurs
in nature in certain volcanic regions, as in those of Naples, as the
mineral misenite.

When acid potassium sulphate is heated it melts at 200®. When
heated in a vacuum to 300"* it passes over into the pyrosulphate : —

2 KHSO4 = H,0 -f K,S A.

When heated still higher it decomposes, giving water and sulphur

trioxide : — ^ KHSO4 = H,0 + K^04 + SO^

or, K AO7 = K^04 -f SO,.

a46 PRINCIPLES OF IXORGA^^C CHOnSTRT

The sulphur trioxide thus set free dissolves oxides of metals
converting them into sulphates^ decomposes insolable silicates
forming soluble compounds, and in general is a powerful reagent.
When it is desireii to remove insoluble substances adhering to plati-
num vessels, the best method is to partly fill the vessel with acid
potassium sulphate, and heat the mass until there is a copious evolu-
tion of sulphur trioxide.

Acid potassium sulphate in aqueous solution shows a strongly
acid reaction. This is due to the fact that sulphuric acid is a strong
acid and the second hydrogen ion begins to dissociate.

KHS04 = K,H,S0^

Acid potassium sulphate is of interest in connection with the
development of the conception of mass action. Heinrich Rose, who
pointed out the action of carbon dioxide on silicates over the surface
of the earth, also called attention to the following facts. When a
boiling solution of acid potassium sulphate of medium concentration
is crystallized, the crystals have the composition expressed by the
formula 3 K2SO4.H3SO4, a portion of the sulphuric acid having been
split off to combine with the water. If these crystals are redis-
solved in more water, and the solution evaporated to crystallization,
the neutral salt will separate, showing a further splitting off of sul-
phuric acid due to the mass action of the water.

We can now understand how this reaction takes place. In an
aqueous solution of acid potassium sulphate there are both potassium

ions and sulphuric ions SO4. When the conditions of concentration
are properly established, these ions combine directly and form the
neutral sulphate.

Potassium Carbide, KsC^. — Metallic potassium acts directly upon
acetylene, forming the carbide of potassium : —

C,H,+ 2K=C,K, + Hj^

In the presence of water this breaks down, yielding potassium
hydroxide and acetylene.

Potassium Carbonate, KgCOa. — Potassium carbonate was obtained
for a long time mainly from the ashes of certain plants. When the
plants were burned the potassium remained behind in the form of
the carbonate. This was obtained by leaching the ashes with water
and evaporatin<2^ the solution, when the impure carbonate crystal-
lized out. The impurities are in the main potassium sulphate and
chloride and salts of sodium, -all of which are less soluble than
potassium carbonate. The carbonate is purified by means of the

POTASSIUM 847

difference in solubility between the impurities and the salt in ques-
tion. It dissolves, and for the most part leaves the impurities
behind. This impure mass is known as potash, the purified car-
bonate as jmrified potash.

Potassium carbonate is also obtained from the residues of the
beet-sugar industry.

Potassium carbonate can also be prepared by the action of mag-
nesium carbonate and carbon dioxide under pressure on potassium
chloride. There is formed the double carbonate of potassium and
magnesium, KHCO3, MgC08.4H,0, which, when decomposed with
water at a high temperature, yields potassium carbonate in solution.

Potassium carbonate is also prepared by the electrolysis of
potassium chloride. Around the cathode in a concentrated solution
of potassium chloride, potassium hydroxide is formed. Carbon
dioxide is conducted into this solution, and potassium carbonate is
formed. Potassium carbonate is remarkably soluble in water, one
part of water at O"" dissolving 0.83 parts of the salt, and the solu-
bility increases rapidly with rise in temperature.

Potassium carbonate takes up water from the air, cr is deliques-
cent. From the cold, aqueous solution a salt separates, containing
three molecules of water to two of potassium carbonate, 2 K2CO3. 3 HjO.
Potassium carbonate melts a little above 1000''.

The aqueous solution of potassium- carbonate shows a strong
alkaline reaction. This is due to the hydrolysis of the salt of the
weak carbonic acid : —

K, K, CO3 + H,0 = k, OH + HCba, k,

which takes place to a very considerable extent, yielding a large
number of hydroxyl ions, which react strongly alkaline.

Acid, or Primary Potassium Carbonate, KHCO3. — The acid salt is
formed by conducting carbon dioxide into the solution of the neutral
salt : — ^^QQ^ ^ QQ^ ^ H^Q ^ 2 KHCO3,

or, if we express this in terms of the ions : —

K, K, CO3 + CO, + H2O = k, HCO3 + HCO3, k. .

Dilute solutions of acid potassium carbonate react alkaline. At
first sight it may seem peculiar that an acid salt should show an
alkaline reaction, when there is still one acid hydrogen present in
the molecule. This was entirely ijnexplained until the theory of
electrolytic dissociation arose. Now we know that it is simply due
to the hydrolysis of the acid salt by water, forming hydroxyl ions: —

KHCO, + H,0 = k, OH. + HjCOa.

348 PRINCIPLES OF INORGANIC CHEMISTRY

When acid potassium carbonate is heated it decompoBes, forming
the normal carbonate, carbon dioxide, and water : —

2 KHCO, = K/::0, + H,0 + COjr.

If an aqueous solution is boiled the same decomposition takes
place. This continues in either ease until the carbon dioxide has
reached a certain pressure or density, or better expressed, concentra-
tion, which is proportional to pressure.

When, at any given temperature the pressure of the carbon diox-
ide has reached a certain value, equilibrium is established, and as
much carbon dioxide is absorbed in any given time as is set free. If
the pressure of the carbon dioxide is increased at this temperature,
the following reaction takes place : —

KjCOa + CO, + H,0 = 2 KHCO,.

This will be recognized to be exactly the reverse of the
above decomposition, so that we have here an excellent example
of a reversible reaction. If, when equilibrium is established,
pressure of the carbon dioxide is diminished, more of the acid
carbonate will decompose, until the equilibrium pressure is again
established.

If when equilibrium is established at any one temperature, the
temperature is not kept constant but varied, the equilibrium will be
destroyed ; and more of the acid carbonate will decompose, or will be
formed, until the pressure of the gas is such as to establish equili-
brium at the new temperature. When the carbon dioxide is allowed
to escape as fast as it is formed, its pressure is practically zero, and
nearly all of the acid carbonate can be transformed into neutral car-
bonate under these conditions.

Phosphates of Potassium. — Potassium combines with phosphorus
when the two elements are heated together, and forms the compound
KPa — potassium phosphorus.

Potassium, like sodium, forms three salts with phosphoric acid,
— the primary, KH2PO4; secondary, KJHPO4; and tertiary, KaPO^,
phosphates.

These phosphates do not call for any special comment. They are
all readily soluble in water, yielding potassium ions and phosphoric
acid ions, both of which are needed for the growth and seeding of
plants. They are, therefore, among the most valuable compounds
known as artificial fertilizers. When these compounds are heated,
they undergo decompositions which are similar to those suffered by

POTASSIUM 849

the corresponding sodium compounds; the secondary phosphate
yielding a pyrophosphate: —

2 K,HP04 = HjO + K4P A.

The primary phosphate gives off water and forms the metaphos-

phate : — KH,P04= H,0 + KPO,.

Silicates of Potassium. — Finely powdered sand, fused with potas-
sium hydroxide, or potassium carbonate, becomes soluble in water.
The thick, syrupy mass is supposed to be made up of a number of
compounds from which no one substance has thus far been isolated.
The syrup dissolves readily in water, and when the aqueous solution
is allowed to dry a vitreous mass is left behind. It is, therefore,
known as potassium water-glass. When inflammable objects are
covered with water-glass they become more or less fire-proof, since
the covering prevents access of oxygen to them unless they are sub-
jected to high temperatures. When a solution of the silicates of
potassium is treated with an acid, a heavy, white precipitate of
silicic acid is thrown down.

Potassium Silicoflaoride, E!s8iF«. — This salt is formed when hydro-
fluosilicic acid is added to a solution of a potassium salt : —

. KjS04 + HjSiFe = H^04 + KjSiF«.

This salt is comparatively insoluble in water and is, therefore, use-
ful in detecting the presence of potassiuuL The index of refraction
of the salt is almost exactly the same as that of water, and it is,
therefore, very difficult to see when suspended in water. Unless
precaution is taken to examine in different lights the liquid in
which the precipitate may be suspended, its presence can be easily
overlooked.

When potassium silicofluoride is treated with a strong alkali the
compound is decomposed into the alkaline fluoride and silicic

* KjSiFe -h 4 NaOH = 2 KF -|- 4 NaF -f Si(0H)4.

Potassium Pyroantimoniate, Xfihjdj, is formed by fusing anti-
monic acid with an excess of potassium hydroxide. When treated
with water it breaks down into potassium hydroxide and the salt
KsHsSbgO?. Although this salt is not very soluble in water, it is
far more soluble than the corresponding sodium salt. When a solu-
tion of the potassium salt is added to a solution of a sodium salt,
the insoluble sodium pyroantimoniate, as we have seen, is thrown
down.

3S0 PRINCIPLES OF INORGANIC CIIEiMISTUY

Pota^Biuiii Cyaaide, KCH* — Potassium cyanide is formed when
organic substances eonLaining uiti-ogen and carbon are heated with
metallic potassium. This reaction with potassium or sodium is fre-
quently made use of to detect the presence of nitrogen in organic
compounds. It is formed on the large scale by passing ammonia
over a mixture of carbon and jjotassium carbonate. It is also formed
in the blast^funuice where caibon and nitrogen are brought together
at very high temperatures. It is obtained in pum condition from a
salt which we shall study when we come to iron — potiisslum ferro-
eyanide^ K4Fe(CN)<^. When heated alone it gives potassium cyanide
— one-third of the cyanogen being lost. By heating th is with metallic
potassium iron is thrown out and potassium cyanide remains; —

K,Fe(CN)a -h 2 K = 6 KCN -h Fe.

Potassium cyanide is readily soluble in water. Being a salt of a
weak acid it is hydrolyzed by water, and its aqueous solution always
smells of hydrocyanic acid : —

KOK + HjO^K, Oil H- HON.

Hydrocyanic acid is very slightly dissociated by water, existing
in solution mainly as molecules* The solution has a certain vapor*
tension of hydrocyanic acid, and this is sufficient to produce a
detectable odor.

Potassium cyanide is a very powerful poison. It is a good reduc-
ing agent at an elevated t-emperature, taking ufr oxygen and pashting
over into the ayanitte^ KOCN; or snlphur, and passing over into the
mdpho€2iayiate^ luSCy. Witen potassium sulphocyanate is dissolved
in water a large amount of lieat is absorbed, and a marked refriger-
ating effect is prod viced. Potassium cyanide is extensively used in
connection with the extraction of gold from ores which are not very
rich, and in connection with the electrolytic deposition of many of
the metals.

Osalatas of Potasaiam. — Potassium forms three well-defined and
stable compounds with oxalic acid. These are the best-known saltfl
of oxalic acid, since they occur in abundance in certain plants. The
normal ojtxdale, K/_\O^.UJd^ can be easily prepared by neutralist-
ing oxalic acid completely with caustic potash. The ami oxahiU^
KHCsO^.^KsO, is prepared by treating the neutral salt with one
equivalent of oxalic acid. Like the neutral salt it occurs abun*
dantly in certain plants as the wood-sorrel, from which it can be
readily extracted.

Potassium forms still another oxalate with oxalic acid^ which ia

POTASSIUM 351

more acid than the acid oxalate. This is the tetroxalate of potassium,
KHC2O4 . H2Cjf04 . 2 HjO. In this compound there is one equivalent of
potassium to two molecules of oxalic acid, or the oxalic acid is just
one-fourth neutralized by potassium. It is readily prepared by bring-
ing either of the above oxalates together with the necessary amount of
oxalic acid, and allowing the salt to crystallize from the hot, concen-
trated solution. Although the molecule is fairly complex it is per-
fectly stable in the air, and is very useful in analytical chemistry,
since the salt can be weighed. It is now used in many analytical
operations where oxalic acid was formerly employed, since oxalic
acid gives off water when exposed to the air, and we are never quite
certain whether it possesses just two molecules of water of crystalli-
zation. When a standard solution is desired containing just so much
oxalic acid, the corresponding amount of tetroxalate is employed.

The tetroxalate is now extensively used in standardizing solu-
tions of potassium permanganate, and in many similar operations
where oxalic acid was formerly employed, for the reason indicated
above.

Detection of Potassium. — Potassium is most readily detected by
means of the flame reaction. When a potassium salt or a substance
containing potassium is introduced into the flame, the latter gives out
a reddish-violet light which is very characteristic. If the potas-
sium salt contains a sodium salt mixed with it, the intense yellow of
the sodium flame may entirely mask the far less intense color of the
potassium flame. In such cases it is necessary to cut out the sodium
light in order to see whether there is any of the potassium flame
present. This is accomplished by allowing the light to pass through
blue, cobalt glass, which cuts off all of the yellow, but allows the short
wave-lengths sent out by the potassium to pass through. When
the flame emitted by both sodium and potassium is examined
through cobalt glass, the sodium yellow is not seen at all, and the
potassium flame appears redder than when alone and examined with
the naked eye. The flame test, correctly made, is a very sensitive
means of detecting the presence of small quantities of potassium.

We have now studied sodium and potassium with some thorough-
ness. The remaining alkali metals, lithium, caesium and rubidium,
are comparatively rare substances and will be treated briefly.

CHAPTER XXIX

LrreiuM, RUBmniM, c^ssium, (ammonium)

LITHIUM (At Wt.^ 7,03)

Diicovery, Froparation, and Froperties. — Lithium waa discovered
as early as 1817^ but was not isolated until 1855, when Buosea
obtained it by electrolyzing the chloride. Lithium is widely distrib-
uted in nature but occurs only in relatively small quantities. Lith-
ium occurs as the silicate in the minerals lepidolite spodumtn^^
tourmaline, etc., as the phosphate combined with other phosphates in
amhhjgoitUe aiul triphfUfe. As alreatly stated Bunsen electrolyzed
the dhloride. The best results are obtained by dissolving the anhy*
drous chloride in some solvent which does not act chemicaUy upon
the metah Pyridine is such a solvent. A pyridine solution of the
chloride conducts the current very readily, and from thia solution
lithium readily separates upon the cathode.

Lithium resembles sodium very closely in its properties. It is
of nearly the same color, but is much harder than sodium andean be
drawn into wire. It quickly taniishes in contact with the air, due
to the combination with oxygen* It acts upon water, forming the
hydroxide and liberating hydrogen^ which, however, is not ignited,
and the metal is not melted by the heat set free. Lithium is the
lightest of all known metals, having a specific gravity of 0,59, It
melts at 180", and does not take fire in the air until about 200** is
reached.

Lithium forms the univalent cation Li, which acts like the univa-
lent sodium and potassium cations, combining with the anions of
acida, forming salts, A few of these substances will be considered,

Compotinds of Lithium, — The compounds of lithium closely re-
semble those of sodium and the other alkalies. In a few cases,
however, differences appear which are worthy of note. Unlike the
remaining alkalies lithium forms an oxide, Li^O, which dissolves only
slowly in water forming the hifdrnxide^ LiOH, The hydroxide is
disaoeiated by water in the same manner as the hydroxides of the
remaining alkalies : — + -^

LiOH = Li, OH,

352

LITHIUM, Ki'iiihr::: ' .^<' jf rsi:

There is a large numljfr cf :.; -.--.t; /-. \^- -- ^
therefore, reacts stroiij[^ly al x *..;.-

Lithium forms the hy:- fr '...':'. x -"-..'•

follQws : —

LiPr^H/' = J:V^:. .-'

The h'thhtm halifleif resei;.?/!- r:->< '.f •• »
deliquescent or absorbing moi-: ^:-r f: •.-.•'«,
the halidrs of the other alkiai.'r-. .-.••-.?
of alcohol and ether. Lithiti!.-! 5 :.; • •* • -

LioCI. Tho//(onV/<? of litliiiinj :^ : .'- -• :.
thus resembles more closely the n . . r. . -
sodium and the remaining alka':*--

The carbonate of lithium, L:J '* ».. ; •*- ' • ' -'-

est. Unlike the carbonates of ih«: ^ » <. -' ■ ''

water, one part of the salt req ;:;..•.:• f • -
water to dissolve it. In this re-ij^^-r -'. •■*--. -^
the alkaline earths. The hirarhocnf^. l^ •:'
than the carlxmate, and in this. a:»'?5-%. . '^^-

earths, as we shall see.

The tertiary or unrmnl lifhi^m ///''y^y*'- '* ' - '^
the same difference from the jJiO"?j>;.^*-r'. •. * -• -. -
insoluble in water ; one part of th'r :,;.•/*;,'..> •-
2i)()0 parts of water to diss^^lve it: hjy. .' t. . .
part of the phosphate requires al/>-t i-'O '/^'^ ' -
to dissolve it. This salt furnishes u\ »,*..'. fc •:>^'-
presence of a small amount of lithiurii. ''.-.>-: t ;
ble i)h()spliate, such as disodium pho';].:-%.V' • "'W.^r
of a lithium salt, lithium phosphate i'. ;^f *^.- ; '^>''

Xa,HPO^-f3LiCl = 2NaC: - J:'' - - .^'

The precipitation takes place in the pr^'b*^**^* '^
acid, which is set free. Here again lit;,. -'s- •*rv#»^i-. --
earth metals, wliich form phosphat#;H tl-^t *?* t^ir - -•

Lithium urate is soluble in watT, ajiC .* x:^** '>«^"
by drinking water containing considew/l* *'*^» ' '*'- *
posited in joints, muscles, et<;., could J^ trttf-*-.*'^--**^
lithium salt and removed from the >xyjy :•. wj :".'•
truth there is in this assumption it is imi^^»» '^* '<*' '''•

Lithium gives two characteristic lineft is* ti* »->«*■ -
the yellow and the other in the red. It i^ycr* *
color to the flame.

354 PRINCIPLES OF INORGANIC CHEMISTRY

RUBIDIUM (At. Wt. = 85.5)

Occurrence, Preparation, Properties. — Rubidium occurs widely
distributed in nature, but nowhere in large quantities. It occurs
in canmllite, lepidolite, leucite, etc., and with potassium salts at
Stassfurt. It is also present in small quantities in the waters of
certain salt- wells, and was tirst discovered here by Bunsen and
Kirchhoff in 1860. As the result of the application of the spectro-
scope, with which they had accomplished so much, they found in
the liquor from the DUrkheim salt-wells lines which they could not
identify as belonging to any known substance. They succeeded in
isolating two substances, one of which gave two lines in the dark
red, and was, therefore, called rubidium.

Rubidium is prepared by heating the hydroxide with magne-
sium : —

2RbOH + 2Mg = 2Mg0 4-H,-t-2Rb.

Rubidium resembles potassium in appearance and properties. It
decomposes water with even greater violence than potassium, form-
ing the hydroxide. It is soft at ordinary temperature, melting
atl58^

Compounds of Bnbidinm. — Rubidium unites with dry oxygen at
ordinary temperatures, forming the dioxide, Rb02. It forms dark-
brown plates, melting at about 500**. When treated with water rubid-
ium dioxide forms i*ubidium hydroxide and liberates oxygen. The
hydroxide is also formed by treating rubidium carbonate with cal-
cium hydroxide, or rubidium sulphate with barium hydroxide. It
is a slightly stronger base than potassium hydroxide, dissociating to
a slightly greater extent at the same dilution.

The hcdides of rubidium resemble in general those of potassium,
but are more soluble. In addition to the ordinary chloride, bromide,
and iodide, rubidium forms compounds containing two halides, such
as RbICl4, and RblBrj. These are formed by bringing chlorine and
bromine, respectively, in contact with rubidium iodide. In such
compounds rubidium seems to have a valence much greater than
unity.

Notwithstanding the greater solubility of rubidium salts in gen-
eral, the perchlorate, RbC104, is far less soluble than potassium
perchlorate.

The sulphate, Rb2S04, and acid sidphate, RbHS04, of rubidium
resemble the corresponding potassium salts. When the acid sul-
phate is heated it readily yields the pyrosulphate : —

2 RbHS04 = H,0 H- Rb,SA,

LITHIUM, RUBIDIUM, CAESIUM, (AMMONIUM) 355

but this is far more stable than the corresponding potassium com-
pound.

Rubidium forms insohible compounds with hydrofluosilicic acid,
hydrochlorplatinic acid, etc. ; and it is, therefore, very difficult to
separate rubidium from potassium. The salt with hydrochlorplatinic
acid is, however, still less soluble than the corresponding potassium
salt, and this difference has been utilized to effect a partial separa-
tion.

CiESIUM (At. Wt = 132.9)

Occurrence, Compounds. — Caesium was first discovered by Bunsen
in 1860, in the waters of the DUrkheini salt-wells. In addition to
the two dark lines which led to the discovery of rubidium, he ob-
tained, after boiling down an enormous volume of the mineral water,
a few grams of the chloride of a substance which, when examined
spectroscopically, showed two lines in the blue. From the blue
color of its spectrum lines he termed this element caesium. It was
first isolated in 1881 by the electrolysis of the fused cyanide.
Caesium also occurs as the silicate in the mineral poUux from the
isle of Elba. The metal melts at 26^5.

The compounds of csesium resemble closely those of potassium
and rubidium. They are in general a little more soluble than the
corresponding compounds of rubidium. The hydroxide is even a
little stronger base than rubidium hydroxide. Caesium is, there-
fore, the strongest base-forming, or most electropositive, of all the
elements. Like inibidium, it apparently shows a valence greater
than unity towards certain of the halogens, especially iodine. It
forms with iodine the pentaiodide Cslj. Certain of the double com-
pounds of caesium are less soluble than the corresponding compounds
of rubidium, and these are used in separating the two elements.

AMMONIUM

The group ammonium, although not an element, closely resembles
in its properties the alkali metals. It forms a univalent cation,
NH4, which has the power to combine with the anions of acids and
form salts, which resemble in many respects those of the alkali metals.
As has already been mentioned, it combines with mercury like the
alkalies and forms an amalgam, which, however, is very unstable.

Ammonium HydrozidO) B^OH. — Ammonia combines with water,
forming the hydroxide NH4OH : —

NH3H-H,0 = NH40H.

356 PRINCIPLES OF INORGANIC CHEMISTRY

In the presence of water this compound is dissociated to some extent
into the ammonium ion, NH4, and the hydroxyl ion, OH : —

NH40H = NH4,OH.

It is, therefore, a base, but it is a very weak base. The small
amount of its dissociation is shown by its small conductivity.

V

Mr

10

3.1

100

9.2

1000

26.0

10000

61.0

60000

70.0

The concentration of hydroxyl ions in a normal solution of am-
monia as compared with a normal solution of sodium hydroxide
is about as 1 to 100. Ammonium hydroxide is, therefore, a rela-
tively weak base.

Before we had the conductivity method of measuring the disso-
ciation of bases and, therefore, their relative strengths, ammonium
hydroxide was regarded as a strong base, probably in part on account
of its action on the olfactory nerves and mucous membrane. This
error has been once for all corrected by the conductivity method.

When an aqueous solution of ammonium hydroxide is boiled it
breaks down into ammonia and water : —

NH,OH = NH8 4-HjO.

This fact is made use of in detecting the presence of ammonia or an
ammonium salt. The ammonium salt is treated with a strong base
like caustic soda, when it is broken down into the sodium salt, and
ammonia which is given off when the solution is heated. This can
be detected by the odor when present in considerable quantity, or
by holding a piece of moistened red litmus in the escaping vapors,
when the ammonia is present in small quantity. This becomes
colored blue.

Although a solution of ammonium hydroxide is only slightly dis-
sociated, it forms salts with practically all acids. Some of these
have characteristics which are sufficiently interesting to merit special
consideration.

Ammoninm Chloride, irH4Cl. — When hydrochloric acid is neutral-
ized with ammonium hydroxide and the solution evaporated, am-

UTHIUM, RUBIDIUM, CESIUM, (AMMONIUM) 357

monium chloride or sal ammoniac is obtained. This salt is a beau-
tifully crystalline compound, which is readily soluble in water.
Although ammonium hydroxide is only slightly dissociated, the salt
with hydrochloric acid is among the most strongly dissociated com-
pounds. This is true in general of the salts of ammonia with strong
acids. They are nearly as strongly dissociated as the corresponding
salts of strong bases like potassium or sodium. A concentrated
solution of ammonium chloride is, therefore, a concentrated solution
of ammonium ions and chlorine ions. If into such a solution contain-
ing a large number of ammonium ions ammonia gas is conducted,
the ammonium hydroxide formed will be dissociated far less than in
pure water at the same concentration. This could have been pre-
dicted from the law of mass action, and what has already been said
(p. 318) of the effect of one substance on the solubility of another
with a common ion. The presence of ammonium ions diminishes
the number of such ions which can be formed from the ammonium
hydroxide in the same solution, or, as we say, drives back the disso-
ciation of the ammonium hydroxide.

When ammonium chloride is heated it volatilizes at about 450^
Some of the most interesting phenomena connected with ammonium
chloride have to do with the condition of the substance in the form
of vapor. When ammonium chloride is volatilized it is dissociated
in part into the molecules NH3 and HCl by heat. The experimental
methods by which this is proved have already been discussed (p. 89).
The higher the temperature, the greater the amount of the salt broken
down into its constituent molecules.

If an excess of either ammonia or hydrochloric acid is present, the
dissociation of the ammonium diloride by heat is greatly diminished.
Here again we have an example of the influence of mass on chemical
activity. If ammonium chloride is volatilized into an atmosphere
which contains either of the products of dissociation, NH3 or HCl,
the amount of the dissociation is diminished. This is the same law
with which we have already become familiar in connection with
phosphorus pentachloride.

The volatilization of ammonium chloride is particularly interest-
ing, in that water plays such a prominent r51e in connection with
the dissociation of its vapor. Dry ammxmium chloride is only slightly
dissociated into ammonia and hydrochloric acid when volatilized. The
presence of water-vapor accelerates the dissociation of the ammo-
nium chloride. This is just the opposite of what we might expect,
since the presence of water is absolutely necessary in order that
ammonia gas should combine with hydrochloric acid gas.

358

PRINCIPLES OF INORGANIC CHEMlSTllY

Ammonium chloride in water shows a slightly acid reaction. This
is due to the hydrolysis of the salt of the weak base ammonia by
the water; — _ +

NH.Cl + H^O = CI, H + NH.OH.

Hydrochloric acid being strongly dissociated^ while ammonium
hydroxide is only weakly dissociated, there are more hydrogen ions
in the solutioii than hydroxl iona, and^ constHxuentlyj the solution
shows an acid reiiction.

AmmoniuiQ Hydrazoate or Triazoate, H^H^.^ — This salt of hydra^
zoic acid in reuirirkablc un account of its conii>osition. It containa
the same number of hydrogen and nitrogen atoms* It is obv^ioualy
the ammonium salt of hydrazoio acid : —

or.

HK, + NH.OH = H,0 + K^H^

As ire would expect from its composition, this salt is explosive, and

it explodes with violence on account of the large volume of gases
which it yields. N,H, = 2N, + 2H,

This salt is most readily prepared by means of complex organic
reactions, which it would lead us too far to take up in detail*

Ammonium Nitrite, NH^lfO^ — ^The nitrite is conveniently pre-
pared by the action of ammonium chloride on silver nitrite, insoluble
silver chloride being formed : —

NH.Cl + AgNOa = AgCl + l^H^KO^

When ammonium nitrite is heated it breaks down into nitrogen

and water : — KH^O, = 2 H,0 -h N^

It is of impoTtanec, as a means of preparing pure nitrogen. ^Vh€n
a solution of ammonium nitrite is heated, nitrogen is given oflf. In
preparing nitrogen in this way, it is only necessary tiJ mix solutions
of an ammonium salt and a nitrite and to heat the mixture.

Ammonium Nitrate^ NMJX%. — The nitrate of ammonium is
formed by the action of ammonium h^-droxide on nitric acid: —

KH4OH + HNO,=^ H,0 + KH,^^0^

The salt is very soluble in water, producing a marked lowering
of temperature. The dry salt is decomposed b;^ heat into nitrous
oxide and water : —

KH,N0, = 2H,0 + iSrA

LITHIUM, RUBIDIUM, CJiiSIUM, (AMMONIUM) 359

Ammonium nitrate is coming into use as an explosive. It decom-
poses when quickly heated to a high temperature, yielding water-
vapor, nitrogen, and nitric oxide, all of which are gaseous. The
volume of the gases set free is thus very great, and its power as an
explosive thereby increased. Further, the compound leaves no resi-
due when it explodes, and, therefore, the explosion takes place with-
out any appreciable amount of smoke or solid matter to contaminate
the gun. This substance has the further advantage that it is quite
stable under ordinary conditions.

Ammoninm Hydrosnlphide, Sulphide, and Polysnlphides. — When
a solution of ammonium hydroxide is saturated with hydrogen sul-
phide, the hydrosiilphide NH4HS is produced: —

NH4OH 4- HjS = H,0 + NH,HS.

This substance can be obtained in the form of crystals, most read-
ily by allowing ammonia gas and hydrogen sulphide to react in the
proper proportions. When volatilized, ammonium hydrosulphide,
like ammonium chloride, breaks down into its constituents — ammo-
nia and hydrogen sulphide.

The sulphide of ammonium^ (^Yi^^y is formed by treating the
hydrosulphide in solution with an equivalent of ammonia : —

NH4HS 4- NH4OH = (NHO^S -f H2O.

It is also formed by the action of ammonia gas on hydrogen sulphide.
Like the hydrosulphide, it can be obtained in the form of a solid,
which readily volatilizes. The vapor, like that of ammonium chlo-
ride and hydrosulphide, is dissociated by heat into the constituent
molecules," ammonia and hydrogen sulphide, as is shown by the
abnormally small vapor-density.

The aqueous solution of ammonium sulphide is colorless when
freshly prepared, but when allowed to stand for a time it becomes
deep-yellow in color. This is due to the oxidation of the sulphide
by the oxygen of the air setting sulphur free : —

(NHO^S 4- 0 = HjO 4- 2 NH3 4- S.

The sulphur thus set free acts on more ammonium sulphide, form-
ing the polysnlphides of ammonium. There are supposed to be
several of these compounds.

Ammonium sulphide is a very useful reagent in qualitative analy-
sis. The sulphides of metals which are soluble in hydrochloric acid
are readily throvm down by ammonium sulphide. The polysnlphides
of ammonium combine with sulphides of arsenic, antimony, gold.

360

PRIKCIPLES OF INORGANIC CHEMISTRY

platinum, and tin, and form sulpho-salts* These salts are soluble,
and this reagent h therefore useful to dissolve the sulphides of the
above five eletueiits and separate them from other substances.

Ammoniuni Sulphate, {KE^yfiO^. ~ This is one of the moat impoi^
tant salts of ammonia, as being one of the chief sources of the ammo-
nia used on a commercial scale* It is readily soluble in water, and
is hydtolytically dissociated by it When heated it yields the acid
sulphate : — (NH.VSO = NH, -h NH4HSO,-

Ammonium Carbonate, (irH^)jCO;,.H,XJ* — When a mixture of cal-
cium carbonate antl ammonium sulphate is distilled, and ammonia
passed into the aqueous solution of the product, normal ammomum
carbonate is formed. This is not very stable, and breaks down
readily into ammonia and the acid carbonate, I^KjHCOj: —

(KH,),C03 = NH, + NH.HCO^

The acid carbonate is also formed by the action of carbon dioxide
on aqueous ammonia. This is a much more stable substance than
the normal carbonate.

The two carbonates combine, and form what is known as the
sesquimrbt}nal€f (KH4)jC03*2KH4HCOb- Ammonium carbonate usu-
ally contains also the salt of an acid which bears a simple relation
to carbonic acid. The salt has the composition, lfl^H4C03NHj, and ia
known as ammonium earham^ale. It is obviously ammonium car-
bonate minus waters —

(NHO^COa - H,0 = KH,CO,KH|,

Ammonium carbonate is also formed by the action of ammonia gaa
on carbon dioxide.

Phosphates of Ammomum. — The primary and secondary phos-
phates of ammonium, NH4H^P04 and (KH|)3tHP0|j are well-known
substances. The tertiary phosphate (NH4)3P04 is probably formed
when concentrated ammonia is brought in contact with a concen-
trated solution of phosphoric acid. It is, however, very nnstable,
readily losing ammonia and passing over into the secondary phosphate.

When the solution of the secondary phosphate of ammonium is
boiled it loses ammonia and passes over Into the primary phosphate : —

(:N^H4),HP04 =^ NH, -h NHiH^PO*,

The double phosphate of ammonium and sodium, NaKH4HP0|,
or miGroco^mic salt, haii already been referred to (p, 329). When
heated it yields sodium metaphosphate: —

KH^NaHPO, - H,0 + KH^ +KaP03,

LITHIUM, RUBIDIUM, CiESIUM, (AMMONIUM) 361

Characteristics of the Alkali Metals in General. — From the fore-
going study of the alkalies we can draw general conclusions as to
their chemical behavior. In the first place, they are all strongly base-
forming elements^ which is the same as to say that when they are dis-
solved in water they form strongly electro-positive cations, there being
a corresponding number of hydroxy 1 anions present in the solution.

The alkalies form only xmivalent ions, which means that they
can carry only one electrical charge. We have already seen that
Faraday's law lies at the basis of chemical valence. This can be
tested directly in the case of the metals. When a given amount
of current is passed through a solution of any alkali chloride, the
amount of the corresponding hydroxide formed shows that the cation
is univalent. If we insert a solution of a silver salt into the current,
and allow the current to flow through this solution and then through
the solution of the alkali chloride, we would find that for every gram-
atomic weight of silver which separated, a gram-molecular weight of
the hydroxide of the alkali would be formed. Since a gram-atomic
weight of silver contains the same number of silver ions as a gram-
molecular weight of potassium hydroxide contains potassium ions,
or a gram-molecular weight of sodium hydroxide contains sodium
ions, it follows that an ion of silver carries just the same electrical
charge as an ion of potassium or an ion of sodium.

From the law of Faraday, then, all univalent ions carry exactly
the same electrical charge. Since it is mainly the ions that react
chemically, the question naturally arises, what connection exists be-
tween the electrical charges which the ions carry and their chemical
behavior. According to our present conceptions, the connection is a
very close one. The electrical attraction of these oppositely charged
parts, or ions, is undoubtedly an important factor in conditioning
chemical union.

This raises one further question. If all the alkalies are univ-
alent, from Faraday's law they all carry the same amount of elec-
tricity. Is there, then, no difference in the electrical energy carried
by a sodium ion from the electrical energy carried by a potassium
ion ? There may be a marked difference, and this is an important
point to note. Electrical energy, like every other form of energy,
is made up of two factors, a capacity factor, or quoMtityy and an in-
tensity factor, ox potential. While the quantity of electricity carried
by all univalent ions is the same, the potential of the charge varies
from ions of one kind to those of another. This explains why ions
of one kind will give up their charge under conditions which would
not remove the charge from ions of another element

862 PRINCIPLES OF INORGANIC CHEMISTRY

The ions of each of the alkali metals have certain characteristic
properties which enable them to be distinguished from one another.
Some of these have already been referred to. The more important,
from the standpoint of analysis, will now be summarized.

TTie sodium ion forms diflScultly soluble compounds with the

anion of hydrofluosilicic acid, SiF^ and the anion of pyroantimonic

acid, H^bA.

Tfie potassium ion forms difficultly soluble compounds with the

anioQ of chlorplatinic acid, PtCl«; with the anion of perchloric

.aoid, CIO4; ivith the anion of hydrofluosilicic acid, SiFa; with the

:anion of tartaric acid, H(C4H40«), and with the cobaltinitrite ion,

'Co(NO,)t.

The lithium ion forms difficultly soluble compounds with the anion

of carbonic acid, CO3, and with the anion of phosphoric acid, PO4.
The ammonium ion forms difficultly soluble compounds with the

anion of chlorplatinic acid, PtCl«; with the anion of tartaric acid,

HC4H40e, and with the cobaltinitrite ion, Co(N02V

The flame-tests for these several elements were considered when
each element was studied in some detail.

CHAPTER XXX

THE ALKALINE EARTHS

CALCIUM, STRONTIUM, BARIUM

CALCIUM (At. Wt. =40.1)

The metals which we have thus far studied are all univalent, or
their ions carry one electrical charge each. In the second group of
the metals the ions are nearly always bivalent, and in the calcium,
strontium, barium sub-group they are always bivalent. The salts
which these elements form with the anions of acids are of the general
type MClj, M(N03)2, MSO4, MC^Os, Mn(P04)2 and so on. With this
conception in mind we may now proceed to study in some detail the
compounds formed by the several members of the group.

Occurrence, Preparation, and Properties of Calcium. — Calcium
occurs very widely distributed in nature and in large quantities.
The carbonate occurs in great abundance as marble if well crystal-
lized and pure, or if impure as lirnestone or chalk. If in combination
with magnesium carbonate we have dolomite. Calcium phosphate
occurs in considerable quantity in certain phosphate beds. Gypsum,
or calcium sulphate, occurs in considerable quantity, while calcium
fluoride, or fluor-spar, exists in certain localities. Calcium comes
next to aluminium and iron in the order of abundance in the earth.

Calcium is best prepared by decomposing the iodide by metallic
sodium : — ^^ j^ 4. 2 Na = 2 Nal + Ca.

Calcium is a silvery-white metal, which decomposes water slowly
at ordinary temperatures. It combines with the oxygen of the air,
and also with the halogens at elevated temperatures. It melts in
a vacuum at 760**.

Calcium Hydride, CaHj. — This compound is formed by the action
of hydrogen on hot calcium. Barium and strontium form similar
compounds.

Calcium Oxide, or Lime, CaO. — Calcium combines with oxygen,
forming the oxide, CaO. This is most conveniently prepared by
heating the carbonate : —

CaC03 = CO, + CaO.

364 PRINCIPLES OF INORGANIC CHEMISTRY

Calcium oxide is a white, amorphous powder, which is extensively
used in a number of chemical operations. It is used as a carrier of
chlorine in the form of bleaching-powder, and in general wherever a
cheap base is desired. It also acts chemically upon various rocks
and minerals in the soil, liberating their constituents in soluble form,
so that they can be taken up by the plants.

Calcium oxide does not melt until a very high temperature is
reached (about 3000''). ' It is, therefore, used in constructing the
Drumvwnd tight. When the flame from the oxyhydrogen blowpipe
is allowed to play upon a cylinder of lime, it becomes highly heated
and at this high temperature gives out an enormous amount of light.
It resembles in this respect the oxides of thorium and cerium used
in the Welsbach light, the latter, however, giving out large amounts
of light energy at a much lower temperature.

When lime is brought in contact with moisture it takes up water
and forms calcium hydroxide : —

Ca0 4-H,0 = Ca(0HV

This process is known as slaking,

Calcinm Hydroxide or Slaked Lime, CaCOH)]. — When lime or cal-
cium oxide is thrown into water a large amount of heat is evolved,
and calcium hydroxide is formed as stated above. This is a white
powder, soluble in water only to the extent of 0.002 part in one part
of water. This is known as lime water, A mechanical suspension of
the finely divided calcium hydroxide in water is known as milk of lime.

Calcium hydroxide is a strongly dissociated compound, as is shown
by the following molecular conductivities. It dissociates thus : —

Ca(0H)2 = Ca,()H,dH.

p

Mr(«°)

04

381

128

440

256

419

512

427

Its solution contains a large number of hydroxyl ions, and it is,
therefore, a very strong base. It is, however, not quite as strongly
dissociated as the hydroxides of the alkalies. When a clear solution
of calcium hydroxide is allowed to stand exposed to the air for a
short time, it takes up carbon dioxide from the air, forming flakes of
the insoluble calcium carbonate: —

Ca(OH), -f CO, = CaCOa 4- H^O.

THE ALKALINE EAETIIS 365

When lime is exposed to the air the same reaction takes place to
some extent. The calcium oxide takes up moisture from the air, form-
ing the hydroxide, and this combines in part with carbon dioxide,
forming the carbonate. The white powder formed when lime is ex-
posed to the air, known as air-slaked lime, is, then, a mixture of cal-
cium oxide and calcium carbonate. Lime mixed with caustic soda
is known as sodarliine,

Compoimds of Calcimn with the Halogens. — Calcium combines
with the halogens, forming compounds of the general type CaAj,
where A re})resents a halogen anion.

The chloride^ CaClj.GHjO, is very soluble in water, producing
when dissolved a considerable lowering of temperature. By mixing
this salt in the proper proportions with finely powdered ice, a tem-
perature as low as — 30° to — 35° can be produced. The salt is most
readily prepared by dissolving marble, which is pure calcium
carbonate, in hydrochloric acid, and evaporating the solution to
crystallization. The salt when heated loses water, and if highly
heated hydrochloric acid, forming calcium oxide. When heated in
an atmosphere of dry hydrochloric acid gas all of the water can be
removed from calcium chloride without the salt undergoing any
decomposition.

On account of its attraction for water, anhydrous calcium
chloride is frequently used as a drying agent, especially for gases.
These are passed slowly through tubes filled loosely with calcium
chloride, and most of the water is removed from the gases and
absorbed by the chloride. Calcium chloride, however, does not
remove all of the water from substances. Indeed, it is not as good
a drying agent as sulphuric acid, and still less than phosphorus
pentoxide, which is the best drying agent known to the chemist.
Calcium chloride, however, cannot be used at all to dry ammonia
gas, since it combines with ammonia, forming definite compounds,
such as CaCl2.4NH3, and CaClj.SNHg.

Calcium bromide, CaBr^, and calcium iodide, Calj, present few
points requiring special comment. They are not very stable
compounds, the iodide especially breaking down in the pres-
ence of the oxygen and carbon dioxide in the air, yielding free
iodine.

Calcium fluoride, CaFj, or fluor-spar, is a beautifully crystalline
substance, practically insoluble in water, and is the chief source
of hydrofluoric acid and fluorine. Many varieties of fluor-spar are
strongly fluorescent, i.e. have the power of lengthening the wave-
lengths of the light which is allowed to fall upon them.

366 PRINCIPLES OF INORGANIC CHEMISTRY

Calcium Hypochlorite Bleaching-powder, Ca(0Cl)2. — When chlo-
rine is conducted into lime it is absorbed by the lime. The reaction
which takes place may be represented thus : —

(1) 2 Ca(OH)j 4- 2 CI, = CaCl, 4- Ca(OCl)j + 2 HA

giving a mixture of calcium chloride and hypochlorite ; or it may be
represented thus : —

(2) 2 Ca(0H)2 4- 2 01, = 2 Ca < ^^'^ -f 2 HA

forming one compound, which is half chloride and half hypo-
chlorite.

It is difficult to decide between these two possibilities. The fact
that bleaching-powder is not deliquescent, while calcium chloride
is strongly deliquescent, would indicate that there is no calcium
chloride in bleaching-powder. In aqueous solution the two reactions
would obviously yield exactly the same ions : —

(1) CaCla + Ca(0Cl)2 = C^ CI, c'l, 4- C^ OCl, OCl,

(2) 2 Ca < ^^^ = c4 CI, CI 4- Ca, OCl, OCl.

The difficulty in deciding between equations (1) and (2) is thus
apparent. All things considered, however, it seems probable that
bleaching-powder is a definite chemical composed of the composition

OCl
expressed by the formula Ca < .,, .

When treated with an acid, bleaching-powder gives up all of its
chlorine.

Ca < ^^4- 2 HCl = CaCl, 4- H,0 4- CI,,

Ca < 2P 4- H,S04 = CaS04 4- H.O 4- CV

Bleaching-powder is thus a very convenient means of transport-
ing chlorine without loss, since all the chlorine which was taken up
by the lime is set free when the bleaching-powder is treated with an
acid. This chlorine can be used for bleaching or for antiseptic pur-
poses. When bleaching-powder is exposed to the air it always has
the odor of chlorine. This is due to the action of the carbon dioxide
in the air, forming calcium carbonate and liberating chlorine.

on

Ca<^j 4- CO, = CaCOa 4- C\^

Calcium carbonate, being a stable solid not very soluble in water, is
formod.

THE ALKALINE EARTHS 367

When bleaching-powder is heated it forms calcium chlorate and

calcium chloride : —

on
6 Ca<^j = Ca(C103)j + 5 CaCV

Calcium chlorate can be used as the starting-point in preparing
potassium chlorate.

When bleaching-powder is brought together with certain
compounds rich in oxygen, like hydrogen dioxide, it gives up
oxygen:— ^p,

Ca<^j + HjO, = H^O + CaClj + Oj,

Half of the oxygen set free comes from the dioxide and half
from the bleaching-powder. This is the most convenient method of
determining the strength of a solution of bleaching-powder.

Sulphides of Calcium. — Calcium Hydrosulphidej Ca(SH)s, is formed
when hydrogen sulphide is conducted into a solution of calcium
hydroxide : — q^^q jj^^ + 2 H,S = Ca(SH), -f 2 HA

This salt, which is also formed when calcium sulphide dissolves
m water, — g CaS -f 2 H,0 = Ca(OH), -f Ca(SH)2,

has never been isolated from the solution. When an attempt is
made to obtain it, hydrogen sulphide is given off and calcium
sulphide remains behind : —

Ca(SH),= H,S-t-CaS.

Calcium sulphide was supposed for a long time to have the power
of emitting light in the dark, or to be luminescerU or phosphorescent.
This property has been shown to be due to the presence of small
quantities of the sulphides of certain metals, such as manganese.

Calcium Sulphate, CaSOi. — The sulphate of calcium containing
two molecules of water of crystallization, and known as gypsum —
CaS04.2H20 — occurs abundantly in nature. It dissolves in water
to some extent, 2 parts in 1000, and is frequently found in solution.
It also occurs in the solid form in many localities.

Gypsum is useful chiefly on account of the transformations which
take place when its water is removed by heat, and the anhydrous
salt is brought again into contact with water. When gypsum is
heated to 107° it loses one and a half molecules of water: —

CaS04.2 H,0 = CaSO^.i HjO H- 1^ H,0,
or, 2CaS04.2H,0 = 2CaS04.H,0 -f 3H,0.

868 PRINCIPLES OF INORGANIC CHEMISTRY

Gypsum which is thus partially dehydrated is a flowery powder,
and is known as plaster of parts. When brought in contact with
water, plaster of paris takes it up again and forms gypsum. The
mass, however, is now finally divided, and hardens after a few
minutes. It is used extensively for making mouldings and casts
of objects, especially of marble statuary.

If the temperature to which the gypsum is heated is at all high
(200°), it loses all of its water. When the completely dehydrated
product is brought in contact with water it combines with the water
very slowly, and is useless as far as making mouldings is concerned.
Such gypsum is known as " hard burned" or " dead burned " gypsum.

Calcium sulphate occurs in nature also in the anhydrous form. In

this condition it is known as anhydrite, and usually occurs in the

salt-beds deposited from seas which have evaporated. It can be

prepared by fusing together calcium chloride and potassium

sulphate : —

^ CaCl, -h K^04 = 2 KCl + CaS04.

Calcium Carbide, CaCj. — Calcium carbide has come into very
great prominence recently on account of its method of preparation,
and because when brought into the presence of water it readily yields
the illuminant acetylene.

The carbide of calcium is prepared by heating a mixture of finely
divided carbon and lime in an electric furnace : —

3 C -f CaO = CO -h CaC^

Calcium carbide has been prepared in the form of transparent
crystals. The product as it comes on the market is a grayish solid,
which, when exposed to the air has the odor of acetylene.

Its commercial value depends entirely upon the fact that it de-
composes with water, giving acetylene gas : —

CaCj -h H,0 = CaO + CjH^

Acetylene gas has recently come into great prominence as an
illuminant. This is due to the ease with which it can be made from
calcium carbide. Water is admitted slowly to the carbide when the
acetylene is desired, and a slow or rapid current of the gas generated
at will.

Calcium Carbonate, CaCOs- — The most abundant and important
compound of calcium which occurs in nature is the carbonate ccUcUe.
It occurs in beautifully transparent, crystalline masses in Iceland, and
is known as Iceland spar. Another crystalline variety of calcium car-

THE ALKALINE EARTHS 869

bonate is aragonite. This crystallizes in a different crystallographic
system from Iceland spai-, and we therefore have diamorphmn rep-
resented by calcium carbonate. When aragonite, which is ortho-
rhombic, is heated, it passes over into hexagonal Iceland spar. The
latter is, therefore, the stable, the former the metastable, phase. We
have here a case somewhat analogous to those already met with in
oxygen, sulphur, phosphorus, etc., and it is probable that the two
forms of calcium carbonate contain different amounts of intrinsic
energy in their molecules.

Calcium carbonate occurs in. great abundance in less beautifully
crystallized condition. Marble is a crystallized fgrm of calcium
carbonate. Limestone is also calcium carbonate and is usually crys-
talline, but the crystals are generally smaller than in marble and
are contaminated with various impurities. Clialk is a very fine-
grained variety of calcium carbonate, formed chiefly of the shells of
microscopic organisms. Indeed, calcium carbonate is frequently of
organic origin, consisting of shells of animals which have been more
or less metamorphosed by the geological processes to which they
have been subjected.

Calcium carbonate can be readily formed in the laboratory by
treating a soluble calcium salt with a soluble carbonate : —

CaCl, + Na,COs « 2 NaCl + CaCO,,
CaCNOs), + K,CO, = 2 KNO3 + CaCOj^.

When calcium carbonate is heated it undergoes decomposition into
lime and carbon dioxide, as we saw when we wei*e studying lime : —

CaCOs = CaO + COj.

At a given temperature this decomposition takes place until the car-
bon dioxide acquires a certain pressure. When this pressure is
reached the carbon dioxide combines as rapidly with the lime to
reform calcium carbonate as the latter decomposes. The pressures
of carbon dioxide at which equilibrium exists for several tempera-
tures are given below : —

Trmpkkatubbs

Pkxmubbs op Cabbost Diozidb

740^
810<>
865*'

4.6 millimetres of Hg

25.5 millimetres of Hg

67.8 millimetres of Hg

133.3 millimetres of Hg

2b

370

PRmCIPLES OF IXORGANIC CHEMISTRY

CftCO«.CO«

r,Oi.

C»^

ChO.COi

If we plot these results in a corve it Trill liave tlie form repre<
sen ted in Fig, l^S. We liave two components, Ciiiboii dioxide aud
limei and tlii-ee phaseis, carbou dioxidti, lime, and ealciiim carbonate*

We have, from the
phase rule, one degree
of freedom J and can
vary either the tem-
perature or pressure
along tiie curve with-
out destroying the
equilibrium.

As hajs already been
stated, lime is used
extensively in agricuU
ture. It is also used
extensively in archi-
tecture, as ino7-tarf for
holding together bricks
and stone. Mortars are made by mixing calcium liydroxide aud sand
in the presence of enough water to form a pasty mass. This is placed,
beneath each brick or stone^ upon the brick or stone below, and then
allowed to harden as it is said. The lime in in the form of the
hydroxide. When this comes in contact with the carbou dioxide in
the air it is retransforjned into carbonate : —

Ca(0H)5 + COj ^ H,0 + CaCO^,

and this is the chemical process which takes place when mortar
hardens. There is a large amount of water given off in the sense
of the above equation, and this agrees with universal experience
that a house which is freshly plastered is always damp. The harden-
ing or setting of mortar could be effected much more rapidly by
exposing it to an atmosphere rich in carbon dioxide, as by genemt-
ing large amounts of carbon dioxide in rooms which had been freshly
plastered.

When limestone, containing as impurities clay and magneainm
carbonate, is heated, the prcMluct forma with water a very hard mass.
This is known as kiffhaMlk cemenL

Portland cenmtit is made either from a mixture of pure calcium
earbonate and clay, or from marl, which is a mixture of silicates and
calcium carbonate.

Calcium carbonate is not readily soluble in pure water, but dis-
solves to a very considerable extent in water containing carbon dioxide.

THE ALKALINE EARTHS

871

Primary or Acid Calcium Carbonate, Ca(HC0s)2, is formed when the
normal carbonate is dissolved in water containing carbon dioxide: —

CaCOa 4- H,0 + CO, = CaCHCOs),.

Into water containing normal calcium carbonate in suspension
conduct a current of carbon dioxide, and the calcium carbonate will
pass into solution as the acid carbonate. The acid carbonate cannot,
however, be isolated. Indeed, when its solution is boiled the acid
carbonate is decomposed into the normal carbonate and carbon diox-
ide is given off.

When a solution of the acid carbonate is evaporated at ordinary
temperatures the normal carbonate is deposited. This is the way in
which the stdlaijina in caverns are formed. The waters containing
acid calcium carbonate in solution percolate through the roof of a
cavern, and evaporating, deposit calcium carbonate on the ceiling.

"* " *^^« solution of the carbonate trickling over the out-

-_i.i,. ^t^Tpjj hftautiful

il-
^he
of
ast
po-
ind
of
fre-
lag-
The
ced,
jium
3h is
vhile
»sum.
rable
n the
.tmos-
n car-

Pensioners under Act

^^ NO. OF
CERTIFICATE.

RATE.

m car-
' being
iich is
ers are

THE ALKALINE EARTHS 871

Primary or Acid Calcinm Carbonate, Ca(HC08)29 is formed when the
normal carbonate is dissolved in water containing carbon dioxide: —

CaCOs 4- H,0 + COi = CaCHCOsV

Into water containing normal calcium carbonate in suspension
conduct a current of carbon dioxide, and the calcium carbonate will
pass into solution as the acid carbonate. The acid carbonate cannot,
however, be isolated. Indeed, when its solution is boiled the acid
carbonate is decomposed into the normal carbonate and carbon diox-
ide is given off.

When a solution of the acid carbonate is evaporated at ordinary
temperatures the normal carbonate is deposited. This is the way in
which the staJagma in caverns are formed. The waters containing
acid calcium carbonate in solution percolate through the roof of a
cavern, and evaporating, deposit calcium carbonate on the ceiling.
This continues, the solution of the carbonate trickling over the out-
side of the deposit already formed, until frequently very beautiful
hanging columns result Such formations suspended from the ceil-
ing or sides of a cavern are called stalactites. The solution of the
carbonate frequently drops off of the stalactite, since the rate of
evaporation in a closed space beneath the surface of the earth must
be very slow. Where it drops on the floor of the cavern it evapo-
rates and deposits its carbonate, and we frequently find columns and
pillars of calcium carbonate growing upward from the floor of
caverns. Such growths are known as stalagmites. It not infre-
quently happens that the stalactite grows downward and the stalag-
mite upward until the two meet, and continuous columns result. The
beautiful and fantastic decorations of many caverns is to be traced,
then, to the action of water containing carbon dioxide on calcium
carbonate, resulting in the formation of the acid carbonate, which is
fairly soluble. This is true in caverns like Luray in Virginia, while
the stalagma in the Mammoth Cave of Kentucky are mainly gypsum.

Natural waters frequently contain carbon dioxide in considerable
quantity. This is produced by decomposing vegetable matter in the
soil through which they i)ercolate, and is also taken from the atmos-
pheric air. Consequently, many natural waters contain calcium car-
bonate in solution. Such are known as hard waters.

Where the hardness is due to the presence of acid calcium car-
bonate this is removed by boiling the water, the acid carbonate being
decomposed, as we have seen, into the normal carbonate which is
precipitated, and carbon dioxide which is set free. Such waters are
known as temporarily hard.

PRINCIPLES OF DfORGANIC CHE^MISTRY

Waters not infrequently contain in solution other salts of calcium,
as the sulphate, and also the salts of other metals. ^Vllen such water
is boiled these salts are not precipitated, and hence such waters are
permanenthf hard.

Phosphates of Calcium. — The three calcium salta of phosphoric
acid are all known, They are the normal salt, C%(P0^)2, the sec-
ondary salt, CaHF04, aud the primary salt, CaH4(P04)j.

Ti'iadcium phoisphuie, Cdt^ilX^t).^, is found in large quantity in the
hones of animals, and is therefore very important in connect ion with
animal life. When bones are heated to a high temperature in con-
tact with the air the organic matter is destroyed, and the calcium
phosphate and other mineral matter in the bones remain behmd in
the bone-asL

The normal calcium phosphate also occurs in nature mphospho-
ritef in combination with chloride or fluoride as ajfafite; and in
addition large beds of phosphate rock which are mainly of animal
origin occur in certain regions of the world, especially in the southern
parts of the United States, in Georgia, B-iorida, South Carolina, and
Tennessee.

The phosphoric ions — tertiary, secondary, and primary, PO4,

HPO4 or HhP04 — are of fundamental importance for the groi^iih, and
especially for the seeding, of certain plants and grasses. Among
these are the very valuable cereals wheat and corn. These plants
gradually remove the phosphoric acid ion from the soil, and the
latter would soon become impoverished in this substance, were it not
supplied to the soil artificially. The most important constituent of
commercial fertilizer is phosphoric acid ions. These, however, are
not supplied in the form of the tricalcium phosphate, since this salt
is not sufficiently soluble in water.

We shall see that normal calcium phosphate is readily trans-
formed into calcium pho.sphates which are soluble in water.

Normal calcium phosphate is formed when a soluble calcium salt
is added to a solution of disodium phosphate containing ammonia: —

3 CaClj + 2 Na^HPO* + 2 NH3 = 4 NaCl + 2 NH.Cl + Ca,(PO,)^

When the normal calcium phosphate is treaterl with an acid it read-
ily dissolves. If the acid is weak the secondary salt is formed; if
it is strong the primary salt is produced.

Secondanf mkium photqtlnU^^ CanPOj, is formed by the action
of a soluble calcium salt on disodium phosphate in the presence of a
little acid. If no aeid is added the tricalcium phosphate is first
formedi but since in this reaction acid is formed^ this acid reacts

THE ALKALINE EARTHS 378

slowly on the triphosphate, converting it into the secondary phos-
phate:—

2 Na,HP04 + 3 CaCl, = Ca,(VO,)t + 4 NaCl + 2 HCl,
Ca8(P04), -f 2 HCl = CaCl, + Ca,H,(PO02.

s

In this reaction we pass from the trivalent phosphoric ion PO4,
to the divalent, secondary, phosphoric ion HPO4.

If a little acetic acid is added to the solution of disodium phos-
phate and calcium chloride then introduced, the tricalcium phosphate
is formed at once : —

3 CaClj + 2 NaaHPO^ + (CHsCOOH) = C2l,(P0,\ 4- 4 NaCl + 2 HCl.

If tricalcium phosphate is treated with a strong acid, i.e. with a
concentrated solution of hydrogen ions in the proper proportion, the
primary calcium phosphate is formed : —

Ca3(P04)2 + 2 H8SO4 = 2 CaS04 + CaH4(P04),.

This is the commercial "superphosphate." In the presence of
iron or aluminium compounds it reverts, as it is said, probably form-
ing aluminium and ferric phosphates.

In preparing commercial fertilizer the tertiary calcium phosphate
is the starting-point in obtaining phosphoric acid ions. Although
this is acted upon slowly by the carbonic acid, and organic acids
formed from decomposing vegetable matter in the soil, and con-
verted into the more acid phosphates which are soluble in water,
this process is too slow to secure the best results.

The tricalcium phosphate in ground bone, or in finely ground
phosphate rock, is treated with sulphuric acid and converted into
the secondary and primary phosphates, which are somewhat soluble
in water. This is known as " soluble or available " phosphoric acid,
while the phosphoric acid in the form of tricalcium phosphate is
known as " insoluble " phosphoric acid.

In analyzing a commercial phosphate a given amount of the salt is
treated with a given amount of water at a given temperature, and
shaken for a given length of time. The phosphoric acid dissolved
by the water can then be determined by precipitating with ammo-
nium molybdate, dissolving in ammonia, and precipitating with the
"magnesia mixture" as ammonium magnesium phosphate. This is
heated and weighed as magnesium pyrophosphate. This is known
as "w^ater soluble" phosphoric acid. The residue is then treated
with a given amount of a standard solution of ammonium citrate, in
which the secondary calcium phosphate is soluble. This is known
as " citrate soluble " phosphoric acid. The part of the phosphr

ST4

PRINCIPLES or INORGANIC CHEMISTRY

ftoid insoloble in water and ammonium citrate is in combination with
eiJeitiiiit as triaUciuia phosphate, and 13 known as ** insoluble" phos^
phoHc acid* The " water soltibltj*' pbisj the "citrate soluble" phoa-
phL>ric acid are known as the " available phoapboric aftid.*'

Calciom SiMoata. CaSiOj, — The silicate of calcium occurs In na-
ture us i^^MoMmiite, It also occurs with other silicates in such
well-kuowil minerals as mim and garnet. Its chief importance is
in ooimeeliiKi with the tnannfacture of glass. Glass is, in general,
fta amorphous mixture of the silicate of calcium with the silicates of
the alkalies. There are a number of varieties of glass^ and a few of
these will be reconsidered,

(?/aA* is wade by fusing togetlier sand, calcium carbonate, and the
carUatiate of tlie alkali desired. If stHliuni cat-lionate is used we
have stMla glass, if potassium carbonate is employeil, potash glase, etc.

SoiUi i^htss is prepared by fusing together sodium carbonate, cal-
cium carbonate^ and sUicon dioxide, Soda glass is readily fusible^
and is easily attacked by chemical reagents such as boiling alkalies.
It is blown into cylinders, which are then opened aud flattened, and
cut into ordinary window panes. This is known as " soft glass,'^
because it is easily worked in the blast-lamp, and has applications in
the chemical labiiratory, although j on account of its solubility it is
not well adapted for bottles for holding chemical reagents.

Bohemian or ftottusmm giaits is a potassium calcium silicate. It
is much harder than the soda glass, fuses at a much higher tempera-
ture, and is much more resistant to chemical reagents. It is, there-
fore, valuable to the chemist, and is extensively employed in the
manufacture of apparatus which is to be heated to a high tempera*
ture, such as combustion-tubing and the like,

FHnt-glaAH consists of potiissium and lead silicates, the lead tak-
ing the place of ciilcium in ordinary smla or potassium glass. It has
such a high refnu^tive power that it is used in making optical lenses,

nftUiitm ^flint-ijlam contains thallium instead of potassium, and
has still higher refractivity and greater dispersion than ordinary
flint-glass,

Stmnit is a silicate of lead, potassium, and sodium, and also con-
tains some boric acid. It has such high refractivity that it is used
in making itu station gems.

The rokn'ed gimi^es are prepared by aflding to the fused silicates
oxldea of certain metals which give the desired color to the glass,
YfilQW glm99B owe their color to uranium or antimony ; him glasses
|o i^halt and mauganeae, red ghjHHes to copper, iron, or sometimes gold,
yrwn ij^aaaftf to chromium or copper, and so on. Glasses of almost

THE ALKALINE EARTHS 375

every shade of color have been prepared by using different coloring
constituents or mixtures of these constituents.

Most of the glass objects with which we are ordinarily familiar are
blown. A large ball of molten glass is taken on the end of a hollow
metal tube, through which the breath can be blown. The tube is
then moved rapidly backwards and forwards through the air beneath
the glass blower, who drives air through the tube at the desired rate
and time. The glass takes the form usually of a hollow cylinder,
which is blown out to the desired thickness. This is cracked, flat-
tened, and cut into the desired size. In this way flat panes of glass
are made. Bottles are blown into moulds, and other glass objects
of definite shape are either blown into moulds or moulded. Such
objects must be annealed.

The property of the glass is largely conditioned by the way in
which it is annealed. If glass is cooled with moderate rapidity it
has the properties which we ordinarily associate with it. If cooled
very rapidly, however, it has very different properties. It is under
a considerable strain, and when the surface is fractured in any way,
the glass flies in pieces almost with explosive violence. The Prince
Rupert drops, made by dropping molten glass into water, are exam-
ples of this condition.

On the other hand, when glass is cooled very slowly, as by intro-
ducing it when hot into hot oil, or by placing it in an oven which is
cooled slowly, it is much less easily broken than ordinary glass. It
acquires considerable elasticity^ and can be struck a fairly hard
blow without injury.

Calcinm Oxidate, CaCA- — We have already studied a number of
salts of calcium which are practically insoluble in water. These
include the phosphate, carbonate, and, to some extent, the sulphate.
Another compound of calcium, which is only slightly soluble in pure
water, is calcium oxalate — the calcium salt of the dibasic, organic
acid H2C2O4. When a solution of any calcium salt is treated with a
solution of ammonium oxalate, calcium oxalate is precipitated : —

CaClo -f. C AC^H^), = 2 NH4CI + CaC A.

Ammonium oxalate is used to precipitate calcium oxalate and not
free oxalic acid, since calcium oxalate is soluble in strong acids. If
free oxalic acid were used in the above case, instead of ammonium
oxalate, hydrochloric acid would be liberated, and this would redis-
solve the oxalate. Oxalic acid is a dibasic acid, and the above equa-
tion should be written : —

Ca, CI, Ci -h NH4, NH4, C A = NH4, Ci 4- NH4, Ci + CaC A.

876 PRINCIPLES OF INORGANIC CHEMISTRY

Calcium is frequently preeipitated as the oxalate in quantitative
determinations of, this element. The oxalate loses more and more
water of crystallization as the temperature is raised, and is, there-
fore j not a good salt to weigh. The oxalate is easily decomposed to
the oxide, which is the form in which the calcinm is most convea*
iently weighed : - ^^^^^^ ^ ^^^ ^^^^ ^^^

It may also be decomposed into the carbonate by careful heating,
and the carbonate theii converted into the oxide : —

CaCA = CaCOa + CO.

Detection of Calcium. — In addition to the insoluble compounds
formed by calcium, the spectroscope furnishes a vaUiable means of
detecting calcium. When a calcium salt is introduced into the color*
less flame of a Eunsen burner, the flame is colored orange-red, and
can easily bo recognized with a little practice. If examined with the
spectroscope the calcium flame shows a heav3' line in the orange-red,
another in the green, and a still fainter Hue in the blue,

STRONTIUM (At Wt* = 87.8)

The element strontium resemblea calcium very closely in its prop-
erties ami in the properties of its compounds. It will, there fore,
be treated very briefly, certain differences between the two being
pointed out.

Occurrence, FreparatioB, and Properties of Strantium. — Strontium
occurs in nature ebiefiy in the form of two salts which are well-known
minerals. These are strontium carbonate or stroniiamtet and stron-
tium sulphata oTcekstite.

T]je element is prepared most conveniently by electrolyzing the
fused chloride, the metal separating at the cathode.

Strontium resembles calcium in its appearance and properties*
It has a yellowish tint, com bines with the oxygen of the air, acts
upon w^ater setting hydrogen free and forming the hydroxide, and
in general so closely resembles calcium that it is unnecessary to
describe its properties in detail. ^^

Salts of Strontlnm* — Strontium forms the divalent ion 8r, which

is strictly analogous to the calcium ion Ca. It combines with two
hydroxy! ion a forming Btrontinm hifdmxkhf Sr(0H)2, whose aqueous
solution is strongly basic. This substance is more soluble in water
than calcium hydroxide, and crystallizes from the aqueous solution
with eight molecules of water: Sr(0H),.8HA Strontium niinxte^

THE ALKALINE EARTHS 377

Sr(N08)2, like strontium salts in general, gives a beautiful red color
to a colorless flame. It is used, because of this property, to produce
red light in fireworks and other pyrotechnic displays. Strontium
nitrate is insoluble in alcohol, and thus differs from calcium nitrate^
which dissolves readily in this solvent. This fact is made use of
to separate strontium from calcium.

The strontium ion combines with the anions of acids, forming in
general the same insoluble compounds as calcium. Certain differences
in the degree of solubility however, manifest themselves. Strontium

combines with the carbonic ion CO3 forming strontium carhoriatey
SrCOg. As already stated, this occurs in nature as the mineral stron-
tianite, and is the source of the element strontium. It is practically
insoluble in water and is, therefore, precipitated when a soluble
carbonate is added to a soluble strontium salt : —

Na^COs + SrClj = 2 NaCl + SrCO,.

Strontium carbonate dissolves readily in hydrochloric or nitric
acid, forming the corresponding chloride or nitrate. It is difficult to
decompose strontium carbonate, a very high temperature being re-
quired.

The chloride contains six molecules of water of crystallization,
SrClj-GHgO. It readily takes up water from the air but not as
readily as calcium chloride. Strontium chloride is easily soluble in
alcohol, and thus differs from barium chloride, which is insoluble in
this solvent.

Strontium combines with the gulphuric ion SO4, forming strontium
sulphate, SrSOi- This salt occurs in nature as celestite, and since it
is only slightly soluble in water, is formed when a soluble sulphate
is added to a soluble strontium salt : —

Na^SO* 4- SrCl, = 2 NaCl 4- SrS04.

Strontium sulphate is much less soluble in water than calcium sul-
phate, and is practically insoluble in a mixture of water and alcohol.
The strontium ion forms difficultly soluble compounds also with

S IB

the ions PO4 and HPO4. The normal phosphaie, Sr3(P04)8, and
secondary phospJiate, SrHP04, are quite insoluble. With the ion of

IB

oxalic acid, Cj04, the strontium ion combines, forming insoluble stron-
tium oxalate, SrCj04.

The chromxUe of strontium, SrCr04, is fairly soluble.

Detection of Strontium. — Strontium is easily detected by the
color which it imparts to the flame. It produces an intensely dark-

878 PRINCIPLES OF INORGANIC CHEMISTRY

red flame^ which can be easily recognized. When this flame is
examined with the spectroscope it is found to contain a number of
lines in the red and orange-red, and one characteristic line in the
blue. Since strontium is the only common substance which gives a
deep-red flame, its presence is usually detected by the flame reaction
alone. Methods of separating it from calcium and barium will be
considered a little later.

BARIUM (At. Wt. = 137.4)

An element closely allied to calcium and strontium is barium.
This element occurs chiefly as the sulphate, BaSOf, which is known
as barite or heavy Rpary and as the carbonate, BaCOa, known as
wUherite,

Barium, like strontium and calcium, is prepared by electrolyzing
the fused chloride. The metal barium is white, takes up oxygen
from the air, and decomposes water with evolution of hydrogen.
The reaction with water is more vigorous than that of calcium or
strontium, and in this it more closely resembles the alkalies.

Oxides of Barium. — Barium forms two oxides — the normal
oxide, BaO, and the dioxide, BaOj. Barium oxide, BaO, is formed by
heating the nitrate.

2 BaCNOs), = 2 BaO + 4 NO, -f 0^

It can be prepared from the carbonate, not conveniently, however,
by heating directly, since the carbonate decomposes only when
heated to a very high temperature. It can, however, be readily
prepared by heating the carbonate with carbon : —

BaCOs + C = 2 CO + BaO.

Barium dioxide, BaO,, is formed by heating the oxide to 500® in
a current of air or oxygen : —

2 BaO + 0, = 2 BaO,.

Barium dioxide is an excellent "carrier of oxygen,'^ since at a
somewhat higher temperature it gives up its excess of oxygen and
forms barium oxide again : —

2 BaO, = 2 BaO -fO,.

We have already become familiar with this substance in connec-
tion with the preparation of hydrogen dioxide. When it is treated
with an acid the following reaction takes place : —

BaO, + 2 HCl = BaCl, -f H^O,.

THE ALKALINE EAUTIIS

379

With water, barium dioxide forms the hydrate with eight mole-
cules of water of crystallization — BaO,.8 H,0.

Barium Hydroxide, Ba(OH)s. — The hydroxide of barium is formed
when the oxide is dissolved in water, BaO -h HjO = Ba(OH),. Barium
oxide is much more soluble in water than strontium oxide, which in
turn is more soluble than calcium oxide.

The hydroxide of barium crystallizes in beautiful, white plates,
containing eight molecules of water — Ba(0H),.8H,0,

The hydroxide is readily soluble in water, especially at an ele-
vated temperature. It dissolves in three parts of boiling water and
in twenty parts of cold water ; a solution saturated when hot, there-
fore, crystallizes out in abundance when cooled. A solution of
barium hydroxide in water is known as baryta water. This solution
is strongly basic, showing that barium hydroxide readily dissociates
thus: — 4.+ _ .

Ba(OH),= Ba,OH, OH.

This is shown by the large values of its conductivities.

r

M,(«y>)

8

340

04

402

512

437

1024

440

As a strong base it is frequently used to neutralize acids, and to
standardize solutions of acids by titration, using some of the indi-
cators already studied to determine when the neutralization is just
complete. A solution of baryta water is frequently used to detect
the presence of carlnm dioxide. On account of the insolubility of
barium carbonate, a minute trace of carbon dioxide can be detected
by means of this reagent : —

Ba(On), -h CO, = BaCO, -f 11,0.

Barium Chloride, BaCl, . 2 H,0 — The chloride of barium is
formed by dissolving the carbonate (witherite) in hydrochloric

^^^ '' ~ BaCO, + 2 HCl = H,0 -h CO, -h BaCl,.

Also by treating barium sulphide with magnesium chloride and
water : —

BaS -f- MgCl, -h 2 HiO = H,S -h Mg(OH), -f- BaCl,.

380 PRINCIPLES OF INORGANIC CHEMISTRY

Barium chloride is less soluble in water than strontium chloride,
which is less soluble than calcium chloride. This is exactly the re-
verse of the solubilities of the oxides^ barium oxide being the most
soluble, strontium oxide less soluble, and calcium oxide the least
soluble of the three.

Barium Sulphate, BaS04. — The sulphate of barium, or heavy spar^
occurs in nature as stated above. It is readily formed whenever a
soluble sulphate is added to a soluble barium salt : —

Ba(N03), -f K8SO4 = 2 KNOs -f BaS04.

It is the most insoluble sulphate known, and is, therefore, the
form in which sulphuric acid is precipitated and weighed in quanti-
tative determinations of this acid. It is also the most insoluble salt
of barium, and the form in which barium is determined quantita-
tively.

Barium sulphate is used as a white pigment, under the name of
pe^'manent white.

One decomposition of barium sulphate is of more than ordinary
interest. When the sulpJiate is boiled with a solution of sodium car-
bonaie, it is transformed in part into barium carbonate : —

BaS044- Na^COs = Naj^04 -f BaCOa-

This transformation is at first only partial. If the solution of
sodium carbonate and sodium sulphate is poured off after a time, and
a new solution of sodium carbonate added, the decomposition of the
barium sulphate into carbonate will proceed farther. By repeating
this for a few times, practically all of the barium sulphate can be
transformed into carbonate. This is one of the very best examples
of the effect of mass on chemical activity. By renewing the solution
of sodium carbonate, i,e, by increasing its mass, and by pouring off
the solution of sodium sulphate formed, i.e. by decreasing its mass,
the transformation in the sense of the above equation can be made
practically complete.

Barium Carbonate, BaCOs. — The carbonate of barium, or witherite,
is the most convenient starting-point in the preparation of compounds
of barium. It is easily formed by bringing together a solution of a
soluble carbonate with a soluble barium salt : —

Na^COa + BaCl, = 2 NaCl -f BaCOj^

We have seen that strontium carbonate is far more difficult to
decom])oso into the oxide and carbon dioxide than calcium carbonate.
Barium carbonate is still more difficult to decompose than strontium

THE ALKALINE EARTHS 381

carbonate, giving off only a little carbon dioxide when heated to a
white heat. ^^

Phosphates of Barium. — The barium ion, Ba, combines readily
with the ions of phosphoric acid, forming insoluble compounds.
The normal phosphate is formed by the union of the barium ion,

Ba, and the phosphoric acid ion, PO4, and has the composition
Ba3(P04)j. The secondary phosphate is formed when the barium ion,

Ba, and the secondary phosphoric ion, HPO4, unite. It has the com-
position BaHP04.

Other Insoluble Compounds of Barium. — The barium ion com-
bines with the oxalic ion, C2O4, forming insoluble barium oxalate,
BaC204. It also combines with the chromic ion, Cr04, forming
insoluble barium chromate, BaCr04, and with the hydrofluosilicic
ion, SiFe, forming insoluble barium fluosilicate, BaSiF^

Detection of Barium. — Barium is easily detected by its charac-
teristic green flame. This flame persists for a long time when a
barium salt is held in a Bunsen burner. When examined spectro-
scopically, a bright green line appears, of a definite wave-length, and
this is characteristic of barium. The barium line has a slightly
shorter wave-length than the corresponding calcium line.

Belations between Calcium, Strontium, and Barium. — The three
alkaline earth metals resemble one another closely in their chemical
properties. Certain differences, however, manifest themselves. The
different solubilities of the salts of these elements with a given acid
are very important, especially in connection with the separation of
these elements from one another.

Take the hydroxides of the three elements : Calcium hydroxide
is the least soluble ; strontium hydroxide is more soluble than cal-
cium ; while barium hydroxide is still more soluble than strontium
hydroxide.

When we turn to the sulphates we find exactly the opposite rela-
tions. Barium sulphate is the most insoluble of the three ; then
comes strontium sulphate, and finally calcium sulphate.

All three of these elements form insoluble carbonates, while
barium alone forms an insoluble chromate.

These differences in solubility are made use of to detect the alka-
line earths when in solution in the presence of one another.

Detection of the Alkaline Earths — Calcium, Strontium, and
Barium. — Given a solution containing calcium, strontium, and
barium ions, how would these be detected in the presence of one
another ? As we have seen, all of these ions form insoluble com-

382 PRINCIPLES OF INORGANIC CHEMISTRY

pounds with the ion of carbonic acid. Therefore, if ammonium car-
bonate is added to a solution containing these substances, calcium,
strontium, and barium carbonates are precipitated.

The three carbonates are filtered off, washed, and dissolved in a
little dilute nitric acid. The solution is evaporated to dryness, and
the residue heated until all traces of nitric acid have disappeared.
The residue is then treated with a mixture of equal parts of abso-
lute alcohol and ether. Calcium nitrate dissolves, while strontium
and barium nitrates remain undissolved. The solution is then fil-
tered off from the residue and treated with a few drops of dilute
sulphuric acid, when calcium sulphate is precipitated.

The residue is washed carefully with the mixture of alcohol and
ether to remove every trace of calcium, and then dissolved in a little
water. A part of the solution is then treated with a few drops of
acetic acid, and a solution of potassium chromate added. Barium is
precipitated as the chromate.

To the solution from which all the barium has been precipitated
as the chromate, add ammonium carbonate and ammonia, when the
strontium will be thrown down as strontium carbonate.

CHAPTER XXXI

THE MAONESIUM OROX7P

GLUCINUM, MAGNESIUM, ZINC, CADMIUM, MERCURY

GLUCINUM (At. Wt. = 9.1)

The first two members of this group, glucinum and magnesium,
correspond rather closely in their properties to the elements of the
calcium group. The first of these elements, glucinum or beryllium,
is comparatively rare. It occurs chiefly in the mineral beryl, which
is a silicate of aluminium and beryllium. It has the composition
Al,(Si03)j-h3GlSi03. It also occurs as clirysoberyl, having the
composition GIO. AI2O3.

The element is conveniently prepared by electrolyzing the chlo-
ride. It is white, and decomposes water only at an elevated tempera-
ture. It decomposes water only slowly under these conditions, and
thus differs from the metals of the calcium group.

Glucinum forms with water the hydroxide Gl(OPI)j, which has
basic properties and, therefore, yields hydroxyl ions. The hydrox-
ide dissociates in the presence of water as follows : —

Gl(OH)a = Gl, OH, OH.

The dissociation of glucinum hydroxide is not as great as that of
calcium hydroxide, or, in a word, it is not so strong a base. This is
shown by the fact that in the presence of a strong base like sodium
hydroxide it acts as an acid, forming salts with the sodium ion. We
shall see that this property is manifested repeatedly by succeeding
members of this group, and shall discuss it more fully in connection
with them.

Glucinum combines with the anions of strong acids, forming salts

in which the glucinum ion, Gl, is bivalent. The chloride of glucinum
has the composition GlClj. The sulphate, GISO4, crystallizes both
with four and with seven molecules of water, forming GISO4.4H2O,
and GISO4.7H2O. The carbonate of glucinum, GICO3, is soluble in
water.

883

S84

PRINCIPLES OF INORGANIC CHEMISTRY

Glucinum readily combines with silicic acid forming ailieates
wMch have value m gems.

MAGNESIUM (At. Wt = 24.36)

The second member of the series in the order of mcreasmg atomic
\Tveights, but the first member in the order of importance, and the one
tfrom whieh the aeries takes its name, is magnesium, Magnettium
^occurs in nature in the form of several salts. Magnesium earljonate,
mftgmsite^ MgCOf»j is fairly widely distributed, while the double car-
bonate of magnesium and calcium, dolomite ^ forms whole mountain
ranges, Magnet>ium also occurs as the sulphate, MgSO^.HjO, — ktes-
erite. It occurs in large quantities its Jcahiiie the double sulphate
of magnesium and iwiassium^ which also contains magnesium chlo-
ride — MgSO^. K^SO|, MgCl-i. 6 H^O. dmutUUe— the double chloride
of magnesium and potassium — MgKCl 3. Ti HaO, is also found iu
large quantities in certain salt deposits. The silicates of magnesium
and other metals constitute some of the best-known mineralS| such
as hombkndef sei-f^entlnef tah, mbefitofty etc.

Magnesium is prepared by elect ro]y?/mg fused, anhydrous, carnal-
lite. It is a white, ductile metal, which can be readily drawn into
wire. In contact with the air at ordinary temperatures it cond>ines
with oxygen very slowly. When heated it takes fire and burns with
a brilliant white light, which on account of its richness in light of
short wave-length is frequently used for illuminating purposes ia
photography. Magnesium, on account of its power to combine with
oxygen at elevated temperatures, is an excellent reducing agent, and
is frequently used, as we have seen, to remove oxygen from the
oxides of other elements.

Magnesium decomi>oses water very slowly, indeed, at ordinary
temperatures, but quite rapidly when the water is hot. It thus dif-
fers from the elements of the calcium group, and resembles more
closely the Inter memlwrs of this series. Magnesium combines
readily with the halogens.

Magnesium Oxide, MgO, and Magnesium Hydroxide, MgiOB.)^ —
Magnesium oxide, or magnesia , is formed when magnesium is burned
in the air, or when the carbonate or hydroxide of inagtjesium is heated.
Magnesia melts only at enormously high temperatures. It is, there-
fore, used to line vessels which are subjected to high temperatures.

Magnesium oxide dissolves in water to only a slight extent,
forming, however, magnesium hydroxide. Magnesium hydroxide,
Mg(OH)„ is not very soluble in water, and is precipitated from

THE MAGNESIUM GROUP 385

a solution of a magnesium salt by treating this with a soluble
hydroxide: —

MgClj + 2 NaOH = 2 NaCl + Mg(OH),.

Magnesium hydroxide dissolves readily in ammonium salts, and
this is of imix)rtance in the detection of magnesium.

Magnesium Chloride, MgCls.GHxO. — Magnesium chloride, like the
chlorides of calcium and strontium, contains six molecules of water
of crystallization. It occurs in nature in combination with potassium
chloride as carnaJUte, and can readily be prepared by dissolving
magnesium carbonate or oxide in hydrochloric acid. When heated
it does not give up its water of crystallization, leaving the anhy-
drous salt behind, but undergoes decomposition, giving off hydro-
chloric acid. The decomposition may be represented thus: —

MgCl, -f. HjO = 2 HCl -f- MgO.

This reaction is now of importance for obtaining hydrochloric acid
on a commercial scale.

Magnesium Sulphate, Mg804.7H20. — Magnesium sulphate con-
taining one molecule of water occurs in nature, as we have seen, as
the mineral kieseritej MgS04.H20. The ordinary salt, known as
Epsom salt since it is contained in the Epsom springs, contains seven
molecules of water, MgS04. 7 H,0. Magnesium sulphate when heated
loses six molecules of water at 150®. The last molecule is retained
until a temperature of about 200® is reached. We shall find other
examples of this same behavior — salts containing water of crystal-
lization retaining the last molecule to a much higher temperature
than the other molecules. Magnesium sulphate is used in medicine
as a purgative.

Magnesium sulphate has the power to combine with alkaline
sulphates and form double sulphates. An example of this class of
substances is the mineral schonite, which has the composition
MgSO4.K2SO4.6H2O. This contains one molecule of magnesium
sulphate and one molecule of potassium sulphate. Another double
sulphate of magnesium and potassium is the mineral langheinitej
K2S04.2MgSC)4, containing two molecules of magnesium sulphate to
one of potassium sulphate.

Magnesium Carbonate, MgCOj. — As already stated, magnesium
hydroxide is a weak base and carbon dioxide a weak acid. When
the two are brought together they form a salt, but there is a strong
tendency to form a basic salt. W^hen a soluble carbonate is brought
together with a soluble magnesium salt, the normal carbonate of
2g

38U

I'RINCtPLES OF INORGANIC CHEMISTRY

k

nargaesium is not formed, but cai-bon dioxide escai>es, and there
results a basic carbonate, whose eomi>o3ition depends upim the dilu-
tion and temperature of the solutions. The greater the dilution of
the solutions, and the higher the ten^iierature when they are mixedj
the more basic the carbonate which is precipitated. The carbonate
formed by adding to a solution of magnesium sulphate a solution of
an alkali carbonate has approximately the composition

Mg,(OH),(CO^t = 2MgC0, + Mg(OH)^

This is known as magneski <Ma,

When carboti dioxide is passed into water containing magnesia alba
in suspension, the normal earbooate of magnesium, MgC03,3HjO,

crystallizes out of the solution.

^[agnesium carbonate, like magnesium sulphate, forms double
salts. The best known of these double carbonates is that formed
with calcium carbonate. As already mentioned, magnesium carbo-
nate and calcium carbonate crystallize together as dolomite. Tliis has
the composition MgCO^^CaCO^j and occurs in great abundance in
nature. These two carbonates can apparently crystallize together in
all proportions, since many such mixtures are known having very
different proportions of the two substances present in them.

Phosphatei of Hagnefiimu. — Magnesium forms the primary^
Mg(HJHX)^; the secondary, MgHPO,; and the teitiary, or normal,
phosphate, Mg^^(PO<)j. These resemble the phosphates of calcium so
closely that a detailed discussion is unn<?ces8ary. One phosphate of
magnesium, however, is of special interest. This is the ammonium
magnesium phosphate alre^ly mentioned — NH^MgPO^.G H^O. This
compound is of special importance, since it is the form in which
phosphoric acid is precipitated in quantitative determinations of this
acid. It is also the form in which magnesium is precipitated in
quantitative determinations of magnesium* When heated it decom-
poses as follows ; —

2 NH.MgPO^ = 2^H, ^ H3O + Mg,P,0,,

yielding magnesium pyrophosphate, which is a very stable sub.
stance and of constant composition. It is, therefore, an excellent
form in which to weigh either phosphoric acid or magnesium,

Silicttei of Hagnesium, — Magnesium combines with silicic acid
or the poly silicic acids, forming some of the best-known minerals*
Ordinary kth is a silicate of magnesium having the cojupoHitiun
Mg^H^i^Oj,, Serjjentine is a silicate of magnesium liaving the composi-

THE MAGNESIUM GROUP 387

tion MggSijO; . 2 H,0. Olivine is the normal silicate of magnesium,
Mg2Si04. Soapstone is a more complex silicate.

Magnesium silicate combines with other silicates, forming such
well-known minerals as hornblende, tourmaline, etc.

Magnesium Nitride, Vg^sy is formed by the direct action of
nitrogen on red-hot magnesium. Magnesium nitride decomposes
with water in the sense of the following equation : —

MgaN, + 6 H,0 = 2 NH3 -h 3 ^rg(OH)^

Separation of Magnesium from the Elements of the Calcium
Oroup. — In an ordinary qualitative analysis magnesium would be
precipitated as the basic carbonate along with the carbonates of cal-
cium, strontium, and barium. To prevent this, ammonium chloride
is added to the solution until on addition of ammonia no precipitate is
formed. The calcium, strontium, and barium are then thrown down
by adding ammonium carbonate, as the corresponding carbonates.
These are filtered off, leaving the magnesium in the filtrate.

To the solution containing the magnesium, a solution of disodium
phosphate is added. The magnesium is precipitated as a secondary
phosphate when no ammonia is present, but since some ammonia is
added, as magnesium ammonium phosphate.

Magnesium does not give any characteristic flame reaction and
cannot, therefore, be detected by this means.

ZINC (At. Wt. = 05.4)

An element closely allied to magnesium in many of its properties,
but differing from it markedly in others, is zinc. This well-known ele-
ment occurs in nature in abundance in a number of compounds. Zinc
blende, which is the sulphide ZnS, is one of the most common. It also
occurs as the carbonate ZnCOs smithsonite, the silicate Zn2Si04.HjO
calamine, and in combination with iron oxide as fratiklinite.

Zinc is prepared by reducing with carbon the oxide, which is
usually obtained by roasting the sulphide in the air : —

ZnO 4. C = CO -f Zn.

The metal boils at 950®, and is, therefore, easily distilled off when
once set free. It condenses in the form of a fine powder or dust,
which is known as zinc diist. The metal can be melted and cast into
sticks, or poured into water while molten, and produces granulated zinc.

Zinc manifests remarkable behavior on warming. At ordinary
temperatures it is brittle. When heated above 100® it becomes mal-
leable, and can be rolled into sheets, drawn into wire, etc. Having
once acquired this property, it retains it even at ordinary tempera-
tures. When heated still higher, — to somewhat above 200**, — it

PRIXCIPLES OF INORGAXIC CHEMISTRY

becomes bi'ittle again, and can be readily pulverized. Having
become brittle again, it retains this property when the temperature
is reduced to the ordinary. Zinc melts at 420^, and combines with
oxygen when heated to a much higher temperature, forming zinc
oxide. Zinc decomposes water very slowly at ordinary tempera-
tures^ but more rapidly at elevated temperatures* The vapor-density
dhows that the molecule in the form of vapor contains one atom.
Zinc is dissolved by hydrocdiloric and nitric acids. In the former
case, hydrogen is evolved; in the latter it acta upon more nitric
acid, reducing itj liberating oxides of nitrogen. In sulphuric acid
pore zinc dissolves very slowly, if at all. If the zinc is impure^
however, it dissolves readily in sulphuric acid with evolution of
hydrogen gas. If a pieoe of platinum wire is wrapi^d around pure
zinc and the whole plnn^ed int^ sulphuric a«.?id, the xinc dissolves
rapidly, and hydrogen is evolved on the platiimui. If a little pW
tinic chloride is added to the solution of the acid aronnd the zinc,
the same result is produced. This is due to electrical action, the
hydrogen from the acid separating more easily upon tlie impurity,
or upon the platinum^ than upon the zinc. Zinc is used for covering
objects which, if exposed to air or water, w^ould rust. Iron objects
thus protected are known as galmnized. Iron objects are galvanized
by dipping them into molten zinc. Zinc is also extensively used in
constructing primary cells* The chief uses of zinc, howe%^er, are in
alloys with other metals. The best known of these is hrctss^ an alloy
with copper. When nickel is added, we have the alloy of zinc, cop-
per, and nickel, known as German silver. Zinc reatlily combines with
mercury when the clean surface of the metal is brought in contact
with the mere my, forming zinc amahfam.

Zinc OzidCj ZnO^ and Hydroside, Zq(OH)]. — Zinc oxide is formed
Trhen zinc burns in the air< It is known from its general appearance
as phUosophefs wooL It is prepared more conveniently by heating
basic zinc carbonate. This loses carbon dioxide and water, and
zinc oxide remains behind. Zinc oxide is used as a pigment
under the name of zinc uMie. Zinc oxide is almost insoluble in
water, so that zinc hydroxide is not prepared by treating the oxide
with water.

Zinc hydroxide, Zn{OH)j, is formed when a soluble zinc salt is
treated with a small amount of an alkaline hydroxide: ^ —

ZnClt + 2 KOH = 2 KCl + Zn(OH)^

Zinc hydroxide is a white, floeculent precipitate, which dissolves
readily in acids, forming the corresponding salts. It is therefore

TUE MAGNESIUM GROUP 389

basic with respect to acids, and as it dissolves^ must dissociate into
zinc and hydroxyl ions, thus : —

Zn(OH), = Zn, OH, OH. (1)

Zinc hydroxide also dissolves in strong alkalies, forming salts with
these, in which the hydroxide plays the role of an acid. If zinc
hydroxide is treated with potassium hydroxide in excess, the white
precipitate dissolves, forming potassium zincate : —

Zn(OH), + 2 KOH = K,ZnO, + 2 H,0.

The zinc hydroxide, in the presence of a strong base, acts then as an
acid, and must dissociate so as to yield hydrogen ions : —

Zn(OH), = H, H, ZnOa. (2)

The kinds of ions into which certain compounds dissociate is condi-
tioned, then, by tJie nature of the substance into whose presence they are
brought. Thus, zinc hydroxide in the presence of hydrogen ions, dis-
sociates in the sense of equation (1) ; in the presence of hydroxyl
ions, in the sense of equation (2).

Zinc Chloride, ZnClj.HsO. — The chloride of zinc is readily formed
by dissolving zinc in hydrochloric acid and evaporating the solution
to crystallization. When the salt is heated to remove the water it
loses hydrochloric acid, and a basic chloride remains behind : —

ZnCl,4- H,0 = HCl + Zn(OH)Cl.

To obtain pure anhydrous zinc chloride, the salt with water of
crystallization must be heated in an atmosphere of dry, hydrochloric
acid gas.

Zinc chloride combines readily with water, and is, consequently,
a *•' mild, dehydrating agent." Like the bromide and iodide of zinc,
it is readily soluble in water, and is dissociated much less than the
chlorides of the alkalies or alkaline earths.

Zinc Sulphide, ZnS. — The sulphide of zinc has been referred to
as o(^curring in nature as zinc blende, and as being one of the most
important ores of zinc. It is formed as a white precipitate when-
ever hydrogen sulphide is conducted into a dilute solution of a
soluble zinc salt: —

ZnCl, + Hj^ = 2 HCl + ZnS.

The sulphide of zinc is soluble to some extent in strong acids,
i.e, in the presence of hydrogen ions. In a reaction like the above
where a strong acid is liberated, the precipitation of zinc sulphide is

390

PRINCIPLES OF INOEGJLNIC CHEMISTRY

not complete, unless the solution is very dilute, since a portion of it
is dissolved in the hydrochloric acid formed as the result of the re-
action. If sodinm acetate is added to the solution, the hydrochloric
acid may be regarded as reactiJig with this salt, forming sodium
chloride which remains dissociated into its ions, and acetic acid: —

H, CI + Na, CH.COO = Ka, 01 + CH3COOH.

Acetic acid is a very weak acid^ which is the same as to say that
it is very little dissociated. The hydrogen ions com bine with the

acetic acid anions, forming the molecLile of the acid, and thus we
remove the hydrogen ions from the solution and prevent them from
dissolving the zinc sulphide.

If the solution of the zinc salt is treated with a solution of an
alkaline sulphide, the zinc is completely precipitated, since no free
acid is formed as the result of the reaction.

ZnCl, + K^ = 2 KCl + ZnS.

Zinc Sulphate, ZnSO^.THaO. — The sulphate of zinc is obtained

when a bar of pure zinc is wrapped with a piece of platinum wire,
dissolved in pui^ sulphuric acid, and the solution evaporated to
crystallization. The salt is formed on a large scale by heating the
sijphide in the presence of oxygen : —

ZnS + 2 0a = ZnS04.

Like other sulphates containing a lai-ge number of molecules of
water of crystallization, it retains one molecule to a much higher
temperature than the remainder. When zinc sulpbate is heated
slightly above 100" it loses six molecules of water* The last
molecule is retained until a considerable higher temperature is
reached. On account of its color and conj position the salt is known
as ivhite vitrioL

Zinc Carbonate, ZnCOj, is an important ore of zinc* When
heated it decomposes into the oxide and carbon dioxide: —

ZnCOj, s= ZnO + CO,.

Zinc carbonate is formed when a soluble zinc salt is treated with
a solution of an alkaline carbonate. The precipitate formed under
these conditions is not the normal carbonate of zinc, but a basic
carbonate containing more or less hydroxyl groups, depending upon
the conditions under which it is precipitated.

This compound is of importance m connection with the quanti-
tative determination of zinc. The zinc is precipitated as the basic

THE MAGNESIUM GROUP 391

carbonate by means of potassium or sodium carbonate, the precipitate
filtered off, washed, and ignited. It decomposes into the oxide on
heating, and is weighed in this form.

USES OF ZINC IN PRIMARY BATTERIES

Zinc is used more frequently in constructing primary cells than
any other known element. This is due to the fact that of all the
very common metals zinc has the greatest power to send ions into
solution. Whenever a metal is immersed in water or in a solution
of one of its salts, it exerts a tension or pressure which tends to
drive ions off from the metal into the solution. This is known as
the solution-tension of the metal. That this tension exists has been
demonstrated beyond question in the case of mercury — the only
metal liquid at ordinary temperatures.

Demonstration of the Solution-tension of Metals. — A demonstra-
tion of the solution-tension of metals has been furnished by Palmaer.
Mercury is a metal whose solution-tension is very small. Even
when in contact with a very dilute solution of a mercury salt, the
solution-tension of the mercury is less than the osmotic pressure of
the mercury ions in the solution ; and some of the mercury ions will
separate from such a solution.

Given a vessel whose bottom is covered with metallic mercury,
and over this is placed a solution of mercurous nitrate having a
volume of 2000. A few mercury ions will separate from the solution
and give up their positive charges to the mercury. The positively
charged mercury will attract electrostatically a few negative nitric

ions, NO3, to form the double layer. This will be continued until a
certain difference in potential has been reached, when equilibrium will
be established. If a drop of mercury is now let fall into the solution,
a few mercury ions will separate upon it, charge it positively, and it

will then attract an equal number of negative nitric ions, NOg, and
drag them down with it through the solution. The next drop of
mercury will behave in exactly the same manner, and thus the top
of the solution will become continually poorer and poorer in the salt.
When the drop of mercury comes in contact with the mercury at
the bottom of the vessel where equilibrium is already established,
what will happen ? When the drop has united with the mercury,
this will contain an excess of positive electricity, and, therefore, a
small quantity of mercury ions will pass into solution. And, indeed,
exactly the same number as there are nitric ions, KO^, brought down

392

PRINCIPLES OF INORGANIC CHEMISTRY

Jl

from the top to the bottom of the solution. The solution will thus
become more coucentrated just above the layer of luercury ou the
bottom of the vessel.

A fiue glass tube from whic^h mercuTj flows is known as a droi>-
eleetrode* To produce changes in concentration sufficient for the
purposes of a demonstration, a very powerful drop*eIectrode must be
used. This is in axle by inserting a conical glass stopper into a conical
glass tul)e, so that the jutictiou is mercury -tight. A large nuinl>er
of fine grooves are then etched on the outside of the stopi>er, so that
the mercury will stream through as a fine mist To assist this
process the mercury is subjected to four or five atmo&pheres of
pressure.

Under these conditions, however, the mercury cannot be allowed
to flow directly into a vessel filled with a dilute solution of a mer-
cury salt, and containing mercury at the bottom, sinee there would
be too much commotion in the solution. The
arrangement which was used is shown in
Figt 37. The drop-electrode T dips into the
funnel-shapetl vessel 0, which is connected
by a narrow tube and a rubber tube with tlie
larger vessel 3L This is in turn coutiected
with the vessel Uj where the change in
concentration ean be observed* When the
mercury has been allowed to flow for five
minutes under a pressure of five atmos-
pheres, distinct changes in concentration ean
be detected,

Fahnaer gives data which @how that the
concentration above had been diminished as
much as fifty per cent, and increased below
as much as forty per cent.

This will L>6 recognized at once to be a
very remarkable experiment, and before our
modern |>hysical chemical theories were pro-
posed would have been entirely inexplicable. The results of this
experiment were predicted before the experiment was tried.

The Eelative Solution-tensioni of Soma of the More Common
Metals* — It wuuhl lead us much too far to discuss here the
mtjthod of determining the relative solution-tensions of metals.
To study the principle and method some work on pliysical chem-
istry must be consulted* A few of the results obtained are the
following : —

Fio. 3?.

THE MAGNESIUM GROUP 393

Mktal

SOLmOlC-TKMSION
IW

ATMosruKBim

Magnesium

10**

Zinc

10"

Alaminium

10"

Cadmium

3xl0«

Iron

10* .

Cobalt

ax 100

Nickel

1x100

Lead

io-»

Mercury

10-16

Silver

10-"

Copper

io-»

Magnesium and alaminium as well as zinc have high solution-
tensions, or a great power to send off ions into solution, but these
are far less abundant and more expensive substances than metallic
zinc.

Solution-tension of Metals and Primary Cells. — The question
which arises is, What has the solution-tension of metals, or their
power to throw ions off into the solution, to do with the production
of electricity in a primary cell ? Why is zinc used in primary cells
because it has a high solution-tension, or stands near the top of the
tension-series, as the above order of the metals is termed ?

We have seen that all dissolved substances exert an osmotic
pressure. When a bar of zinc is dipped into a solution of a zinc
salt, say zinc chloride, the osmotic pressure of the zinc ions in the
solution tends to drive these ions out of the solution on to the bar
of metal, in the atomic condition. In order that an ion may become
an atom it must give up its electrical charge to something, and in
this case the only thing to which it can give it up is the bar of zinc.

The solution-tension of the metal exerts exactly the opposite in-
fluence. It tends to drive atoms of metal off from the bar into the
solution as ions. What will happen when a bar of metal is immersed
in a solution of one of its salts, depends upon which of these forces
is the greater. Let us call the solution-tension of the metal P, and
the osmotic pressure of the cations in the solution p. If P>p the
metal will send ions into the solution, and in doing so will become
negatively cJiargedy since it gives up electricity to convert atoms into
ions. If, on the other hand, P<p ions will separate out of the solu-
tion on to the bar as atoms, and in doing so will give up their charge

894

PRINCIPLES OF INORGANIC CHEMISTRY

to the bar, which will become positively charged. If P = p nothing
will happen when the metal is immersed in the solution of its salt.

The object, then, in using a metal with a high solution-tension
on one side of the cell is to have this electrode charged negatively
with respect to the surrounding solution.

With these conceptions of the relation between osmotic pressure
and solution-tension of the metals we can easily understand the
.action of a primary cell.

Concentration Element. — The simplest form of primary cell in
vwhich zinc is employed is the following: A bar of zinc, B (Fig. 38),
lis immersed in a solution of zinc chloride of a definite concentration,
.-say one-tenth normal. At any concentration of solution of a zinc
i0idt P>Pj since P is so very great. The bar of zinc will send ions

Fio. 38.

into the solution until a definite difference in potential is established
between the metal and the solution, the zinc being negative.

On the other side of the cell a bar of zinc, 5x, is immersed in a
solution of zinc chloride of a different concentration, say one-hun-
dredth normal. Here the zinc will throw ions into the solution and
will do so to a greater extent than the other bar, since the osmotic
pressure of the zinc ions in the solution, which is the counter force,
is less. The bar will become negative with respect to the solution,
and still more negative than the first bar. Let the two sides of the
cell be connected. A current will flow from the bar B to the bar
Bij since the latter is negative with respect to the former. Zinc ions
will continue to separate on B from the solution of the zinc salt
around it, and to dissolve from Bi as the current flows ; the chlorine

THE MAGNESIUM GROUP

395

anions moving against the current in the solution from B to Bi,
This cell is called a concentration element, since its whole action
depends on the difference in the concentration of the two solutions
of the same electrolyte which are used. During the flow of the cur-
rent the more concentrated solution becomes more dilute and the
more dilute solution more concentrated, and the cell will cease to
be active when the two concentrations have become equal.

The electromotive force of such an element is calculated from an
equation derived very easily from the fundamental equation: —

^ IT = 0.058 log^\

Pi is the osmotic pressure of the zinc ions in the more concentrated
solution of zinc chloride, and p2 the osmotic pressure of the zinc ions
in the more dilute solution of zinc chloride.

If the action of this, the simplest form of primary cell, is under-
stood, we are in a position to understand easily the action of any
form of primary battery.

The Daniell Cell. — We shall discuss briefly one very common
form of cell in which zinc is used as one electrode — the Daniell
cell. This is sketched diagrammatically in Fig. 39. It consists of a
bar of zinc immersed in a
solution of zinc sulphate on
one side, and a bar of cop-
per surrounded by a solu-
tion of copper sulphate on
the other. While zinc is a
metal with very high solu-
tion-tension, copper has a
very small solution-tension.
A bar of zinc immersed in
any solution of a zinc salt
will send ions into solu-
tion and become negatively
charged. A bar of copper
immersed in almost any
solution of a copper salt will receive ions from the solution and
become positively charged. W^hen the two electrodes are connected,
and the two electrolytes connected by a siphon filled either with
the zinc sulphate or the copper sulphate, we have a Daniell cell.

ZnS04

Fig. 39.

Cu SO4

1 It would lead too far to derive even this equation here,
work on physical chemistry.

Consult some

896 PRINCIPLES OF INORGANIC CHEMISTRY

The zinc passes into solution, and the copper separates from the
solution. The zinc is, therefore, negative and the copper positive,
and the current flows in the direction shown by the arrows. Copper
is used because it has a very low solution-tension and occurs in
abundance.

The electromotive force of this element is calculated by means
of the equation —

,r = .0291og--.0291og-*
P Pi

in which p and pi are the osmotic pressures of the zinc and copper
ions respectively, and P and Pi the solution-tensions of zinc and
copi>er. The solution-tensions of the two metals come into play
in the calculation of the electromotive force of the Daniell cell,
since they have different values. In the calculation of the electro-
motive force of the concentration element, the solution-tension of
the zinc is the same on both sides of the cell, and, consequently,
enters the equation twice, but with equal value and opposite sign.
It, therefore, entirely disappears from the equation.

To show that zinc is used in almost all of the best-known forms
of primary cells it is only necessary to mention the Grove, Leclanch^,
and bichromate cells.

CADMIUM (At. Wt. = 112.4)

A comparatively rare element which is closely allied in its prop-
■ erties to zinc is cadmium. It usually occurs associated with zinc

I ores, either in the form of the oxide or sulphide. The oxide of cad-

mium, like that of zinc, is easily reduced by heated carbon : —

I CdO + C = CO-hCd.

The metal boils at 778® and, therefore, distils over before the zinc.

The vapor-density of cadmium shows that the molecule jconsists of

I one atom. There are a few cases where the molecule of an element

; is monatomic, but very few, and most of them are in this group of

elements.

Cadmium unites with oxygen at an elevated temperature, forming
cadmium oxide CdO. This is a brown powder, as obtained by direct
! combination of the elements, and also by decomposing the carbonate

by heat. When the nitrate is heated the residue is cadmium oxide,
but this is very dark brown, or brownish-black. Cadmium hydroxide
is formed when a soluble salt of cadmium is treated with an alkaline
hydroxide : —
' CdCNOj), -h 2 NaOH = 2 NaNO, + Cd(OH)^

THE MAGNESIUM GROUP 397

Cadniiiim hydroxide is not soluble in an excess of the alkaline
hydroxide, showing that the acid properties which are manifested by
zinc hydroxide are lost in cadmium with the higher atomic weight.
While cadmium under ordinary conditions forms only the ion Cd,
yet it is capable of forming compounds in which it acts as a univa-
lent ion. The suboxide of cadmium, Cd^O, has been prepared. Also
the subhydroxide, CdOH.

Salts of Cadminm. — The cadmium ion Cd unites with the anions
of acids, forming salts which closely resemble the corresponding com-
pounds of zinc.

The cMoride, CdClj, 2H2O, is readily formed by the action of
hydrochloric acid on the metal or on the carbonate. Its chief inter-
est from a physical chemical standpoint is in connection with its
dissociation. While the halides of zinc are less dissociated than those
of the alkalies and alkaline earths, the halides of cadmium ai-e much
less dissociated than those of zinc. Of these, cadmium bromide is
dissociated less than the chloride, and the iodide less than the bromide.

Cadmium chloride, like the chloride of zinc, combines with the
chlorides of the alkalies and alkaline earths, forming such compounds
as the following : —

KCLCdCl^ 2 KCl.CdCl„ CaCl^CdCVKBr.CdBr^ BaBr^CdBrj^

These are the so-called double halides, of which many hundred
examples are known.

Cadmium sulphide, CdS, occurs in nature as greenockUe and is
formed when hydrogen sulphide is passed into a solution of a cad-
mium salt: —

CdCl, -h H,S = 2 HCl + CdS.

This is a beautiful, yellow precipitate, and is soluble to some ex-
tent in strong acids. Although soluble in strong acids it is much
less soluble in acids than zinc sulphide, and is thrown down nearly
quantitatively from its neutral salts by hydrogen sulphide. By an
alkaline sulphide cadmium is precipitated quantitatively, and this is
one of the methods of determining the amount of cadmium present.
On account of its fine yellow color it is used as a pigment by artists,
under the name of " cadmium."

The sulphate containing seven molecules of water, CdS04. 7 HjO, is
known. The sulphate which is ordinarily formed, however, has the
composition 3CdS04, SHjO.

The carbonate of cadmium, CdCOj, is formed when soluble cad-
mium salts are treated with soluble carbonates : —

CdCl, -f Na,COa = 2 NaCl + CdCO^.

398

PRINCIPLES OF INORGANIC CHEMISTRY

The compound is of importance because it is the form in which
cadmium is frequently precipitated in analysis.

Cadmium belongs to that group of elements whose sulphides
are precipitated from their neutral salts by hydrogen sulphide, and
whose sulphides are insoluble in ammonium sulphide.

MERCURY (At. Wt. = 200.0)

Mercury occurs in nature in the elementary condition, but much
more abundantly in the form of the sulphide, HgS, cinnabar. The

chief localities are Almadeu in Spain,
California, and Hungary. Mercury is
readily obtained by heating the sul-
phide in contact with the air. The
metal distils over and is condensed,
and the sulphur is oxidized to sulphur
dioxide.

Mercury is purified by passing it in
fine drops through a solution of ferric
chloride, through nitric acid, or by shak-
ing with sulphuric acid containing a little
chromic acid. The impurities are oxi-
dized while the mercury is unattaeked.
The apparatus used to purify mercury-
is shown in Fig. 40. The mercury flows
through the funnel, whose end is fused
so nearly together that the drops of mer-
cury which pass through are very fine.
The tube contains the purifying liquid.
The mercury is collected in the flask.

Properties of Mercury. — Mercury
combines with oxygen only at elevated
'°' ' temperatures, but combines with the

halogens at ordinary temperatures. Mercury forms a univalent

+ ++

mercurous ion, Hg, and a bivalent mercuric ion, Hg. Both of these

ions combine with the anions of acids, forming salts. The former

are known as " mercurous," the latter as " mercuric " salts.

Mercury is distinguished by the fact that it is the only metal

known which is liquid at ordinary temperatures. It solidifies at

— 39°.4. The specific gravity of mercury is 13.595. On account of

its low melting-point and its high density it is very valuable in the

preparation of barometers, thermometers, and other physical and

chemical apparatus. If mercury were not a liquid at ordinary

THE MAGNESIUM GROUP 899

temperatures the science of chemistry would undergo some funda-
mental changes.

Mercury has an appreciable vapor-tension at temperatures not
very far above the ordinary, and since its vapor is quite poisonous
care must be exercised in dealing with it. It boils at 358°, and
passes into a vapor whose density is 99.4 in terms of hydrogen as
unity. Its molecular weight is, therefore, 198.8, which is identical
with the atomic weight. The cUam and molecule of mercury in
the form of vapor are identical, and this is the fourth member of this
group where this relation obtains.

Amalgams. — Many of the metals dissolve readily in mercury,
forming amalgams, which are really solutions of the metals in
mercury as a solvent. Such metals as magnesium, zinc, cadmium,
silver, gold, the alkaline earths, the alkalies, and many others dis-
solve without serious difficulty in mercury. The amalgams of sodium
have already been referred to. The one containing one per cent of
sodium is a liquid, while double the amount of sodium gives a solid
solution in mercury. Such compounds as Hg«Na and HgNa, have
been obtained.

The ammonium amalgam formed by bringing ammonium chloride
in contact with sodium amalgam has also been referred to (p. 207).
These amalgams are often useful in electrochemical investigations,
where it is found to be more convenient to use the amalgam than the
pure metal.

Molecular Weights of Metals in Mercury. — The question naturally
arises. What is the molecular weight of the metal in question dis-
solved iu mercury ? There are two methods available for throwing
light on this question; the freezing-point, and the boiling-point or
the vapor-tension method. The lowering of the vapor-tension is
proportional to the rise in boiling-point, and the one phenomenon can
be used as well as the other for determining molecular weights.

The results obtained for the molecular weights of a few metals in
mercury, as determined by the lowering of the freezing-point of the
mercury, are as follows : —

MoL. Wt.

At. Wt.

Potassium
Sodium
Tliallium
Zinc

26-55

21-25

141-221

61-66

200
65

400

PRINCIPLES OF INORGANIC CHEMISTRY

The mean of the values found experimentally shows that for these
four elements the molecular weights in mercury are the same as the
atomic weigh ts, or the molecules are monatomic.

The results found by the lowering of the vapor-tension of mercury
are even more interesting. A few are given below : —

MoL. Wt.

At. Wt.

Lithium

7.10

7.02

Calcium

19.1

40.1

Barium

76.7

137.0

Magnesium

24-21.6

24.3

Manganese

66.6

66.0

In these and in all other cases investigated the molecular weight
was never greater than the atomic weight. In some cases, as with
calcium and barium, the molecular weight is about one-half of the
atomic weight. There is other evidence that the supposed atom of
calcium can be broken down into two or more parts. This evidence
is based on the shift in the position in the spectrum of the calcium
lines under pressure, but cannot be discussed here.

Mercurous (Hg,0) and Mercuric (HgO) Oxides. — Mercurous
oxide is formed when a mercurous salt is treated with an alkali : —

2 HgCl -h 2 NaOH = 2 NaCl -h H,0 -f Hg,0.

By light, heat, or friction, it is decomposed into mercury and
oxygen. It is a black powder and obviously unstable.

Mercuric oxide is formed when mercury is heated for a long time
in the air. It is also obtained as a red powder by heating the nitrate.

It is also formed by treating a mercuric salt with sodium
hydroxide : —

HgCl, + 2 NaOH = 2 NaCl + H,0 4- HgO.

Prepared in this way it is yellow, but becomes red on gentle heating.
The difference in color in this case seems to be purely mechanical,
depending on the size of the particles.

Mercurous (HgCl) and Mercuric (HgCl,) Chlorides. — Mercurous
chloride is the familiar substance calomel. It is formed by heating a
mercurous salt with a soluble chloride : —

Hg2S04 4- 2 NaCl = Na,S04 4- 2 HgCl.
Calomel is also formed by subliming a mixture of mercury and mer-
curic chloride : - j^^^^ -f Hg = 2 HgCl.

THE MAGNESIUM GROUP 401

It is also formed by reducing mercuric chloride with such a reducing
agent as sulphur dioxide : —

2 HgCl, 4- 2 H,0 -h SO, = H,S04 + 2 HCl + 2 HgCl.

Mercurous chloride is a white powder which can easily be ob-
tained in crystals by sublimation. It is difficultly soluble in water,
and is therefore taken into the system slowly when used as a medi-
cine.

Mercurous chloride is readily oxidized to the mercuric salt, which
is quite poisonous. This oxidation is effected by nitric acid, and the
same transformation is effected by hydrochloric acid and the alka-
line chlorides. Mercurous chloride is partly transformed into mer-
curic chloride by the action of light. When calomel is exposed for a
considerable time to the action of light it contains the very poison-
ous substance, mercuric chloride, and should never be used as a medi-
cine. Calomel, which is to be used for such purposes, should always
be carefully protected from all agents which transform it into mer-
curic chloride. The vapor-density shows that the molecule of
mercurous chloride is HgCl and not HgjClj. Mercuric chloride, or
corrosive sublimate, is formed by subliming a mixture of sodium
chloride and mercuric sulphate. The following reaction takes
place : — ^ ^^^^ _^ jj^g^^ ^ Na,S04 + HgCl,.

Mercuric chloride is a white, crystalline compound, readily soluble

in water, and still more easily soluble in a mixture of alcohol and

ether. It boils at 307®, and its vapor-density shows that it has the

formula HgCV On account of its solubility it furnishes a conven-

+ +
lent means of obtaining mercuric ions, Hg, which are very poisonous

to most forms of life. Mercuric chloride is, therefore, an excellent
disinfectant, and is extensively used in surgery for this purpose.

Mercuric chloride, like the salts of mercury in general, is only slightly
dissociated iiito ions. We saw that the compounds of zinc with the
halogens are far less dissociated than the corresponding compounds
of the alkalies and alkaline earths. Passing to the next member of
this series in the order of increasing atomic weights, cadmium, we
saw that its halogen compounds were dissociated less than those of
zinc. We come now to mercury, the next member of the series, and
find that its halides are dissociated relatively very little. The con-
ductivity of mercuric chloride in water is very small indeed. Mer-
curic chloride has the power to combine with the halides of the alka-
lies and alkaline earths, forming so-called double chlorides. These
have the composition MHgCls and MsHgCl^, where M corresponds to
2d

402 PRINCIPLES OF INORGANIC CHEMISTRY

a univalent alkali metal, and MiHgCl^, where M^ is an alkaline earth
metal. These can be regarded as compomids of complex acids of
the composition HjHgCla and HjHgCl^ and are definite chemical com-
pounds, as has been shown by the way in which they dissociate in
the presence of water. When the double halides of mercury and
especially of cadmium are electrolyzed, the mercury or cadmium
passes in part to the anode, showing that there exists a complex
anion composed of the less alkaline metal and the halogen. This is
the anion of the complex halogen acid of which the above compounds
are salts.

Mercuric chloride is easily reduced to mercurous chloride, just
as the mercurous chloride is easily oxidized to mercuric. The
reducing action of sulphur dioxide has already been considered in
connection with the preparation of mercurous chloride.

Oxalic acid reduces mercuric chloride in the presence of light,
and this reaction has been made use of to measure the intensity of
the light, as in photometry: —

2 HgCl, -h HjC A = 2 HCl -f 2 CO, + 2 HgCl.

Mercuric chloride is also reduced by stannous chloride, first to
mercurous chloride: —

2 HgCl, -h 6nCl, = SnCl4 + 2 HgCl ;

and if enough stannous chloride is present, to metallic mercury : —

HgClj + SnClj = SnCU + Hg.

Merooric Bromide (HgBrs) and Iodide (Hgis). — The bromide of
mercury resembles strongly the chloride. The iodide presents cer-
tain features of special interest. It is formed by rubbing together
mercury and iodine in the proper proportions ; if too much mercury
is used the mercurous iodide will be formed. It is prepared most
conveniently by adding potassium iodide to a solution of a mercuric

^^^^ • — Hg (CI), + 2 KI = 2 KCl -h Hgl,.

It is easily soluble in alcohol, and crystallizes from the alcoholic
solution in beautiful, scarlet-red crystals, which belong to the te-
tragonal system. When the red modification is heated above 126° it
turns yellow, and when more highly heated, melts and forms a yel-
low liquid, which, at a still higher temperature sublimes and forms
yellow crystals belonging to the orthorhombic system. W^hen the
yellow modification cools below 126° it passes again into the red.
If placed in a position where it is not subjected to mechanical dis*

THE MAGNESIUM GROUP 408

turbance the yellow form may exist below 126®. We have here an
eiiantiotropic substance existing in two forms which are mutually
transformable — the transition temperature being 126®.

When mercuric iodide is precipitated at low temperatures the
yellow modification is formed. After a time this passes over into
the red. When either modification is volatilized the vapors con-
dense as the yellow modification. The action of light is to acceler-
ate the transformation from the yellow to the red modification.
Mechanical disturbance causes the yellow form to pass into the red.
If a vessel containing the yellow modification is struck a few times,
the red modification will begin to appear.

Mercuric iodide dissolves readily in a solution of potassium
iodide, forming the double iodide K^Hgl^ : —

HgI,-h2KI = K,HgI,.

This may be regarded as the potassium salt of the complex acid
H2Hgl4. Solutions of this salt do not show the ordinary reactions
of mercury, nor is there any reason for supposing that they would
do so. The mercury is combined with four iodine atoms, forming

the complex anion Hgl4, and in this the mercury is doubtless play-
ing a very different role with respect to energy relations than when
alone. When caustic potash is added to the solution of potassium
mercuric iodide we have Nessler^s reagent, which is a very sensitive
means of testing for ammonia ; minute traces of ammonia forming
a characteristic, yellowish-brown color in a solution of this reagent.
This is due, as we shall see, to the formation of complex compounds
of ammonia and mercury, some of which have very characteristic
colors.

Mercuric Sulphide, HgS. — Only one compound of mercury with
sulphur is known, and this is the one in which the mercury is
bivalent. When hydrogen sulphide is brought into the presence of
a mercurous salt, a mixture of black mercuric sulphide and mercury
is thrown down : —

2 HgCl -f H^S = HgS -h Hg -h 2 HCl.

When mercury and siUphur are rubbed together the sulphide is
formed. It is also formed by passing hydrogen sulphide for a time
through a solution of a mercuric salt : —

Hgaj -f H,S = 2 HCl 4- HgS.

When hydrogen sulphide is passed through a solution of a
mercuric salt at first a white precipitate is formed. This is a

404

PRINCIPLES OF INOEGAKIC CHEMISTRY

compound of the ori^ual mercury salt with mercuric sulphide, such
as HgCla.HgS, Hg(Jlt.2HgS, and so on. By continuing to pass
in hydrogen sulphide the effect of mass comes intx> play, and such
complex compounds aa the above ate gradually broken down ; all of
the mercury being t ran a formed into mercuric sulphide. This is
shown by the graxlual transformation of the white precipitate into
gray, and finally black.

The crystallized form of mercuric sulphide, which occurs in
natui-e as dnnaharj is grayish-red in mass, but when powdered is
red. On account of ita beautiful color the artificial sulphide called
vermilion is used as a pigment.

The black modification is the unstable form gince it is produced
first, and because it passes over slowly into the red modification! ei
cially in the presence of alkali sulphides. Further, when the
modification is heated it becomes dark. If not heated to too high ;
temperature the dark modification becomes red again on cooling.

Mercury sulphide is an important substance, because mercury is
usually precipitated in this form in analjiieal operations. Mercuiy
sulphide does not dissolve in dilute nitric acid. It is thus distin-
guished from all the sulphides which are precipitated in the presence
of dilute atdds. These sulphides are wamied with dilute, nitric
acid, when all dissolve except the sulphidG of mercury.

Kercurouf (HgaSO^) and Hercuric (HgSO^) Sulphates, — Mercurous
sulphate is formed by the action of hot, concentrated sulphuric acid
on mercury : —

2Hg+H^0, = Hg,S0, + H^

Mercuroua sulphate is difficultly soluble in water, and is ex-
tensively used in preparing standard cells with constant electro-
motive force. Tlie normal element of Latimer Clark consists of ,
mercury covered with a thick paste of mercurous sulphate, whiclrl
serves as one electrode^ Above this a thick paste of zinc sulphatei
into which a Imr of zinc is immersed serves as the other electrode.
Such an element has an electromotive fore© of 1.4328 volts, —0,0012
(t — 15*), t being tlie temperature at which the element is used*

Another normal element in which mereurons sulphate is us^ li
the WeMon cadmium element* This consists of mercury covered
with a paste of mercurous sulphate^ and over this a paste of cadmium
sulphate into which a bar of cadmium dips. The electromotive
force of tins element is 1,0186 volts, and it has a very small temper^
ature coefficient.

THE MAGNESIUM GROUP 405

Mercuric sulphate is practically insoluble in water, but is decom-
posed by it, forming a basic salt. This is a yellow compound
having the composition HgS04.2 HgO.

Mercuric Cyanide, Hg(Cir)2, is formed when Prussian blue is
boiled with mercuric oxide. It is also formed when mercuric oxide
is dissolved in hydrocyanic acid. When heated it decomposes as
follows : — Hg(CN), = Hg + 2 CN.

This, as we have seen, is the most convenient method of prepar-
ing cyanogen. Mercuric cyanide combines readily with potassium
cyanide, forming the compound K2Hg(CN)4.

We have seen that the salts of mercury are in general only
slightly dissociated. Also that hydrocyanic acid is one of the very
weakest acids, and is almost undissociated. We would, therefore,
expect that mercuric cyanide would be very slightly dissociated,
and such is the fact. An aqueous solution of pure, mercuric cyanide
conducts the current only a little better than pure water.

This shows that we must not conclude that all salts are dis-
sociated and conduct the current because most of them do so. Here
is an excellent example of a salt which shows practically no dis-
sociation. It should, however, be added that there are only a few
analogous cases.

The mercury compound of an acid whose composition is the same
as that of cyanic acid, HOCN, is explosive. The compound has the
composition HgCgNgOj, and is known as fulminating mercury.

Action of Ammonia on Salts of Mercury. — When ammonia is added
to a mercurous salt a black precipitate is formed. When ammonia
is allowed to act on mercurous chloride the following reaction takes
place : —

2 NHg -h 2 HgCl = NH4CI 4- NH,Hg,Cl.

This is ammonium chloride in which two hydrogen atoms are
replaced by a mercury atom, and is known as mercurous chloramide
— the group NH2 being known as the amido group. It should be
stated that this substance is regarded by some as a mixture of
mercury and the compound NH,HgCl.

When mercurous chloride is treated with gaseous ammonia the
following reaction takes place : —

HgCl + NH« = HgNH,a

which is ammonium chloride, in which one hydrogen atom is replaced
by mercurous mercury. This is known as mercurous ammonium
chloride.

406 PRINCIPLES OF INORGANIC CHEMISTRY

When mercuric chloride is treated with ammonia the following
reaction takes place ; —

2 NH, + HgCl, = NH.Cl 4- HgNH,Cl.

This is ammonium chloride, in which two hydrogen atoms have
been replaced by mercuric mercury. From its properties, its color,
and the fact that it sublimes without melting, it is known as infusible
white precipitate.

If a solution of mercuric chloride is added to a boiling solution
of ammonium chloride in which there is some free ammonia, the
following reaction takes place : —

HgCl, + 2 NH3 =Hg(NH3Cl)^

This consists of two molecules of ammonium chloride, in each
of which one atom of hydrogen is replaced by one-half of a mercuric
mercury atom. It is, therefore, known as mercuric diammonium
chloride^ or since the salt readily melts, 2ls fusible white predpitcUe,

Variable Valence. — We have studied a number of non-metals
which showed different valence under different conditions. This is,
however, the first metal which we have encountered in which two
different valencies clearly manifest themselves. We have a well-
defined mercurous ion Hg, in which the mercury carries only one
electrical charge. This forms, with the anions of acids, a class of
salts which have definite properties. From these salts the mercurous
mercury is precipitated by bases as the black mercurous oxide.
Hydrochloric acid throws down insoluble mercurous chloride.
Hydrogen sulphide precipitates a mixture of mercuric sulphide
and mercury. ^^

The mercuric ion, Hg, has its own characteristic properties,
forming compounds also with the anions of acids. From these
compounds it is precipitated by hydrogen sulphide, as mercuric
sulphide. The alkaline hydroxides precipitate mercuric oxide,
while stannous chloride in excess throws down metallic mercury.

CHAPTER XXXII

THE EARTH METALS

ALUMINIUM AND THE RARE ELEMENTS, — SCANDIUM,
GALLIUM, YTTRIUM, INDIUM, LANTHANUM, YTTERBIUM,
THALLIUM, SAMARIUM

ALUMINIUM (At. Wt. = 27.1)

Occurrence and Preparation. — Of the elements of this group only
one occurs in any abundance, and this is aluminium. The remainder
are rare substances and will be considered briefly.

Aluminium is a very important constituent of the crust of the
earth (see p. 6). The oxide of aluminium, known as corundum, is
very abundant, and the precious stones ruby and sapphire are alumin-
ium oxide colored by a small amount of other substances.

The double silicates of aluminium and the alkalies are among
the most common minerals. Mica is a double silicate* of aluminium
and one of the alkalies, having the general composition MAlSiOf,
in which M is a univalent alkali. The more common feldspars are
silicates of aluminium and one of the alkalies, sodium or potassium.
They have the general composition MAlSijOa. Bauxite is a hydrox-
ide of aluminium containing iron. Kaolin and clay are more or less
pure hydrous silicates of aluminium, while cryolite, found in Green-
land, is a double fluoride of sodium and aluminium, having the com-
position 3 NaF . AIF3. This compound is of importance in connection
with the preparation of the element.

Aluminium, named from alum in which it occurs, was prepared
first by the great German chemist, Wohler, who heated the chlo-
ride with metallic sodium.

Aluminium is prepared to-day by the electrolytic method. The
compound electrolyzed is aluminium oxide. On account of the high
fusing-point of aluminium oxide, a bath of cryolite is used, and into
the fused cryolite the aluminium oxide is introduced as desired.
The cryolite is fused in iron crucibles, which are sometimes lined
with carbon. This serves as the cathode upon which the metal
separates ; the oxygen set free combines with the anode, which con-

407

408 PRINCIPLES OF INORGANIC CHEMISTRY

gists of bars of carbon introduced into the fused cryolite. The mass
is kept molten by the heat generated by the current. Since the
introduction of the electrolytic method of preparing aluminium,
this metal has become quite abundant and its price lessened many
hundred fold. This method, extensively applied at Pittsburg and
elsewhere, is known as the Hall method.

Properties of Alominiam. — Aluminium is a metal with remark-
able properties. It is ductile and malleable and can be readily drawn
into wire or hammered into thin foil. It is very light for a metal,
having a specific gravity of only 2.7. It can, nevertheless, withstand
considerable strain, but by no means as great as was supposed before
it was prepared on a large scale. Its softness also detracts from its
commercial value. It melts at about 700°, and can, therefore, be
readily moulded. It is an excellent conductor of both heat energy
and electrical energy.

Chemically, aluminium is fairly resistant. In contact with moist
air it becomes covered with a very thin layer of oxide, which pro-
tects the metal from further action. It does not act appreciably
upon water, even at elevated temperatures. It dissolves readily in
hydrochloric acid, while nitric and sulphuric acids act only at ele-
vated temperatures. It is readily attacked by alkalies, and this is
another defect commercially.

While aluminium does not combine readily with oxygen at
ordinary temperatures, it combines with great vigor at high tem-
peratures. At these elevated temperatures it is, therefore, an
excellent reducing agent. This reducing action is utilized commer-
cially for preparing certain elements, as well as for producing locally
very high temperatures, since an enormous amount of heat is pro-
duced when aluminium combines with oxygen. When a mixture
of finely divided aluminium and ferric oxide is heated by an ignited
magnesium wire, or by a primer consisting of a mixture of finely
powdered magnesium and barium dioxide, the mixture becomes
intensely hot and the aluminium takes the oxygen from the iron,
leaving the latter in the molten condition. The temperature of
3000** can be reached by means of this reaction. This is utilized to
heat iron bolts white hot, weld iron rails, and the like. The iron
from the reduced oxide remains behind as a fused mass.

By means of this same reaction chromium, manganese, and similar
elements can ho prepared from their oxides.

The (tUoifs of alKminium are frequently of considerable commer-
cial value. The alloy with copper known as aluminium bronze, con-
taining G to 8 per cent aluminium, is used for constructing certain

THE EARTH METALS 409

forms of scientific apparatus. The alloy containing from 12 to 20 per
cent of magnesium is known as magnaliam. It has a speci6c gravity
of from 2 to 2.5, and from its properties promises to be very useful.

Almnininm Amalgam. — Aluminium does not combine as readily
with mercury as many of the metals. If, however, a clean piece of
aluminium is introduced into a dilute solution of mercuiic chloride,
the amalgam quickly forms. The aluminium in the amalgam seems
to be very much more active chemically than ordinary aluminium,
this " active aluminium " decomposes water vigorously at ordinary
temperatures, forming aluminium hydroxide. The aluminium under
these conditions is really no more active than ordinary aluminium.
The difference is that a clean surface of the metal is exposed to the
water. The thin coating of oxide which covers the surface of the
metal is removed by the alloy being liquid, and a fresh surface is
continually exposed to the action of water or the oxygen of the^ air.

Aluminium Oxide (AlsOs) &&d Hydroxide (A1(0H)3). — The oxide
of aluminium occurs in nature, as has already been mentioned. Its
most common form is corundum. Sapphire is a blue variety used as
a gem, while ruby is a red sapphire. A violet variety is known
as oriental amethyst, and a yellow variety as oriental topaz.

Aluminium oxide is easily prepared by heating aluminium
hydroxide : - ^ Al(OH). = Al.O, + 3 HA

Metallic aluminium is prepared chiefly from the oxide, and espe-
cially from the variety containing iron, known as bauxite.

Aluminium hydroxide occurs in nature as hydrargillite. This sub-
stance minus one molecule of water is known as diuspore, HAIO^ It
is readily prepared by treating a soluble aluminium salt with a base: —

AICI3 -f 3 NH4OH = 3 NH,C1 4- Al (011)3.

Aluminium hydroxide is very slightly soluble in water, and is
a very weak base. It can, however, form tlrree classes of salts in
which one, two, and three acid anions, respectively, are present.
Thus, we may have A1(0H),A, A1(0H)A2, 'and AlAg, where A is a
univalent anion. The first two substances are basic salts. Since alu-
minium hydroxide is a triacid base, it must be capable of dissociating
as follows : — ^^^ _ _ ..

A1(0H)3 = Al, OH, OH, OH.

Aluminium hydroxide also acts as an acid towards strong bases.
When aluminium hydroxide is treated with a strong base like
sodium hydroxide, the following reaction takes place: —

A1(0H)8 -h 3 NaOH = 3 H,0 + Na^AlOs-

410

PRINCIPLES OF INORGANIC CHEMISTRY

This compound, known as mdltim o/uwina^e, ia obviously the sodium
salt of the acid HsAHV The question which arises here is, How \
can a eom pound which dissociates like aluminium hydroxide^ yield-
ing hydroxy! ions, have acid jjroperties or yield liydrogen ions ?
The dissociation of alaminium hydroxide depends upon the nature
of the substance with which it is in contact Wlien in contact with
an acid or hydrogen ions, it dissociates as indicated above, yielding
hydroxyl ions. When in contact with a base or hydroxyl ions> it
dissociates as follows : —

Al(0H)3 = H, H,A10^

HA10, = H, AlO^

In a word, it dissociates as a tribasic acid* We have seen a similar
phenomenon manifested by zinc, and other cases will appear later.

Alnminatea. — The compounds in which aluminium hydroxide
plays the role of an acid are known as aluminates. There are three
classes of these substances, as would be expected from the above de-
scribed dissociation of aluminium hydroxide. Thus, we would have
MH2AIO3, MsjHAirivt and M^AlOa, where M is a univalent ion. The
alkali salts are sohible in water, showing an alkaline inaction. This
is due to the hydrolysis of the aluminates by water, aluminium
hydroxide being a weak atud as well as a weak base. The potas-
sium salt most readily formed has the composition KHjAIOs minus
water, or KAlOj. This is the potassium salt of jnetaalumimc acid,
HAlOa, obtained from aluniinic acid by loss of water: —

Al(0H),= H,0+HA10a,

The alkaline earths also form aluminates. Those of barium are
soluble in water, while those of calcium are insoluble. Calcium
forms the normal aluminate Cag(A10,^)j, and also the nietaaluminate
Ca(A102)a. Since these salts harden in contact with water, they are
used in connection with the preparation of hj^draulic cemeuU.

Magnesium forma a metaaluminate, which is the well-known
mineral spineL It has the composition Mg(A10jV Oahnite is the
corresponding zinc salt, Zn(A10a)t, while chnfitciberyl is the corre-
sponding compound of glucinum, Gl{AlOs)2, and ferrous iron forms
the metaaluminate-^>?eonajrf, Fe (AlOj)^

We have a large number of minerals of the type of spinel, in
which the place of the aluminium is taken by mang&nesej Irouj etc
These will be referred to ^ain.

I

THE EARTH METALS 411

The natui-ally occurring aluminates are frequently very stable
compounds, requiring vigorous chemical reagents to decompose
them. This is of importance in connection with tlie analysis of
these minerals. They must be powdered very finely and fused with
acid potassium sulphate, sodium carbonate, etc., before they can be
gotten into solution.

Alnim'niiiin Chloride, AICI3, is formed when aluminium hydrox-
ide is treated with hydrochloric acid. From such a solution it
can be obtained in crystalline form with six molecules of water —
AICI3.6H2O.

The anhydrous chloride is obtained by heating aluminium filings
in a current of dry, hydrochloric acid gas. It is also obtained by
the action of chlorine on a mixture of aluminium oxide and carbon
heated to a high temperature : —

AlA-f-3C-f-3Cl, = 3CO + 2AlCl3.

When the hydrated salt is heated it loses hydrochloric acid, and
aluminium oxide remains behind : —

2 AICI3 + 3 HjO = 6 HCl -f- AI2O3.

Aluminium chloride is very hygroscopic, taking up moisture read-
ily from the air when brought in contact with it.

With chlorides of the alkalies it forms beautifully crystalline
double salts, which are quite stable. The sodium aluminium chlo-
ride volatilizes without decomposition.

Aluminium chloride is strongly hydrolyzed by water, aluminium
hydroxide being precipitated : —

AICI3 4- 3 H,0 = 3 HCl 4- Al (0H)3.

This can be prevented by adding hydrochloric acid to the solu-
tion. By sufficiently increasing the mass of this acid, the reverse
reaction involving the reformation of aluminium chloride can be
carried very nearly to the limit; so that there is only an infini-
tesimal amount of the hydroxide formed.

Aluminium chloride has the remarkable property of causing
hydrogen in one compound to combine with chlorine in another,
forming hydrochloric acid; the residues of the two substances
uniting and forming a new compound. This reaction, known as the
Friedel'Crafta reaction, is of great importance in organic chemistry
for effecting the synthesis of many well-known substances.

The molecular weight of aluminium chloride can be readily

■x:.-'- **

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'i liirt J.- ili« <ji.-! i«i<i;il t;<ij|/)it<l<' w)ii*}j wi; iiav<; ihiLS far encooii-

kinmmmn HulpbuU, Al,/l*0,,, UHA i« fomuMl by treating
iiliiiiiHiiiiiK liyilMfAhlii Will) r>iil|iliiiiMr it'-hl iiiid iMfatiiig the mixture.
()m I v.i|<ifi.ii)ii{/ llii: hifliilinn, ii hiill of dm jiUivi* roiiiposition se|>a-
*••''•» it 1.1 .il.)i« luiiiii'il liy I rial m^ v\\\y willi concentrated sul-
I'liiiiii: .41 hi iihtl [Mil il^ INK ilui prniliiil,. Aluinininni sulphate is

THE EARTH METALS 413

hydrolyzed by water, as we would expect, on account of aluminium
being such a weak base. The solution contains free hydrogen ions
and, therefore, reacts acid. Aluminium sulphate forms with sul-
phuric acid a basic sulphate in which only one hydroxyl of the
aluminium hydroxide has disappeared. This has the composition
(Al (011)2)2804.7 HjO. When we consider that aluminium hydrox-
ide is such a weak base and sulphuric acid a strong acid, it is sur-
prising that such a compound should exist. This basic sulphate
occurs in nature under the name of aluminite.

The Alums. — Aluminium sulphate combines with the sulphates
of the alkalies to a remarkable extent, forming a class of double sul-
phates known as the alums. These have the general composition,
MAI (804)2.121120, in which M is an alkali ion. There are a large
number of these substances ; indeed, every alkali sulphate forms an
alum with aluminium sulphate. The best known of these are potassium
alum, KAl (804)2. 12 H^O, and ammonium alum NH4AI (804)2. 12 H2O.
The alums all crystallize in the same system, as octahedra and cubes,
and are all isomorphous, i.e. will form crystals containing several of
these substances. When a crystal of one alum is suspended in a
solution of another alum, the second alum will be deposited upon it
as upon one of its own crystals.

The term alum has been extended from the double sulphates of
aluminium and the alkalies, to the double sulphates of allied ele-
ments and the alkalies. Thus, we have a series of iron alums of the
general composition, MFe (804)2- 12 H2O, in which M is sodium, potas-
sium, rubidium, caesium, lithium, or ammonium. 8imilarly, we have
a series of manganese alums, MiMn (804)2. 12 HgO, and a series of
chromium alums, MCr (804)2. 12 11,0. These all crystallize in the
regular system, contain twelve molecules of water of crystallization,
and are isomorphous with one another and with the corresponding
aluminium compounds.

When alum is heated it passes into solution in its water of crys-
tallization, and when heated higher loses its water, swells up, and
forms a light powder which is known as burnt alum.

Alum, which is easily prepared by bringing the two sulphates
together in solution and evaporating the solution to crystallization,
is used to-day very largely in sizing paper, and as a mordant in
dyeing. Aluminium compounds, as we have seen, are hydrolyzed to
some extent by water. The aluminium hydroxide formed unites
firmly with the fibre of the substance to be dyed, and also unites
with the dye. In this manner a mordant renders the object
permanently dyed.

414 PRINCIPLES OF INORGANIC CHEMISTRY

The sodium alum is much more soluble than the potassium or
ammonium alum, while the rubidium alum is much less soluble.

These complex sulphates dissociate in dilute solutions just like
the constituent sulphates. A dilute solution of an alum has, then,
the same properties as a mixture of the two sulphates. This is
shown by determining the conductivity of solutions of alum. It is
the same in dilute solution as the mixture of the two sulphates. In
concentrated solution the conductivity of the alum is less than that
of the mixed sulphates, showing that some of the complex ions
persist undecom posed in such solutions.

Alnmininm Carbide (AI4C3) and Carbonate. — Aluminium com-
bines with carbon, forming the carbide, AI4CV This was produced
by Moissan, by heating aluminium in carbon boats in an electric
furnace, also in preparing aluminium by the electrolytic method
where the metal separates upon carbon electrodes.

Aluminium carbide decomposes with water forming methane and
aluminium hydroxide: —

Al.C, 4- 12 H,0 = 4 Al (OH), -f- 3 CH^.

The carbonate of aluminium can exist only at low temperatures.
Even at ordinary teini)eratures it decomposes into the hydroxide and
carbon dioxide. When a soluble carbonate is added to a solution
of an aluminium salt, the hydroxide and not the carbonate is pre-
cipitated : —

2 AlCl, 4- 3 Na,C03 + 3 H,0 = 6 NaCl -f- 3 CO, -f- 2 Al (OH ^

Silicates of Alomininm. — These are very important substances.
Aluminium forms salts not only with normal silicic acid, but with
the polysilicic acids. A salt of the composition AljSiOa is known
as disthene. A comparatively pure form of aluminium silicate is
known as kaolin. This substance has approximately the composi-
tion Al4(Si04)3.4 H,0, being the aluminium salt of normal silicic
acid.

Clay is an impure variety of aluminium silicate. This is formed
as the result of the weathering of the rooks, and, consequently,
many other substances are liable to be present in it. The different
colors of clays are due to different impurities.

3farl is clay containing a large amount of calcium carbonate.

Aluminium silicate readily forms double silicates with the silicates
of the alkalies, and these constitute some of the most important
minerals, the feldspars. We have potassium feldspars, sodium feld-
spars, potassium sodium feldspars, etc. These have the genei-al

THE EARTH METALS 415

composition MAlSisOg. The potassium feldspar is known as ortho-
dose, the sodium compound as albite. The feldspars are continually
undergoing decomposition, affected by the moisture and carbon diox-
ide in the air, and they are the chief source of the soluble potassium
compounds in the soil. This is one important effect of that geologi-
cal process known as weathering which is going on over the surface
of the earth.

Aluminium silicate forms double silicates with many other ele-
ments, especially with those of the calcium and iron group. Among
the double silicates of aluminium are such minerals as the garnets,
mica, zeolites, etc.

Lapis lazuli is a double silicate of sodium and aluminium, con-
taining sulphur. It is a beautiful coloring matter known under the
name of ultramarine. This substance is also prepared artificially for
commercial purposes.

Applications of Alumininm Silicates. — When kaolin, clay, or marl
is mixed with water it forms a thick, viscous mass, which can be
readily moulded or worked into any desired form. When this mass
is dried and heated it becomes very hard and then resists the action
of water. It is from this material that ordinary brick or fire-brick
is made.

Earthenware or stoneware is made from a somewhat purer variety
of aluminium silicate than that employed in the manufacture of or-
dinary fire-brick. The objects are moulded from the purer varieties
of clay or kaolin by mixing with water to the proper consistency.
They are then "fired" or "baked" by heating to a high tem-
perature. The resulting objects are, however, very porous, and
would be comparatively useless in this form. They must be cov-
ered by a non-porous material — must be glazed. There are different
methods of glazing such objects. A mixture of kaolin and feldspar
melts at a comparatively low temperature, and this is sometimes ap-
plied to the surface of earthenware after it has been " burned," and
the object then reheated. This method is seldom applied to ordi-
nary earthenware. After the earthenware has been burned, sodium
chloride is thrown into the furnace. This is decomposed at the high
temperature by the water-vapor, and forms hydrochloric acid which
escapes, and sodium hydroxide — it is hydrolyzed. The sodium
hydroxide then acts on the silicate of aluminium, forming the
double silicate of sodium and aluminium, which fuses and forms a
glassy, impervious coating over the porous earthenware.

Porcelain is made of pure aluminium silicate or pure kaolin.
This is mixed in definite proportions with feldspar to lower its melt^

416 PRINCIPLES OF INORGANIC CHEMISTRY

ing-point, and also with quartz. The mixture is treated with water
until the desired consistency is reached, and is then moulded into the
desired form. It is then dried and "burned." The glaze consists of
a mixture of feldspar, quartz, and lime. After applying the glaze
the vessel is heated again to a very high temperature. The many
.different varieties of porcelain owe their peculiar characteristics in
•jpart to the nature of the materials used, and in part to the way in
w!hieh they are manipulated during manufacture.

Aluminium silicate is also an important constituent of many val-
iUable cements,

^Detection of Aluminium. — Aluminium belongs to that class of ele-
ments whose hydroxides are precipitated from an alkaline solution
by hydrogen sulphide, or whose hydroxides are precipitated from a
neutral solution by ammonium sulphide.

SCANDIUM (At Wt.=44.1)

The special interest connected with the element scandium has to
do with its relation to the Periodic System of the elements. Men-
del^eff, iu 18G9, predicted the existence of an element with an atomic
weight of 44, occupying a position in the Periodic System next to
boron in group III. He termed the element ekahoron. This ele-
ment was discovered by the Swedish chemist, Nilson, in 1879, in
certain Norwegian minerals, — gadoUnite, euxenite, etc., — and named
scandium from the locality in which it was found.

The properties of the element, and especially of its compounds,
were predicted in detail by Mendel^eff, and his predictions have been
verified to a surprising extent.

Scandium forms the oxide ScjOj and the hydroxide Sc(0H)5.
The nitrate has the composition Sc(N03)3, and the sulphate
^€2(804)3. 6 HjO. Scandium forms a double sulphate with the sul-
phate of potassium, having the composition 3K,S04.Sc2(S04),.

GALLIUM (At. Wt. = 70.0)

Gallium has the same interest in connection with the Periodic
System as scandium. Its existence and properties were predicted by
Mendeleeff in 1869. He placed it in his system next to aluminium
and termed it ekaaluminium. It was discovered by Lecoq de Bois-
baudran in 1875, in certain zinc blendes which occur at Pierrefitte,
in France, and named for the country from which it came. It occurs
in extremely small quantities in these ores, and its presence was

THE EARTH METALS 417

detected by means of the spectroscope. It forms compounds of the
general type GaA,, where A is a univalent anion. Thus, we have
Ga(N03)3, GaCls, Ga,(S04)s, Ga^Oj, and so on. Gallium can also
form gallous compounds — GaCl,.

YTTRIUM (At. Wt. = 89.0)

Yttrium also occurs in the Norwegian minerals, gadolinite, yttria-
lite, euxetu'te, etc., and in monazite sand. It resembles the other
members of the group, forming compounds of the general type YAj,
A being a univalent anion. Thus we have Y(N08)a, Y,(S04)a, YjOj,
and so on.

INDIUM (At. Wt. = 115.0)

Indium occurs in certain zinc blendes in Freiberg, but in very
small quantities. It forms compounds which are analogous to those
of aluminium. Thus we have, InClj, Iu(N0j,)8, In,(S04)a, Iii(OH)j,

InA.

LANTHANUM (At. Wt = 138.9)

Lanthanum occurs in cerite, samarskitef monazite sand, etc. It
forms LajOs, La(N0a)8, LaClj, La,(S04)3. It has been prepared as
the double nitrate with ammonium in considerable quantity in con-
nection with the manufacture of Welsbach lights. The Welsbach
Light Co. of Gloucester, New Jersey, owns several hundred pounds
of the almost pure nitrate of lanthanum and ammonium.

YTTERBRIUM (At. Wt = 178.0)

Ytterbrium is found especially in the mineral euxenite. It forms
the oxide YtjOs, and the salts have the same general composition as
those of the other rare elements just considered.

THALLIUM (At Wt =204.1)

Thallium occurs in a number of minerals, but always in limited
quantities. Its chief source is certain zinc blendes. When these
are roasted the thallium passes off and is deposited in the dust in
the flues, and it was here that it was first discovered by Crookes in
1861, by means of the spectroscope. Its spectrum is characterized
by a bright green line, whence the name of the element.

Thallium forms two classes of compounds — the thallous and the
thallic. In the former the thallium is univalent, in the latter triva-
lent. Among the thallous salts are the chloride TlCl, the bromide
2b

418 PRINCIPLES OF INORGANIC CHEMISTRY

TlBr, the iodide Til, and the sulphide T1,S. Thallous chloride is
difficultly soluble in water, resembling in this particular the element
lead. The bromide is less soluble, and the iodide the least soluble
of the three. Among the thallic compounds are the chloride T1C1»
the carbonate Tl2(COs)3, the sulphide Tl^,, and so on. Thallous
thallium resembles in many respects potassium, forming salts which
are often isomorphous with the corresponding potassium compounds.
Thallic thallium resembles aluminium and the rare elements which
we have just been considering.

SAMARIUM (At. Wt = ISaS)

The position of this element in the Periodic System is not yet
fixed. It occurs in the mineral thorite^ and in some respects
resembles aluminium. It forms the oxide Sa,0„ which is a weak
base, dissolving in acids. Its salts have a sweet taste.

CHAPTER XXXIII

IRON, COBALT, NICKEL, MANGANESB, CHROBdUM, MO-
LTBDANUM, TUNQSTEN, URANIUM

IRON (At. Wt = 55.9)

We now come to a group which contains some of the most impor-
tant elements technically, as well as some of the most interesting
from the chemical standpoint. The first of these, and the one from
which the group takes its name, is iron — the most important tech-
nically of all the elements.

Occurrence and Preparation. — Iron occurs very widely distrib-
uted in nature, but not in abundance in the free state. This is due
especially to its action on water forming the hydroxide. Iron, how-
ever, occurs in the uncombined condition in certain meteorites and
in certain localities, as at Of vivak, Greenland.
Iron occurs chiefly in the form of oxides and
sulphides. It occurs in large quantities as
magnetite^ Fe304, as hematite^ FcjOg, and as
hog-iron orCy limonite, and other hydroxides.
Iron also occurs in large quantities as pyrites^
FeSj, and as the carbonate, FeCOg, or siderite.

Iron is prepared by reduction of its
oxides by carbon : —

FeA4-3C = 3C04-2Fe,
Fe3044-4C = 4CO-f-3Fe.

In preparing iron on the commercial scale
the '* blast furnace " is employed. This is
shown in Fig. 41. The furnace consists of
an iron case lined on the inside with fire-
brick, and has the shape shown in the figure.
It is provided with pipes at the bottom for
introducing hot air under pressure, and with
pipes at the top for carrying off the gaseous products of combustion.
These furnaces are often quite large, being as much as eighty feet

410

Fio. 41.

420 PRINCIPLES OF INORGANIC CHEMISTRY

high. They are filled with coal or coke, the ore, and a flux, which
are mixed when introduced into the furnace. The nature of the
flux, which is used to protect the metal when formed, depends upon
the impurities contained in the ore. If there is much silica in the
ore, limestone is used as the flux. If there is much calcium or mag-
nesium in the ore, a flux containing silica (sand) or aluminium (feld-
spar) is employed. Limestone, however, is almost always used as
the flux in the blast furnace.

The oxide of iron is reduced to the metal by means of the highly
heated carbon and the carbon monoxide formed as the product of the
combustion of the carbon. The combustion of the carbon is increased
by blowing hot air under pressure into the bottom of the furnaces.
Much of the carbon monoxide is not oxidized by the iron oxide to
carbon dioxide, and escapes at the top of the furnace through tubes
provided to receive it, and is used as fuel.

The operation of a blast furnace is continuous ; alternate charges
of coke, ore, and flux are being continually added at the top. of the
furnace, and the molten metal and the slag drawn off at the bottom.
The molten metal coming in contact with the hot carbon dissolves
a part of it, and iron thus prepared always contains some carbon,
as well as silicon and other substances, dissolved in it. The iron is
run into moulds made in the sand, and this impure product is known
as pig-iron^ or cast-iron.

Properties of Iron. — Pure iron is light gray in color, can readily
be drawn into wire, hammered or rolled into sheets. At a bright-
red heat it can be iveldedy or one piece made to adhere to another by
simply hammering the two together. Iron is a good conductor of
heat and electricity, and is one of the most resistant to strain of all
the metals. It is this property, together with its great abundance
and the ease with which it can be prepared, that makes it the most
valuable commercially of all the metals.

When iron is heated in the presence of the air, it readily bums,
uniting with oxygen and forming one of the oxides of iron. Iron
acts upon moist air even at ordinary temperatures, but acts slowly.
This is known as the rusting of iron. It does not act appreciably
upon dry air. Iron acts upon water at all temperatures, forming
the hydroxide: —

2 Fe -f- 6 HjO = 2 Fe(0H)5 4- 3 H^

While the action is slow at ordinary temperatures, it is rapid when
steam is passed over red-hot iron.

Iron dissolves readily in dilute acids, liberating hydrogen and

IRON 421

combining with the anion of the acid, forming the corresponding
salt. As we have seen, this is the same as to say that the hydrogen
ions give up their charges to the iron atoms, converting them into
ions, the hydrogen ions becoming atoms.

When iron is dipped into very strong nitric acid and then into
dilute, the latter is without action upon it. The iron is then in the
passive condition. This has recently been shown to be due to an
electrical condition of the metal, and not to the formation of a pro-
tecting layer of oxide over its surface as was formerly supposed.

Impure or Commercial Iron. — The different varieties of iron
which are used commercially have very different properties. These
are due to the different amounts of impurities in the iron. We have
already seen how pig-iron or cast-iron is made. Cast-iron is very
impure, containing in addition to from 3 to 4 per cent of carbon, 1
or more per cent of silicon, besides phosphorus, manganese, sulphur,
etc. If there is considerable silicon present, and the cast-iron cools
slowly, the carbon separates largely as- graphite, and gives a gray
cast to the iron. This is known as gray cast-iron. If the iron is
cooled rapidly the carbon remains largely in chemical combination with
the iron. Such iron is light in color and is known as chilled cast-iwn.

White cast-iron contains no graphite. It usually contains less
silicon or more manganese or sulphur than any gray cast-iron.

Cast-iron in general contains from 4 to 6 per cent of carbon, and
melts at a much lower temperature than pure iron. It is, therefore,
easily moulded, and gray cast-iron is used extensively for making
objects where great strength is not required. Cast-iron is brittle and
readily broken by a jar, and is far less tough than pure iron. Cast-
iron is not malleable, since it is too brittle, and although it melts
lower than pure iron, does not appreciably soften before it melts.
It therefore cannot be welded like pure iron.

If the ore from which the iron is made is rich in manganese, the
final product is also rich in manganese, and usually contains more
carbon than ordinary cast-iron. This is known as spiegel iron, and
contains from 10 to 15 per cent of manganese and in some cases
even more.

When most of the impurities have been removed from iron we
have wrought-iron. This still contains a small amount of carbon, the
amount, however, usually being less than one-half of one per cent.
Wrought-iron has very different properties from cast-iron. It is
very tough, strong, and malleable. It melts at about 2000®,' but
becomes soft at a bright-red heat, so that it can be hammered, rolled,
or welded. Wrought-iron, while extremely tough, is comparatively

422

PRKCIPLES OF ESOSGAXIC CHEMISTRY

soft, ud bends esiOj^ imdfiT fltrmtn. It is^ tli^^f (H€, not as usefitl as
a form of iioD which eontatiis mote carbon^ and is known as sleeL

Sied is nsnallj inm praetlcalty free from all impuriries except
carbon^ which is pi^setii to from ^M ta 2 per cent. There are two
general methods hj whieh ^eel may be made, — either bj remoring
earboQ and other im parities from ^^st-irotir or bj adding carbon to
WTOtighl-iroQ. The former process would seem to be the simplert
gmce it is necessary to remore the earboo from cast-iron in order to
obtain wronght-iron. The latter process^ howeverj is the one most
frequently made use ot A few methodB of making steel are so
importaat commercially and are so frequently referred to that they
will be briefly described.

The Beuemer ConYerter consists of a pear-shaped vessel of malle-
able iron, lined on the inside with ref ractorj' material. The molten c^^ t-
iron is i»oiired into the converter, and compressed air forced throtigh
the molten metaL The carbon and sili(M>n are completely oxidized by
the oxygen of the air, and the product is similar in eompositioa to
wrought^iron. This is kept above its melting-point by the heat of
combustion of the carbon and silicon. In order to obtain a product
with the desired amount of carbon^ Spiegel iron is added in quantity
anffieient to bring the percentage of carbon up to the desired amount.
The product is BeMsetner st^l, which has found such extensive appli-
cation in the arts.

The SlemeiiB^Martin Frocesi of making steel consists in beating
a mixture of wrought^iron which contains but little carljon with
pig-iron, iron ore being sometimes added. The gas used as fuel is
previously heated.

The Thomaa-Gllehnst ConTerter — The Bessemer process of mak-
ing steel does not remove the phosphorus from the iron. The
presence of an appreciable amount of phosphorus so changes the
properties of the steel as to render it entirely unfit for certain pur-
poses* While they were dependent solely \ipon the Bessemer or
similar processes, only certain iron ores which contain only a small
amount of phosphorus could be used for making steel. This has
largely been changed, due to the Thoin as- Gilchrist conv^erter. This
is essentially a Bessemer converter lined with burned dolomite, which
ifl a mixture of lime and magnesia. This *^ basic lining,** as it is
termed, unites with the phosphoric acid formed by the oxidation of
the phosphorus by the oxygen of the air wliioh is blown through the
molten iron, and forms calcium and magnesinin phosphates. This
material, known as the ** Thomas slag,'* is extensively used as a
source of jihosphoric acid in commercial fertilizers, having a com-

IRON 423

position very similar to the " phosphate rock/' which is so exten-
sively used as a fertilizer in the manner already described (p. 372).

Steel is useful largely because it can be made to assume any degree '
of hardness. When allowed to cool very slowly ordinary steel is soft
and resembles wrought-iron in its properties. When highly heated
and made to cool rapidly steel becomes very hard and brittle, the
degree of hardness depending somewhat upon the amount of carbon
present. The process by which the hardness of steel is regulated is
known as teinpering. The steel is heated moderately, the tempera-
ture being estimated by the color of the layer of oxide which forms
on the bright surface. This color is due to interference, like the
color of thin plates. The steel is then cooled more or less slowly.

Oxides of Iron. — Iron forms two well-known oxides, FcjOj, or
hematite, and Fe804, or magnetite.

Magnetite or magnetic iron ore is so called because it is strongly
magnetic. It is formed by burning iron in oxygen. As it occurs
in nature it is often beautifully crystalline, forming almost perfect
octahedra and cubes. It is sometimes regarded as the ferrous salt
of the hypothetical acid, HFeOj — FeFej04.

Ordinary ferric oxide, FcjOg, occurs in nature in great abundance
as hematite. It is also formed when a ferric salt or ferric hydroxide
is heated. In the form of a fine powder it is known as rouge, which,
on account of its color, is used as a pigment, and on account of its
fine state of division as a polish, where a very highly polished sur-
face is desired.

When ferric oxide is reduced by carbon monoxide, black ferrous
oxide, FeO, is formed.

Ferrons and Ferric Compounds. — Iron forms two kinds of ions,
— one carrying two electrical charges and known as the ferrous ion,
Fe, and the other carrying three electrical charges and known as the

ferric ion, Fe. These are what have hitherto been described as the
ferrous and the ferric condition. In the case of iron we can verify
the statement that Faraday's law is the base of chemical valence. If a
given electric current is passed through a solution of a ferrous, and
a solution of a ferric salt, one and one-half times as much iron will
separate from the ferrous solution as from the ferric. By comparing
the amount of iron which separates from a ferrous solution with the
amount of a univalent metal separated by the same current, it can

be shown that ferrous iron (Fe) is bivalent, or that the ferrous ion

+■♦•■♦• .
carries two charges of electricity. The ferric ion Fe is, therefore,

trivalent, or carries three charges of electricity.

424

PRINCIPLES OF INORGANIC CUEMISTRY

We sball aee that a ferrous ion can be converted into a ferric ion
by oxidatiuo, as it is said, and a ferric ion converted into a ferrous
ion by reduction. All that take^ place when a fen-ous ion is eon-
verted into a ferric ion is the addition of one electrical charge, and
the removal of on© electrical charge from a ferric ion converts it into
a ferrous ion. Oxidation and reduction as used in this sense are
simply the addition or removal of electrical energy, and, like valence,
have their physical basis in Faraday's law,

Ferrous (re(OH).) and Ferric (Fe(OH).i) Hydroxides. — The two
conditious described above are exemplified in the hydroxyl com-
pounds of iron. In one of these the iron holds two hydroxyl groups
in combination ; in the other, tliree. Ferrona hydroxide is precipi-
tated when an alkali is added to a solntion of a ferrous salt: —

FeClg + 2 :NaOH = 2 KaCl + Fe(OH)^

Ferrous hydroxide is white, but unites rapidly with the oxygen
of the air forming ferric hydroxide, which then reacts with ferrous
hydroxide, forming the black magnetite, or a hydroxide of this
substance : —

2 Fe(0H)3 + Fe(OH), = FeFe^O, + 4 H,0.

This black substance, in the hnely divided condition, mixed with
the white, ferrous hydroxide, gives it the well-known green ap-
pearance.

Ferrous liydroxide dissolves readily in acids, forming solutions
of ferrous salts. It does not dissolve in bases.

Ferric hydroxide is formed when an alkali is added to a solution
of a ferric salt : —

FeCl, -h 3 NaOH = 3 NaCl + Fe(OH)g.

It is a reddish-brown precipitate, readily soluble in acids, forming
solutions of ferric salts. Unlike aluminium hydroxide, it does not
dissolve in an excess of the base unless the latter is very conceutrated*

Ferric hydroxidej when freshly precipitated, dissolves readily in
a solution of ferric chloride. Since ferric hydroxide is a weak base
its salts are readily hydrolized*

If the hydrochloric acid is allowed to diffuse through a membrane
as it is set free from the salt, the decomposition of the salt will, in
time, become practically complete. By dfali/sis, it is then poi^sible
to decompose ferric chloride practically completely. The ferric
hydroxide, however, remains in solution, forming a dark-red liquid.
This liquid has the characteristic properties of a coUofdal nolution^
and is probably ferric hydroxide in a very fine state of division*

IRON 425

Ferrous (FeCl^) and Ferric (FeCI,) Chlorides. — Ferrous chloride
is obtained when iron is treated with hydrochloric acid. It forms
crystals containing four molecules of water, FeCl2.4 H,0. The salt
with water of crystallization is green in color. The white anhydrous
salt is formed by heating iron in a current of hydrochloric acid gas.
It forms double salts with the alkaline chlorides.

Ferric chloride is formed by passing chlorine over iron. It is
also formed by passing chlorine into a solution of ferrous chloride.
This reaction is of interest as illustrating a new method of ion forma-
tion. Ferrous chloride in solution is dissociated into a ferrous ion

++ — —
and chlorine ions, Fe, CI, CI.

Chlorine in solution in water is not dissociated. In the presence
of an iron cation with two electrical charges, which can take up a
third positive charge, the chlorine dissociates forming the correspond-
ing anion. The above reaction is then to be represented for sim-
plicity as follows, disregarding the fact that the chlorine molecule is
made up of two atoms : —

Fe, Ci, Ci -f- CI = Fe^ C^, C~l, CL

A cation takes up another charge converting an atom into an anion.
This is a method of ion formation not infrequently met with. Ferric
chloride crystallizes from aqueous solution with six molecules of
water, FeCls.GHjO. Its vapor-density shows that the molecule is
FeClj. It is readily transformed into ferrous chloride by reducing
agents.

Chemical Action at a Distance. — Ferrous chloride can be oxidized
to ferric chloride without the chlorine coming in contact with the
ferrous salt. The beaker containing a solution of ferrous chloride is
connected by means of a siphon with a beaker filled with a solution
of potassium chloride, into which chlorine gas has been conducted.
The siphon is filled with a solution of potassium chloride free from
chlorine, so that no free chlorine comes in contact with the feirous
chloride. A platinum electrode is immersed in each beaker and the
circuit closed. The current flows from the ferrous chloride to the
solution of chlorine in potassium chloride. The iron takes up
another charge of electricity from the electrode, passing into the
ferric condition, and we have chlorine atoms on the other side of the
cell passing over into ions. The result is the transformation of fer-
rous into ferric chloride, effected by chlorine which is not in contact
with the ferrous salt. As the current flows the iron moves with the
current over into the solution of chlorine, and the chlorine moves

426 PRINCIPLES OF INORGANIC CHEMISTRY

against the current over into the solution of the ferrous salt. Tl
however, is a secondary act, the oxidation being effected right aroi
the platinum pole immersed in the solution of the ferrous salt

Sulphides of IroiL — When iron filings and sulphur are hea
together the two combine and form ferrous gulphide, FeS. It is a
formed by the action of ammonium sulphide on a ferrous s
Heated in contact with the air it forms ferrous sulphate. Trea
with acids hydrogen sulphide is liberated.

Ferric sulphide or iron sesquisulphidey FcsSg, is formed bj heat
ferrous sulphide with sulphur. It is also formed when bydro^
sulphide is passed over iron heated to about 100**. Iron disulphi
FeSj, is the familiar substance pyrites^ which occurs so widely c
tributed in nature and iu great abundance. On account of its co
it is frequently taken for gold, and hence has acquired tbe name
fodCs gold. It can be prepared by passing hydrogen sulphide o^
iron oxide heated to a temperature somewhat above 100®.

Ferrons Sulphate, FeSOf.TE/). — Ferrous sulphate, also eai;
" iron vitriol " on account of its composition, or " green vitriol '^
account of its color, is the most important ferrous compound. It
formed by the action of sulphuric acid on iron or iron sulphide. It
made commercially by allowing ferrous sulphide to take up oxyg
from the air. If pyrites is used, one-half of the sulphur is roast
out, and the ferrous sulphide is then moistened and allowed to ta
up oxygen from the air : —

FeS +20, = FeS04.

The salt is then extracted with water. Ferrous sulphate fori
light-green crystals, and is extensively used in dyeing, in pharmac
and as a disinfectant.

Ferrous sulphate, like other sulphates already studied, gives i
six molecules of water at a comparatively low temperature. The la
molecule is not set free until a temperature of about 300® is reache
If ferrous sulphate is allowed to crystallize from a solution at 8(
the salt which conies down contains only four molecules of water -
reS04.4 HjO. Ferrous sulphate readily forms double salts witb tl
alkaline sulphates.

Ferric Sulphate, Fej(804)3, is formed by dissolving ferric oxide <
hydroxide in sulphuric acid : —

2 Fe(0H)3 -h 3 HSO, = Fe,(S04)3 + 6 HA

It is also formed by the addition of a half-equivalent of sulphur
acid to ferrous sulphate, in the presence of an oxidizing agent lil
nitric acid. Ferric sulphate forms double sulphates with the alka

'I

IRON 427

sulphates, which, in composition and crystalline form resemble the
aluminium alums, and are termed iron alums.

Potassium Ferrooyanide, S^Fe(CH')e. — Although iron does not
combine directly with the cyanogen ion and form cyanides, it
forms double cyanides which are beautifully crystallized compounds.
When potassium cyanide is allowed to act upon iron in the presence
of water the following reaction takes place : —

Fe + 6KCN -f 2H,0 = 2K0H -f H, + K^FeCCN),.

Potassium ferrocyanide is usually formed by heating nitrogenous
matter with iron filings and potash. When the mass is digested
with water and the solution evaporated, beautiful yellow crystals are
formed, having the composition K4Fe(CN)e.3HjO. This is potas-
sium ferrocyanide, known commercially as yellow prussicUe of potash.

When this compound is heated it is decomposed into potassium
cyanide, iron carbide FeCj, and nitrogen. When it is treated with a
strong acid a white solid is thrown down : —

K4Fe(CN)e -h 4 HCl = 4 KCl -f H^FeCCN)..

This substance, H4Fe(CN)«, hydroferrocyanic acid, is the acid of
which potassium ferrocyanide is the salt. It dissolves in water,
forming a strongly acid solution. This fact, together with the com-
position of the potassium salt, shows that it is dissociated by water
in the following manner : —

H^FeCCN), = H, H, H, H, Fe(CNV

The anion of this acid is interesting on account of its composi-
tion. In addition to the six cyanogen groups, each of which carries
a negative charge, it contains ferrous iron with two positive charges.
The result is an anion with four negative charges. It is interesting
to note also that a positively charged constituent (iron) may form a
part of an auion. The iron in this complex anion has lost its char-
acteristic properties, as we are accustomed to say ; ue. it no longer
possesses the properties of iron when alone in the ferrous or ferric
condition, nor is there any reason to expect that it should.

Hydroferrocyanic acid forms well-characterized salts. The most
important of these is the ferric salt : —

3 H4Fe(CN)e -h 4 FeClj = 12 HCl -h Fe4(Fe(CN)eV

This is the well-known substance Prussian blue or Berlin Wue,
which is valuable as a pigment, and is formed whenever ferric ions

come in contact with the anion of hydroferrocyanic acid, Fe(CN)t.

428 PRINCIPLES OF INORGAKIC CHEMISTRT

This reaction is, tliereforej a very sensitivB test for tlm presence of
fm^ric iom.

The copper salt of this acid has acquired physical chemical dis-
tinction in connection with the demonstmtion and measurement of
osmotic pressure. This substance is formed by the action of any
soluble eupric salt upon potassium ferrocyanide ia sointion : —

K4Fe(CN)a + 2 CuSO, = Cu^Fe(CN)a 4- 2 K^SOv

Copper ferroci/ankk is a reddish-brown gelatinous solid, resembling
in appearance ferric hydroxide. When deposited in the walls uf
porous cups it allows water to pass through but prevents the dissolved
substance from passing. It is, therefore, used in the construction of
sefni'permeabh menubmnes. Prussian blue has been used in the same
connection but far less successfully. Calcium phosphate has also
been used, but copper ferrocyanide gives by far the best results, as
we have already seen.

Potassium Ferricyanide, XjFe(CH)4, is formed by the action of
oxidlaiiig agents on potas^sium ferrocyanide. If chlorine is passed
into a solution of potassium ferroeyanidCj potassium ferricyaiiide is
formed : — ^ ^^^ ^^^^^ + CI, = 2 KCl + 2 K,Fe (CN),.

The compound KaFe(CK)^ is known also as red prussime of

p^ctshf from the deep-red color of its crystals. When this c^jm-
pound is treated with an acid hydroferrkyanic add is liberated ; —

KsFe (CN)ft + 3 HCl = 3 KCl + H^Fe (CN)^.

This dissolves in water, dissociating as follows : —

HaFe(CN)«^ 6, H, H, Fe(Cl!^,.

The anion of this acid is the same in composition as the anion
of hydroferrocyanic acid. The differeuee is that here the anion
carries three electrical charges, while in the ferrocyanides it carries
four. In the ferri compounds there are^ therefore, three univalent
cations, and in the ferro compounds four*

When a ferricyanide is treated with a ferrous salt, the ferrona
compound of hydroferri cyanic acid is formed; —

2 K,Fe(CN)a + 3 FeCl, = 6 KCl + Fe,(Fe (CK).)*,

This substance is known as TumbuWs blue, and this reaction ia
one of the mo^t sensitive tests for ferrous ions,

Prussian blue and Turnbiiirs blue are interesting as showing the
same metal in the same compound in two different ionic conditions.

I

IRON 429

In these compounds we have iron in the cationic condition, and also
united with cyanogen as a part of an anion. The iron in the former
condition shows the reactions which we are accustomed to ascribe to
iron ; the iron in combination with cyanogen in the anions does not
show these reactions. This illustrates another fact, that the reao
tions which we are accustomed to ascribe to an element are reactions
of the ions of that element, and of the ions alone.

Change in Color with* Change in Electrical Charge. — An ion
having the same chemical composition does not always have the
same color. Take the ion Fe(CN)e; in potassium ferrocyanide it
is yellow and gives the yellow color to a solution of this salt. The
ion in this case is formed by the dissociation of the salt K4Fe(CN)«

-♦-+ + + s

into K, K, K, K, and Fe(CN)j, which carries four negative charges.
The ion Fe(CN)e, obtained by the dissociation of potassium ferricya-
nide, is red. The compound K3Fe(CN)e dissociates as follows : —

KaFe (CN)e = K, K, K, Fe (CN)^

The ion Fe (CN)e, in this case, carries three negative charges, and
the difference of one charge changes the color of the ion from yellow
to red.

To take a simpler example: The iron ion in the ferrous condi-

'*■"'"..
tion, Fe, is green, as is seen in solutions of ferrous salts ; while the

+++
iron ion m the ferric condition, Fe, is yellow, as is seen in solutions

of ferric salts. A large number of examples of changes in the color

of ions with change in the electrical charge which they carry, might

be given.

Other Salts of Iron. — The ferric salt of sulpho-cyanic acid is of
importance in connection with the detection of iron. Ferric sulpha-
cyanaJte, Fe(CNS)8, is deep blood-red in color when undissociated.
A solution containing molecules of this substance has, therefore,
a characteristic red color. Such a solution is prepared by adding
to a ferric salt an excess of potassium sulpho-cyanate, which is the
same as adding an excess of sulpho-cyanogen ions. In accordance
with the general principle with which we are familiar, this would
drive back the dissociation of the ferric salt and bring out the color
of its molecules.

Sodium nitropruasiate, NajFe (CN)5N0 . 2 HjjO, is formed by treat-
ing sodium ferrocyanide with nitric acid. It is useful in testing for
the alkaline sulphides, with which it gives a purple color.

Iron like nickel combines with carbon monoxide forming car-

430 PRINCIPLES OF INORGANIC CHEMISTRY

bonyl compounds. Several of these are known having the composi-
tions Fe(CO)«, Fe(C0)4, etc.

Ferric Acetate, Fe (CHaCOO),, is deep red in solution, and is a
fairly sensitive test for iron. Like the salts of weak acids in general
it is hydrolized by water, and when its solution is boiled the basic
acetate is precipitated. This ifi of importance in connection with
the quantitative determination of iron.

Ferrates. — We have seen that aluminium can act as an acid-
forming element, the hydroxide being soluble in sodium hydroxide.
Iron can act in the same capacity, but as an acid-forming element
has a valence of six. Ferric acid has the composition H,Fe04, and
is, therefore, analogous to sulphuric acid, and^ as we shall see, to
chromic acid.

The potassium salt of this acid is formed by the action of strong
oxidizing agents on iron, in the presence of potassium hydroxide.
When iron oxide is treated yrith chlorine in the presence of potas-
sium hydroxide, potassium ferrate, KjFeO^, is formed. This same
compound is also formed when iron is heated with potassium nitrate.
Other salts of ferric acid are known.

CHAPTER XXXIV

COBALT AND NICKBL
COBALT (At Wt = 69.0)

The chief sources of cobalt are the mineral smaUite, which is the
arsenide of the composition CoAs^, and the mineral cobaltUe, which
has the composition CoAsS. The element is prepared by reducing
tne oxide either with highly heated carbon, or with hydrogen. Co-
balt resembles iron in many respects. It has a somewhat lighter
color, with a distinctly reddish tint. It melts at about the same
temperature as iron. It forms a coating of oxide in contact with
moist air, but not with dry air. Like iron it decomposes water
readily at a high temperature. It dissolves slowly in hydrochloric
and sulphuric acids, and readily in nitric acid.

Cobaltoiu and Cobaltic Compounds. — Cobalt, like iron, forms two

kinds of ions. The cobaltous ion Co, and the cobaltic ion Co.
Unlike iron, however, the cobaltous ion is the more stable condi-
tion, while the ferric condition is the more common for the iron ion.
The cobaltous ion combines readily with the anions of acids, form-
ing solutions of cobalt salts. The cobaltic ion also can form salts
with certain anions, but the cobaltic condition is especially met with
in complex compounds.

Oxides and Hydroxides of Cobalt. — Several oxides of cobalt are
known. Cobaltous oxide, CoO, is formed when cobaltous carbonate or
hydroxide is heated. It is a greenish powder, easily reducible to
the metal. Cobaltic oxide, or cobalt sesquioxide, C02O3, is formed by
gently heating the nitrate. It is a dark-brown powder, which passes
over, when heated, into cobalto-cobcdtic oxide.

When a cobaltous salt is treated with an alkali, cobaltous hydrox-
ide is formed : —

CoCl, -h 2 KOH i= 2 KCl -h Co(OH)^

At first a basic salt which is blue is formed, but this gradually
decomposes into the red hydroxide. Oxidizing agents readily con-
vert tMs into cobaltic hydroxide, Co(OH)a, which is black. Cobalt also

431

b

482 PUmCIPLES OF INORGANIC CIlEMISTUr

forms an acid which corresponds to the hydroxide Co(OH)# minus
water — HtCoOjp Certain salts of this acid are known and are called

cobalt tf*s,

Cobaltons Salts. ^ — When metallic cobalt is heated in chlorine,
cobaU diforidtt COClu is formed- This crystallizes from an aque-
ous solution with six molecules of water — CoCl^.BHjO- The solu-
tion of tlie salt is recL When tlie water is removed the salt is deep
blue in color, due to the dn%^ing back of the ions into molecides,
which are blue. Cobalt chloride has, therefoiT, heen used as sifm^tor-
thetk htk. When a solution of the salt is used for writing on paper,
the writing is practically colorleas^ due to the nearly colorless nature
of an aqueous solution of eobaltous chloride. When the pajjer is
wanned the blue color appears, and the writing l)ecomes plainly
legible. When the blue material is allowed to stand in contact with
the air, it takes up moisture, l>ecoming again invisible.

Cobalt nitrate^ Co(N08)^ is one of the most common of the
cobalt salts. It crystallizes with six molecules of water, forming
beautifully red prisms.

CobaU auljyhaJe, CoHOiJ 11*0^ '^ a beautifully crystallized com-
pound, and like so many other sulphates contains seven molecules of
water of crystallimtltin. It is isomorphous with ferrous sulphate.

CobaU mdphidef CoS, is formed when ammonium sulphide is
added to a solution of a col*altons salt. When once formed cobalt
sulphide does not dissolve in dilute acids, and this fact is made use
of in separating it from other sulphides of the same group. Cobalt
sulphide, however, is not precipitated from a solution of a neutral
colialfc salt, unless some method is adopted to remove the free hydro-
gen ions which would be formed as the result of the reaction. This
is effected by adding to the solution sodium acetate^ when acetic
acid is formed, and this is practically nndissociated.

Cobalt forms blue ^illimfeji^ Wlien a cobalt salt is added to color-
leas glass and the mass fiiscdj the resulting glass is deep blue iu
color. Cobalt glass is finely powdered and used as a pigment under
the name of smalL Cobalt glass cuts oit the yellow rays of lights and
it will be remembered that it is used for this purpose in qualitative
analysis to detect the presence of jwtassium when sodium is present.

Cobalt also colors the microcosmic head, or the borax bead^ deep
blue when heated in the flame of the blowpipe. This reaction is
made use of to detect cobalt in blowpipe analysis.

BoublB Cyanides of Cobalt. — ^ Cobalt foruis two donljle cyanides,
which arc strictly analogous to the two compounds of iron. l^Tiert*
eobaltous cyanide is dissolved in potassium cyanide the two

COBALT AND NICKEL 433

combine and form potassium cobaltous cyanide, K4Co(CN)e. This

is the analogue of potassium ferrocyanide, and dissociates into
+ + + + =

K, K, K, K, Co(CN)^ The anion which contains the cobalt and six
cyanogen groups is quadrivalent, like the ferrocyanogen ion.

When a solution of this compound is boiled in the presence of
the oxygen of the air, the compound is oxidized to potassium cobalti-
cyanide, K3Co(CN)8, which is analogous to potassium ferricyanide.
In the presence of water this dissociates into K, K, K, Co(CN)e, the
cobalticyanogen ion being trivalent. The acid HsCo(CN)e is well-
known. From neither the cobaltocyanogen nor the cobalticyanogen
ion do the ordinary precipitants of cobalt throw down the cobalt The
cobalt in these ions, like the iron in the corresponding ferro- and fer-
ricyanogen ions, does not have the ordinary properties of cobalt.

Double Nitrite of Cobalt. — Cobalt forms a double nitrite with
potassium, having the composition K8Co(N02)e. This is obviously
the salt of the acid H3Co(N02)6, which, however, has never been
isolated. It is formed by adding potassium nitrite to a solution
of a cobalt salt. The difficultly soluble potassium salt is thus
precipitated as a yellow powder. The corresponding sodium salt
Na3Co(N02)e is readily soluble in water.

Action of Ammonia on Solutions of Cobalt Salts. — When solutions
of cobalt salts are treated with ammonia and exposed to the action
of the air, a number of complex compounds are formed. These have
been studied extensively, and the composition of a number of them
established. Thus, compounds of the composition [Co(NH3)4A'j]A
(where A is an anion, CI, KDa, etc., and A' an acid-forming atom or
group), etc., are known as praseo compounds. Another series of com-
pounds are known having the composition [Co(NH8)5A']Aj, and are
termed purpureo compounds, while still another series exists, having
the composition [Co(NH3)e] A^, and are known as the luteo compounds.

NICKEL (At. Wt. = 58.7)

An element which resembles cobalt in many respects is nickel.
It occurs in nature chiefly in combination with arsenic, NiAs, as
nkcolite, and with arsenic and sulphur, NiAsS, as gersdorjffUe. The
silicate occurs in abundance and is known as gamierite.

Nickel is prepared by reducing the oxide with carbon at a high
temperature, and by reducing the oxide in a stream of hydrogen.

Nickel is light in color, with a yellowish cast. Although hard, it
is malleable. It melts at about the same temperature as iron. It is
2f

434

PRmCIPLES OF INORGANIC CHEMISTRY

oxidized in the air with difficulty, but is dissolved by hydrochloric
and sulphuric acids, and especially by nitric acid*

On account of its resiatance to oxidation, nickel is extensively
used to cover metals which are more readily oxidized^ such as iixjn,
etc. The nickel is deposited upon the iron electrolytically. The
iron object is made the cathode of a suitable electric current, and
this is immersed in a solution of a nickel salt, the double suli^hate
with ammonium being frequently used. The anode is of nickel, and
supplies as much nickel to the bath as is deposited on the cathode.
The nickel ions give up their charges to the cathode, and are de-
posited in the form of metal upon the cathode. This method of
covering one metal with another is known as eleciro-phUin^,

Nickel forms valuable alloys with a number of the metals. Chr-
man mher is an alloy of nickel with zinc and copper. Alloys of
nickel and copper are used as coins; our so-ealled "nickel" contain-
ing 75 per cent copper and 25 per cent nickeL

Compoands of IViokeL — While there are a few compounds known
in which nickel plays the role of a trivalent element, it is almost

always present as the bivalent ion NL The omde of nickel^ KiO, is
formed as a black powder when the hydroxide is heated in a limited
supply of air. When there i3 an abundant supply of oxygen the
aesqnioxidey Ni^Oj, is formed.

The green hifdroxidej Ni(OH)aj is formed when a solution of a
nickel salt is treated with a solution of a hydroxide : —

mc\, + 2 KOH = 2 KCl -h Ni(OH)^

When the nickelous hydroxide is oxidized with chlorine, m^eOc
hydroxide, Ni(OH)a, is formed. The chloride of nickel is formed by
heating the metal in a current of chlorine* It crystallizes with nine
molecules of water of crystallization — NiClj,9 H^O.

When the oxide of nickel or the metal is dissolved in dilute sul-
phuric acid, the beautifully green sulphate crystallizes from the
solution. This has the composition liTiSOi.THuO, and readily forms
double sulphates with the sulphates of the alkalies.

A remarkable compound of nickel is the one formed with carbon
monoxide. When nickel and carbon monoxide remain in contact at
30°, they combiue and form a liquid compound which has the com-
position Ki(C0)4 ^nd is known as nickel carhomjl, or nkkd tetmcar-
bontfL Nickel earbonyl boils at 43° and solidifies at — 25°. At &
temperature somewhat above its boiling-point, nickel tetracarbonyl
decomposes into nickel ami carbon monoxide. This method has
been used for purifying nickel, but has the disadvantage that the

I

I
I

COBALT AND NICKEL 435

carbon monoxide set free decomposes in part into carbon dioxide
and carbon.

The cyanide of nickel is formed when potassium cyanide is added
to a solution of a nickel salt: —

2 KCN + NiS04 = KjjS04 + Ni(CN)^

The greenish cyanide which is precipitated readily dissolves in
an excess of potassium cyanide, forming the double cyanide K^i(CN)4.
^Vhen potassium nickelous cyanide is treated with an acid, in all
probability the acid H^i(CN)4 is formed, but this is decomposed
at once into hydrocyanic acid and nickel cyanide.

With potassium nitrite nickel forms the double nitrite^ K^Ni
(NO,V

CHAPTER XXXV

MANGANESE (At. Wt. = 55.0)

We now come to an element which probably forms as large a
variety of compounds as any element known. This is due to the
many degrees of valence which manganese can manifest. Ou the
whole, the element shows many analogies to iron, and undoubtedly
belongs in the iron group. While there are certain analogies between
manganese and chlorine, they are not very striking. Indeed, far less
striking than the differences, and it must be regarded as a weakness
in the Periodic System that manganese falls in the same group with
the halogens.

Occurrence, Preparation, and Properties of Manganese. — Man-
ganese occurs in nature in small quantities in the free condition,
but usually as one of its oxides. The chief soui-ce of manganese
is the oxide MnOj, which is the mineral pyrolusite. Other man-
ganese minerals are hausmannite, Mn304, hraunite^ Mn,Oa, and rhodo-
croisite, MnCOj.

Manganese is prepared by heating the oxides with carbon in an
electric furnace, also by electrolysis of the fused chloride, but more
conveniently by mixing the oxide with finely divided aluminium and
igniting the mixture. This is one of Goldschmidt's mixtures, the
aluminium taking the oxygen, setting free the manganese.

Manganese has but little commercial value, since it is so readily
attacked by chemical reagents. It is oxidized in the air, decomposes
water even at ordinary temperatures, and is readily attacked by acids.

Some of the alloys of manganese are of value. The alloy with
iron known as spiegel iron, has already been referred to. The alloy
with copper containing some zinc is known as manganese bronze, and
is quite valuable.

Oxides of Manganese. — Manganese forms no less than seven
compounds with oxygen. The one containing the smallest amount
of oxygen is manganous oxide, MnO. This is formed by reducing
the higher oxides in a stream of hydrogen, and by heating manga-
nous hydroxide. Manganese sesquioxide, IMn^Og, occurs in nature as
braunite, Manganous-manganic oxide, ^In304, is formed when the

430

MANGANESE 437

other oxides of manganese are heated in the air. Manganese dioxide,
MnOj, occurs in nature as pyrolusite, and is the most important ore
of manganese. There exists a trioxide of manganese, MnOs, and
also a septoxide, MujOy. The latter is formed by treating potassium
permanganate with sulphuric acid : —

2 KMn04 -h H^^O* = K,S04 + H,0 -h IVIn A-
There also exists a tetroxide of manganese — Mn04. Arranging
these oxides in the order of increasing amount of oxygen, we have: —

MnO,
MnO MnA Mn^Oy

Mn304 MnOj Mn04.

Hydroxides of Manganese. — Manganous hydroxide is precipitated
as a white powder when an alkali is added to a manganous salt : —

MnCl, -h 2 NaOH = 2 NaCl + Mn(OH)^

Manganous hydroxide readily takes up oxygen in the presence
of alkalies and passes over into the dark-brown manganic hydroxide,
Mn(0H)3. Manganic hydroxide minus water, HMnOj, occurs in
nature as manganite.

The hydroxide Mn(0H)4 can be prepared by treating a manganous
salt with an alkali in the presence of oxidizing agents. This hy-
droxide minus water is manganous acid, HjMnOs, which forms a
series of salts known as the manganites.

The partial anhydride of the supposed hydroxide Mn(OH)e,
which can be regarded as formed from that substance by loss of two
molecules of wat^r — H2Mn04 — is manganic a^cid. This acid is
unstable and does not exist in the free condition. Salts of this
acid, or the manganates, are well known.

One other hydroxyl compound of manganese merits special
consideration. This is permanganic acid. It has the composition
HMn04, and may be regarded as the partial anhydride of the
hydroxide Mn(0H)7: —

Mn(0H)7 = 3 H,0 -f HMn04.

Permanganic acid is quite stable in aqueous solution, and can be
prepared by dissolving manganese septoxide in water, but far more
conveniently by electrolyzing the potassium salt, as we shall see.
By this method permanganic acid can readily be prepared in any
quantity desired.

Manganons Salts. — The manganous ion, Mn, combines with the
anions of acids, forming salts which are usually beautifully crystal-

438 PRINCIPLES OF INORGANIC CHEMISTRY

lized compounds. Manganous chloride, MnCls.iH/), is formed bj
the action of hydrochloric acid on manganese dioxide : —

MnO, + 4 HCl = 2 H,0 -h CI, -f MnCV

WHien manganous chloride is treated with lime-water, manganous
hydroxide is formed : —

MnCl, -h Ca(OH), = CaClj + Mn(OH),.

When manganous hydroxide is treated with lime and allowed
to stand exposed to the air, it undergoes oxidation and forms the
calcium salt of the acid H^MnOj, which is Mn(0H)4— H,0: —

2 Mn(OH), -h 0, -h 2 CaO = 2 H,0 -h 2 CaMnO,.

In the Weldon process for making chlorine, the above transforma-
tions are effected in order that the manganese chloride which is
formed may not be lost. When calcium manganite is treated with
hydrochloric acid, chlorine is set free, just as when the original
manganese dioxide was treated with hydrochloric acid : —

CaMnOs + 6 HCl = MnCl, + CaCl, + 3 HjO + CV

Manganous sulphide, MnS, occurs in nature as manganese blende.
It is prepared by passing the vapors of carbon disulphide over heated
manganite. It is dark in color as it occurs in nature. It is soluble
in dilute acids, and cannot, therefore, be precipitated from solutions
of manganous salts by hydrogen sulphide. When ammonium sul-
phide is added to a solution of a manganous salt a pinkish precipi-
tate is formed, which is a hydrate of manganous sulphide. When
this is allowed to stand it loses water and forms green manganous
sulphide. When ammonium sulphide is added to a hot, concen-
trated solution of a manganous salt, the anhydrous, green sulphide is
thrown down at once.

Manganous sulphate, MnS04, is formed by dissolving the oxides
of manganese in hot, concentrated, sulphuric acid. It does not
matter which oxide is used, the product is the manganous salt.
This means that under these conditions the higher oxides must lose
oxygen when boiled with concentrated sulphuric acid, and such is
the fact. The salt exists in a number of crystal forms, depending
upon the conditions of its formation. AVhen the temperature of the
solution from which the salt crystallizes is 0® or below, a salt with
seven molecules of water separates, MnS04.7H20. This modifica-
tion is analogous to ferrous sulphate. The modification with five
molecules of water, MnS04.5H20, crystallizes from a solution be-
tween 15° and 20^ This is analogous to copper sulphate. At a still

MANGANESE 439

higher temperature, 20® to 30", the salt MnS04.4H20 is formed;

while at much higher temperatures, extending even above 200% the

salt MnS04.H,0 is stable.

Manganous carbonate ^ MnCOjj, occurs in nature under the name

of manganese spaj-, and is formed when a manganous salt is treated

with a soluble carbonate. When heated in contact with the air, it

forms the compound MugOi, like most manganese compounds.

+++
Manganic Compounds. — The manganic ion, Mn, can unite with

the anions of acids and form salts. The manganic salts are, how-
ever, not as numerous as the manganous, and in general not as stable,
being strongly hydrolyzed by water. A few of these will, however,
be considered. The oxide, AfnjOa, the hydroxide, Mu(0H)3, and the
partial anhydride, HMnO^ have already been referred to. When
manganic hydroxide is treated with hydrochloric acid, there is reason
to believe that manganic chloride, MnClg, is formed. This cannot be
isolated, since it decomposes spontaneously on standing into chlorine
and manganous chloride.

Manganic sulphate, Mn2(S04)8, is formed when manganese dioxide,
MnOj, or manganous-manganic oxide, is dissolved in slightly warmed
sulphuric acid. If the acid is hot, oxygen escapes and manganous
sulphate is formed, as we have seen. The dark-green manganic sul-
phate is comparatively unstable, easily losing oxygen and forming
manganous sulphate. It is, therefore, an excellent oxidizing agent.
It combines with the alkaline sulphates, forming manganese alums
of the general composition, MMn(S04)2.12H20.

Tetravalent Mang^ese. — A few compounds are known in which
tetravalent manganese apparently exists. This is the case with the
oxide, MnOj, the hydroxides, Mn(0H)4, H,MnOj„ the supposed chlo-
ride, MnCl4, and the sulphide, MnS,. The most important of these
substances is manganese dioxide, which, as we have seen, formerly
found extensive application in the preparation of chlorine, and to-day
is largely used in the arts as an oxidizing agent

It is also used in the construction of one of the most efficient

forms of primary cells, the Leclanch4 cell. The action of this cell,

which consists of carbon and manganese dioxide as one pole, and

zinc as the other pole, ammonium chloride being the electrolyte,

depends largely upon the transformation of tetravalent manganese,
++++
Mn, into manganese of lower valence.

Valence and Properties of Manganese. — It should be noted that
as the valence of manganese increases, its basic nature rapidly dimin-
ishes. Bivalent manganese is distinctly basic, forming stable salts

440

PRIXXITLES OF IXORGAXTC CHEMISTRT

I

with the anions of a^iida. Trivalent manganese is ^ery weakly hasic^

its salts being strongly hydrdyzecl by water* Tetravalent manga'
nese is scarcely bfisic at all, its cgini>onn<l with such a strong acid as
hydrochJoric being so unstable that its very existence is doubtfuL

When we pass to manganese with higher valence, not ouly has
all the basic nature been lost, but we find acid properties beginning
to manifest themselves, aud the highest oxidation product of man-
ganese is a strong acid* These acid compounds of manganese wb*h
shall novv study. (

Manganoufl Acid, H^JInO,,, can be regarded as formed from the
hydroxide, Mn(0H)4, by loss of one molecule of water* Salts of the
above compound are known. Calcium manganite, CaMuO^, is formed,
as we have seen, by the action of oxygen on a mixture of nianganous
hydroKide and lime. It is the S4>called Wtklon mndf obtained in ^
the preparation of chlorine by the Weldon proeess, using manganese |
dioxide and hydrocliloric aciiL This compound is not very stable,
easily losing oxygen and passing into the nianganous condition.

Manganic Acid, H^JInO^. — We have now studied compounds of
manganese in which this f^lement has appeared in the capacity of a
bivalent, trivalent, and quadrivalent ion. Fentavalent manganese is
not kiiown^ but hexavalent manganese is well known^ manifesting
itself in salts of the compound manganic acid, the analogue of sal-
pluiric acid. These are formed, as we would expect, by stronglj
o X i <l i zi n g m an gau ese i n t h e prese nee of bases . Potassi w m ma uga » o^c,
KaMuO^j is formed by fusing potassium hydroxide with nianganese
dioxide in the presence of the oxygen of the air, or l)etter, with an
ox:idizing agent such as potassium clUorate, The manganese is osi*
dized from the tetravalent to the heKavalent condition, the potassium
chlorate being reduced to potassium chloride: —

3 MnO. + KCIO3 -h 6 KOH = KCl + 3 H,0 4^ 3 K,MjiO<.

This mass forms a green solution, from which green crystals of
potassium manganate, KfMnO^, separate. This compound is stable
only in alkaline solutions. When brought into the presence of the
air or an acid it decomposes, owing to the instability of manganic
acid itself Indeed, the acid is so unstable that it has never been
isolated* Wh^n i>otassium manganate is treated with an acid the
following reaction takes place : —

3 KjMnO* -h 6 HCl = 6 KCl + 2 HMnO, + MnOt + 2 H,0.

Instead of obtaining manganic acid tins breaks down into manganese
dioxide and pennanganic acid, This same transformation is effected

MANGANESE

441

by carbon dioxide and, consequently, takes place slowly when a man-
ganate is exposed to the air : —

3 iykln04 + 2 CO, = MnO, + 2 KjCOa + 2 KMn04:

The change of color from the green manganate, through blue and
purple to the purplish-red permanganate, is very striking. This was
early observed and termed mineral chameleon.

Permanganic Acid, HMnOf. — The highest oxidation product of
manganese containing hydrogen and oxygen is permanganic acid,
HMnOi, the analogue of perchloric acid and persulphuric acid.
One method of preparing the acid which we have just studied con-
sists in the action of acids on potassium manganate. Another
method which has already been referred to consists in the elec-
trolysis of the potassium salt of this acid. This method, which was
devised by Morse, is carried out as follows: Two unglazed porcelain
cups containing the one water to which a little alkali is added, and
the other water to which a little permanganic acid is added if
available to make the water conducting, are immersed in a beaker
containing a solution of potassium permanganate. The platinum
electrodes are inserted the one in each cup, the cathode in the cup
containing the alkali. The current is passed, when the potassium
ions move toward the cathode, give up their charge, decompose water,
and liberate hydrogen. The alkali formed around the cathode is
easily siphoned off from time to time. The permanganic ions, Mn04,
move to the anode, decompose water forming permanganic acid, and
liberate oxygen. The permanganic acid collects in the cup around
the anode, and after the current has been passed for a sufficient time
can be obtained in perfectly pure condition and in any quantity de-
sired. This method of preparing permanganic acid so far surpasses
all others that they are only of historical interest. Permanganic
acid is a very strong acid as is shown by its large conductivities.

w

f^y

a

16

352.3

93.4 0/,

128

376.0

90.3

612

376.6

00.8

1024

377.3

100.0

Potassium permanganate^ KMn04, is readily obtained by passing
carbon dioxide through a solution of potassium manganate, as

442 PRINCIPLES OF INORGANIC CHEMISTRY

already described. Its solution has exactly the same color as per-
manganic acid, — purplish-red. It crystallizes in beautiful purplish-
red crystals, which are not very soluble in water, one part of salt
requiring about sixteen parts of water to dissolve it at ordinary
temperatures.

Potassium permanganate is characterized chiefly by its oxidizing
power. When its aqueous solution is treated with an alkali in the
presence of a reducing agent, it breaks down as follows : —

2 KMn04 + alkali -h 5 H,0 = 2 KOH + 2 Mn(0H)4 + alkali, + 3 O,

two molecules of the salt giving three oxygen atoms.

In the presence of an acid, however, the reduction goes much
farther, the manganese being reduced to the manganous condition: —

2 KMn04 + 3 H,S04 = K:^04 + 2 MnS04 + 3 HjO + 5 0,

two moleades of the permanganate yielding Jive atoms of oxygen.

The oxidizing action of the permanganates is shown especially
by those permanganates which are very soluble in water. Calcium
and strontium permanganates are extremely soluble in water, one
part of water dissolving 2.9 parts of strontium permanganate and
3.3 parts of calcium permanganate. Concentrated solutions of these
salts oxidize organic compounds with the greatest energy. When
a drop of the solution -of the permanganate is allowed to fall into
oil of turpentine or glycerine, the oxidation takes place almost with
explosive violence.

On account of its oxidizing action potassium permanganate is
used extensively in analytical chemistry. It always yields a definite
amount of oxygen in alkaline solution and a definite amount in acid
solution, and its oxidizing power is therefore known. It is only-
necessary to know the strength of the solution of the permanganate
and the amount used, in order to know the amount of oxidation
which will be effected.

As an example of the uses of potassium permanganate in analyti-
cal chemistry, take its action on oxalic acid. This substance is
oxidized to carbon dioxide and water by the permanganate in acid
solution. Knowing the strength of the permanganate solution,
and the amount employed to just oxidize all of the oxalic acid,
we can calculate at once the strength of the solution of oxalic
acid. The end of this reaction is determined by the appearance

MANGANESE

443

of the color of the permanganate as soon as all of the oxalic acid
is used up.

Color of Permanganates. — The color of the solutions of the per-
manganates is of special interest in connection with the theory of
electrolytic dissociation. The permanganates are compounds of the
metal cations with the anion, MnOf. If we select cations which
are colorless, the color of these permanganates is due entirely to the
permanganic ion, Mn04. They should all, therefore, have exactly
the same color.

This interesting conclusion from the theory of electrolytic disso-
ciation has been tested experimentally by Ostwald. He prepared
solutions of a number of salts of i)ermanganic acid with such color-
less cations as potassium, sodium, ammonium, lithium, barium, mag-
nesium, aluminium, zinc, cadmium, etc., and then studied their
absorption spectra, or the wave-lengths of lights which would be
cut off when white light was passed through their solutions. The
absorption bands were both measured and photographed by Ostwald.
These salts show five absorption bands in the yellow and green, and
four of these were measured by Ostwald for thirteen salts of per-
manganic acid.

The results of Ostwald's measurements are given in the following
table: —

Pebmanoanates. Absorption Bands

I

II

. "I

IV

Hydrogen ....

2601 ± 0.6

2698 ± 0.8

2804 ± 0.7

2913 ± 1.7

Potassium

2600 ±1.3

2697 ± 0.1

2803 ± 0.9

2913 ± 1.1

Sodium

2602 ±1.2

2698 ± 0.8

2803 ± 0.7

2913 ± 0.8

Ammonium

2601 ± 1.3

2698 ± 1.4

2802 ± 0.1

2913 ± 0.1

Lithium

2602 ± 0.2

2700 ± 0.2

2804 ± 0.8

2914 ± 1.7

Barium .

2600 ± 0.9

2699 ± 0.8

2804 ± 0.6

2914 ± 1.8

Magnesium

"

2602 ±0.8

2700 ± 0.6

2802 ± 0.7

2912 ± 1.8

Aluminium

2603 ±0.4

2699 ± 0.9

2804 ± 0.9

2914 ± 0.7

Zinc

2602 ± 0.6

2699 ± 0.7

2802 ± 1.2

2912 ± 1.1

Cobalt .

2601 ± 0.2

2698 ±0.1

2803 ± 0.9

2912 ± 1.7

Nickel .

2603 ±0.6

2700 ± 0.7

2804 ±0.7

2913 ± 1.8

Cadmium

2600 ±0.1

2700 ± 0.2

2803 ±0.8

2913 ± 1.4

Copper .

2602 ±1.2

2699 ± 0.1

2803 ± 0.9

2913 ± 0.8

444

PRINCIPLES OF INORGANIC CHEMISTRY

The spectra of ten of these salts were photographed, the one

directly over the other, and the results are given in the accompanyieg
tigure (Fig. 42). The agreement between the position and character
of the bands is so striking, that there is no room fur doubt that
these salts show the same absorption bands.

Ostwald concluded from these results that the absorption spectra

of all the thirteen salts are exactly the same to within the limit of
error of measurement

This is one of the many beautiful confirmations of the conclusions
led to by the theory of electrolytic diaaociation.

One other element of the iron group is of special importance on
account of the number and variety of the compounds which it forma,
and the importance of soma of the substancea. This element ia
chromium.

CHAPTER XXXVI

CHROMIUM (At. Wt. = 52.1)

Chromium forms a number of series of compoQiids, and many of
these are closely related to iron and manganese. It occurs in nature
largely as chrome iron ore, which is ir(m chromite, having the com-
position Fe(Cr02)2. It also occurs as the lead salt of chromic acid,
PbCr04, or crocoisite.

Chromium is readily prepared by heating a mixture of chromic
oxide and carbon in an electric furnace. The lime which is added
decomposes the carbides of chromium, forming calcium carbide and
metallic chromium. Chromium is prepared most conveniently by
heating the oxide with finely divided aluminium, according to the
method of Goldschmidt.

Chromium is light in color, with a high lustre, and is not attacked
by oxygen at ordinary temperatures. It is very hard and does not
melt until a temperature of 3000® is reached. It dissolves in hydro-
chloric and sulphuric acids, but not in nitric acid. When treated
with acids it does not always dissolve continuously, but frequently
shows periodical or rhythmical phenomena. It dissolves, then be-
comes passive, dissolves again, is again passive, and so on. This
phenomenon, however, is not manifested alone by chromium.

Chromium forms alloys with a number of the metals, such as
aluminium and iron, and amalgams with mercury ; but these com-
pounds are without special interest.

Oxides of Chromium. — Chromium, like manganese, forms a num-
ber of oxides. Chromoua oxide, CrO, is formed by reducing the higher
oxides. It is a green powder, insoluble in water and most acids.
Chromic oxide or chromium sesquioxide, CrjOj, is formed by heating
the trioxide, or by heating chromic hydroxide. It is a green powder,
and when highly heated difficultly soluble in acids. It imparts a
green color to glass. It dissolves in alkalies, forming chromiles,
MCrOa. Chromium trioxide, CrO,, is formed by adding concentrated
sulphuric acid to potassium dichromate in very concentrated solution;

K,CrA+HjrS04 = K,S04 + H,0 -h 2 CrO,.
446

446 PRINCIPLES OF INORGANIC CHEMISTRY

It is a dark-red, beautifully crystalline substance, characterized
by its tremendous oxidizing power. When brought in contact with
organic compounds these are oxidized or burned up, as we say, and
chromium trioxide is reduced to a lower oxide of chromium.

A still higher oxide of chromium, Cr04, is supposed to exist in
solution, but has not been isolated.

Hydroxides of Chromium. — Just as chromium can form a number
of oxides, just so it can form a number of hydroxides. Cliromous
hydroxide^ Cr(0H)2, is formed when a chromous salt is treated with
a strong base : —

CrClj + 2 KOH = 2 KCl + Cr(OH)^

This is a yellow solid which quickly undergoes oxidation on the air.
Chromic hydroxide, Cr(0H)3, is formed by the addition of ammonia
or ammonium sulphide to a chromic salt : —

CrCl, + 3 NH4OH = 3 NH4CI + Cr (OH),,
2 CrCla + 3 (NH4),S + 6 H,0 = 3 H,S + 2 Cr (OH), + 6 NH^CI.

As ordinarily formed it contains two molecules of water, but this can
be easily removed by drying in a vacuum. It readily loses water,
forming the compound HCrOs, which is a weak acid.

The hydroxide H,Cr04, which is Cr(OH),- 2 H,0, is not
known in the free condition. This is chromic add, and its salts are
stable compounds. When it is set free from its salts by addition of
an acid, it loses water at once, forming the anhydride CrO,. Salts
of Siperchromic acid, HCr04, have also been described.

Valence and Properties of Chromium Ions. — It is obvious, from
the composition of the oxides and hydroxides of chromium, that this
element can exist in various conditions of valence. Bivalent chro-

mium ions, Cr, are distinctly basic, as is shown by the hydroxide. The

bivalent ions, however, readily pass into trivalent ions, Cr, which
are very weakly basic towards strong acids, and are acidic towards
certain bases. This is strictly analogous to the ions of iron and
manganese. Those of lower valence, or with the smaller electrical
charge, are basic ; but as the valence increases or as the amount of
electrical energy which they carry increases, the basic property
becomes less and less, and acidic properties begin to manifest them-
selves. When the valence of the chromium ion reaches six, as in
the compound H2Cr04, we have a very strong acid, chromic acid, and,
similarly, when it reaches seven in perchromic acid, HCr04. This is
analogous to iron, and especially to manganese, where the sexavalent

CHROMIUM 447

ion is acidic as in manganic acid, and the septivalent ion strongly
acidic as in permangaQic acid. We shall now study somewhat in
detail these several classes of chromium compounds.

Chromoiu Salts. — The chromous ion, Cr, combines with the anions
of acids, forming salts. These, however, readily absorb oxygen and
pass over into chromic compounds. The chromous compounds must,
therefore, be protected from contact with the air in order to preserve
them pure. They can be prepared by reducing the chromic com-
pounds with zinc and sulphuric acid ; also by the action of acids on
chromium. The yellow hydroxide has already been referred to.

Chromous ddoride, CrClj, is obtained by reducing chromic chloride
with zinc and sulphuric acid, or by heating chromic chloride in a cur-
rent of hydrogen. Its solution is blue, since the chromous ion, Cr,
is blue in color.

Chromous acetatef (CH3C00)jCr, is formed when sodium acetate
is added to a solution of a chromous salt. It is not readily soluble
and is, therefore, precipitated. It is dark red, crystalline, and fairly
stable, and can be used in preparing other chromous salts. Thus,
when chromous acetate is treated with concentrated hydrochloric
acid, the following reaction takes place : —

Cr(CH3C00), -h 2 HCl = 2 CH3COOH -h CrCljj.

Chromous chloride not being very soluble in hydrochloric acid,
can be obtained from this solution in crystals, which are blue in

color.

++

Chromic Salts. — Chromous chromium, Cr, readily passes, as we

+++
have seen, into chromic chromium, Cr. The chromic ion, while not

very strongly basic, unites with the anions of acids forming salts.

Chromic Chloride, CrClj, is obtained in the anhydrous condition by
passing chlorine over a heated mixture of chromic oxide and carbon.
It sublimes and crystallizes in plates of a violet color. Chromic
chloride dissolves very slowly in water unless a chromous salt is
present, when it readily dissolves, giving a green solution. It
crystallizes from the aqueous solution with six molecules of water,
CrClj.e H2O. When the salt with water of crystallization is heated,
it decomposes into hydrochloric acid and chromic oxide. This is
analogous to the conduct of most chlorides of weak bases, which
contain water of crystallization.

When the hydrated chloride is heated in an atmosphere of hydro-
chloric acid the water is given off and the anhydrous salt, which is
violet in color, is formed. This violet salt dissolves in water form-

448 PRINCIPLES OF INORGANIC CHEMISTRY

ing a green solution. If, however, the violet salt is sublimed it
recrystallizes in violet crystals, and these are practically insoluble
in water. It is, thus, obvious that there are two compounds, one
green and one violet. Both have been isolated.

From the green solution silver nitrate precipitates only two-thirds
of the chlorine. When the green solution is allowed to stand it
becomes violet, and then practically all of the chlorine can be pre-
cipitated by silver nitrate. In the violet solution we must have,

-H-+ — — —

then, the ions, Cr, CI, 01, 01, since silver is a reagent only for
chlorine ions. In the green solution one-third of the chlorine is not
present as such in the ionic condition, but must be united to the
chromium forming part of the cation. The green solution must

dissociate thus, OrOl, 01, 01. The vapor-density at 1200° corre-
sponds to the simple molecule CrOl,.

Chromic Sulphate, Cr2(S04),.15 H2O, is formed by dissolving chro-
mic hydroxide in concentrated sulphuric acid : —

2 Or(OH)3 + 3 H2SO4 = 6 HjO +0r,(SO4)s.

These crystals are violet in color and form a violet-colored solu-
tion. The salt Cr2(S04)8.9HjO can also be obtained from aqueous
solution. If the aqueous solution is boiled it becomes green, but
changes back slowly on cooling to violet. From the violet solution
all of the sulphuric acid is precipitated by barium ions. It must,
therefore, contain the ions, Cr, Or, SO4, SO4, SO4.

From the green solution, however, only one-third of the sulphuric
acid is precipitated by barium ions. Therefore, two out of every
three of the sulphuric ions, SO4, are in combination with the chro-
mium, forming part of the cation.

It has been shown that when the violet modification of the sul-
phate passes into the green one molecule of sulphuric acid separates
from every two molecules of chromic sulphate. This reaction can be
represented as follows : —

2 Or,(S04)3 + HjO = H3^04 + Or4(S04)40S04.

The complex green substance 0r4(SO4)4OSO4 dissociates into
Cr4 (804)40 and SO4. It is the sulphuric ion SO4 together with the
ion from the free sulphuric acid which is precipitated by barium ions.

When a mixture of chromic sulphate and sulphuric acid is heated,
the resulting solution gives no precipitate with barium ions, and

shows none of the characteristics of the chromic ion, Or. The

CHROMIUM 449

hydrogen ions of the sulphuric acid, however^ give normal reactions.

+++ -

These facts show that the chromic ions, Cr, and sulphuric ions, SO4,

are combined, forming complexes, which are not dissociated by water

into the simple ions.

Chromic sulphate, like aluminium sulphate and ferric sulphate,
combines with the sulphates of the alkalies forming double sulphates 01
alums. Of the chromium alums, the potassium salt, KCr(S04),. 12 HjO,
is the best known. The double sulphates with chromium manifest
the same general behavior as the simple chromium sulphate. The
violet solutions become green when heated, and the green solutions
have very different properties from the violet. They become violet
on standing, and from the violet solutions the alums crystallize again.

Chromites. — The hydroxide, Cr(0H)3, is a weak base, forming
salts, as we have seen. This substance dissolves in strong bases
showing its acid nature. It loses water and forms the compound
HCr02, which is chromous acid. There are a number of salts of
this compound known, and these are called chromites. Chromite
itself, which occurs in nature in abundance, is the ferrous salt of
chromous acid, Fe(Cr02)2.

Chromio Acid, H2Cr04. — In composition this acid resembles
sulphuric acid, manganic acid, and the like. It may be looked upon
as the partial anhydride of the hydroxide Cr(OH)e : —

Cr(OH)e=F2H20 + H,Cr04.
The compound H2Cr04 is, however, not known. When its salts
are treated with sulphuric acid ' the anhydride CrOs is obtained,
which is chromic acid minus water: —

H,Cr04 - H,0 = CrOs.

In this compound the chromium is obviously sexivalent, and
with its high valence the strongly acid properties begin to come
out. Chromic acid can he prepared also by electrolysis by the method
used by Morse in preparing permanganic acid,

Chromates. — Salts of this acid are formed when ferrous chromite
(chromite) is heated on the air in the presence of an alkali. The
chromium is oxidized to the chromate, which is soluble, and the
iron to ferric oxide which is insoluble in water. The mixture is
treated with water, when the chromate dissolves. If caustic potash
is used the resulting compound is potassium chromate. This is

deep yellow in color, due to the color of the chromic acid ion, Cr04.
It forms crystals which are isomorphous with potassium sulphate.

Barium chromate, BaCr04, is formed when barium ions, Ba,
2o

460 PRINCIPLES OF INORGANIC CHEMISTRY

come in contact with chromic acid ions, Cr04. It is bright yellow
in color, and is used as a pigment (j/ellow ultramarine). Lead
cJiromaie, PbCrO^, which is quite insoluble, is formed when lead ions

Pb, come in contact with chromic acid ions, Cr04. On account of
its fine yellow color and stability it is used as a pigment, and is
termed chrome yellow. The chromates are in general beautifully
yellow substances, silver being an exception, forming a red chromate.
Diohromates. — When a chromate like potassium chromate is
treated with an acid, the color of the solution changes from yellow
to red. The reaction in the case of potassium chromate may be
represented as follows : —

2 KjCrO^ + H^04 = H,0 + KjSO^ + K,CrA.

Potassium dichromaie, KjCrjOy, crystallizes from the solution in
beautiful red crystals, which often grow to unusual size and are of
unusual geometrical perfection. Potassium dichromate is a power-
ful oxidizing agent. It readily gives up oxygen, the chromium being
reduced to the chromic condition. When sulphuric acid is used we

have: —

2 KjjCr A + 8 H,S04 = 8 HjO + 3 0, + 4 KCr(S04)jj.

Potassium chrome alum is formed.

When hydrochloric acid is added to potassium dichromate there
are formed chromic chloride, potassium chloride, water, and instead
of oxygen chlorine is liberated. This is one of the most convenient
methods of preparing pure chlorine on a small scale.

When potassium dichrojnate is treated with caustic potash
potassium chromate is formed : —

KaCr^O; + 2 KOH = 2K2Cr04 -f H,0.

This is made evident by the change in the color of the solution
from red to yellow. It will be observed that the valence of the
chromium in potassium dichromate is the same as in potassium
chromate. The change from the former to the latter by the addition
of an alkali, and the reverse change by the addition of an acid,
therefore, involve neither oxidation nor reduction.

While the chromates of the heavy metals are in general very
insoluble compounds, the diohromates are soluble. If potassium
dichromate is added to a solution of a salt of a heavy metal the
dichromate is not formed, but the chromate, since it is insoluble.

Thus, if potassium dichromate is added to lead nitrate we have: —

2 PbCNOs)^ + KjCr A + H,0 = 2 PbCr04 + 2 KNO, -h 2 HNO,.
There are many cases known which are similar to the above.

CHROMIUM 451

The Ions Cri>4 and CrA- — We have seen that when a chromate is
treated with an acid, i.e. with hydrogen ions, it forms the dichromate.

This means that the chromic ions, Cr04, pass over into dichromic

ions, CrjO;. The reaction is expressed thus: —

2CH)4+H + H = H,0 + CrA.

This reaction always takes place whenever hydrogen ions are
present. It therefore takes place in the presence of pure water,
since pure water is dissociated to a slight extent.

This is in accord with the facts : Chromium trioxide dissolved

in water shows the red color of the CrjOy ions, and Ostwald has

shown by purely physical chemical methods that the ions CrJOy
exist in aqueous solutions of chromic acid.

When a dichromate is treated with an alkali it is transformed

into a chromate, i.e. the ions Cr207 are transformed into Cr04. This
transformation is effected by hydroxyl ions : —

CrA + OH + dH = H20-f2CK)4.

The reciprocal transformations of chromates and dichromates are
then only transformations of the ions Cr04 and CrjOj.

Chlorides of Chromic Acid. — When potassium dichromate is mixed
with sodium chloride and sulphuric acid, and the mixture distilled, a
dark-red liquid passes over which has the composition Cr02Cl,. This
can be regarded as chromic acid in which the two hydroxyl groups
are replaced by chlorine. From its analogy to sulphuryl chloride it
is called chromyl chloride. It is readily decomposed by water, and
easily gives up oxygen. The monoclilor derivative of chromic acid,
or chlorchromic acid, is not known in the free condition. The potas-
sium salt, KCrOgCl, is prepared by boiling a solution of potassium
dichromate with concentrated hydrochloric acid. From this solution
the orange-red salt crystallizes.

Perchromic Acid is formed by treating chromic acid with such
strong oxidizing agents as hydrogen dioxide. A solution of potas-
sium dichromate and sulphuric acid treated with a few drops of
hydrogen dioxide turns a beautiful blue color. This is supposed to
be due to the formation of perchromic acid. The compound formed
is not stable, since the blue color quickly disapi)ears, oxygen being
evolved. The compound, however, is more stable in ether, and when
the above solution is shaken with ether this acquires the beautiful

452 PRIXCIPLES OF INORGANIC CHEMISTRY

blue color of perehromie acid, which persists for a considerable time.
From the composition of the alkali salts this acid seems to have the
eomposition HCrO^

DeCeetion of CkrauiM. — CImHDiam is not precipitated from its
salts by hydrogen sulphide. It is precipitated by ammonium sul-
phide not as chromium sulphide but as chromium hydroxide, as we
have already seen. Chromium, therefore, belongs in the ammonium
sulphide group.

CHAPTER XXXVII

MOLTBDBNUM, TUNQSTBN, AND URANIUM
MOLYBDENUM (At. Wt. = 96.0)

Molybdenum is related to chromium in many respects, and it
seems advisable to study it in this connectioij. It occurs chiefly as
the sulphide, MoSj, molybdenite, and as lead molybdate, PbMo04,
ivulfenite.

The sulphides ai-e roasted and converted into the trioxide M0O3.
When the oxides are heated with carbon, or the oxides or chlorides
heated in a current of hydrogen, the element is obtained.

Molybdenum is steel gray in color, and melts only at enormously
high temperatures. It is not attacked by dilute sulphuric or hydro-
chloric acid, but readily dissolves in nitric acid. It combines with
the oxygen of the air only at high temperatures.

The chemistry of molybdenum is complex, on account of the large
variety of compounds which it forms. A few of these will be briefly
considered.

Oxides of Molybdenum. — Oxygen and molybdenum form three
compounds, MogOg, M0O2, and M0O3. The sesquioxide, M02O3, is
weakly basic, the dioxide has neither acid nor basic properties, while
the trioxide, M0O3, is the anhydride of molybdic acid. The trioxide
is formed when the sulphide is roasted. It is a white powder, prac-
tically insoluble in water.

Molybdic Acid, HSM0O4. — Molybdenum trioxide fused with an
alkaline hydroxide forms salts. These are salts of the normal acid
H2M0O4, or of polymolybdic acid derived from the normal acid by
loss of water. When ammonium molybdate is treated with nitric
acid, molybdic acid, HJM0O4.H2O, is obtained in crystals.

Molybdic acid readily forms complexes by losing water and
several molecules uniting. Thus: —

2 H,Mo04 = H,0 + HjMoA,

3 Hj^oO* = 2 H,0 + HjjMojOio.

Salts of these polymolybdic acids are well known. Molybdenum
trioxide forms complex acids by uniting with oUier acids. The best

458

454 PRINCIPLES OF INORGANIC CHEMISTRY

known is the compound with phosphoric acid — phoaphomolybdic
acid.

Ammonium molybdate treated with nitric acid gives free molyb-
dic acid : —

(NH4),Mo04 -h 2 HNO3 = 2 NH4NOS + H,Mo04.

This dissolves in an excess of nitric acid. If phosphoric acid is
added to this sohition, the very difficultly soluble ammonium phos-
phomolybdate separates as a yellow precipitate. This has the
composition (NH4)3P04.12Mo08.6n20, and is the ammonium salt
of phosphomolybdic acid, H3PO4.I2M0OS.I2H2O, which is obtained
by dissolving ammonium phosphomolybdate in aqua regia and allow-
ing the solution to evaporate. The free acid forms yellow crystals.
This compound is very impoiiant, both in detecting and determining
phosphoric acid quantitatively. The phosphoric acid is thrown down
at first as the phosphomolybdate. This dissolves readily in an excess
of ammonia, and from the ammoniacal solution the phosphoric acid
can be precipitated by " magnesia mixture " as ammonium magne-
sium phosphate. This, when heated, forms the pyrophosphate, which
can be readily weighed.

Compounds of Chlorine with Molybdenum. — Molybdenum forms
an unusually large number of compounds with chlorine. Molybde-
num dichlonde, MoClj, is formed by heating the trichloride in a
current of carbon dioxide. The trichloride, MoCl^, is formed by
reducing the pentachloride in a current of hydrogen. In appearance
it resembles red phosphorus. The tetrachloride, M0CI4, is also formed
when the trichloride is heated. It is volatile. Molybdenum penta-
chloride, M0CI3, is formed when molybdenum is heated in a current
of chlorine. It boils at 268° and melts at 194°. The molecular
weight, calculated from the vapor-density, corresponds to the for-
multt MoCla. The compounds of chlorine with molybdenum are,
then, M0CI2, M0CI3, M0CI4, and MoCla. This is a very unusual
series of substances.

Molybdenum also forms oxychlorides. Thus, we have MoOjCl^
^lOjOgCle, etc.

TUNGSTEN (At. Wt. = 184.0)

An element closely analogous to molybdenum is tungsten, which
is represented by the symbol W (Wolfram). The analogy is to be
found partly in the large variety of the compounds which the two
elements form with chlorine, and the relations between the composi-
tions of the two sets of chlorides.

MOLYBDENUM, TUNGSTEN, AND URANIUM 456

Tungsten occurs in nature as salts of tungstic acid. Thus, we
find calcium tungstate, CaW04, or scheelUe, iron tungstate, FeWOi,
or wolframite, lead tungstate, PbW04, or stolzite, and manganese
tungstate, MnW04, or hiibnerite. Tungsten is prepared by reducing
its trioxide with carbon or hydrogen, or far more conveniently by
means of finely divided aluminium. This is one of the elements
obtained by the Goldschmidt method.

The metal is sufficiently hard to scratch glass, and the compound
with carbon is harder than pure tungsten. It has a very high spe-
cific gravity, 16.5. It is not attacked by the oxygen of the air, and
only slowly by the common acids at ordinary temperatures. At
higher temperatures it forms the trioxide. Tungsten forms alloys,
especially with aluminium and steel. The presence of even a small
amount of tungsten in steel increases its hardness very considerably,
and tungsten steel has come into use for many purposes.

Chlorides of Tungsten. — Tungsten forms a number of chlorides,
which, in general, correspond in composition to the chlorides of
molybdenum. The trichlonde of tungsten, however, is not known,
while a hexachloride exists, and this has no counterpart among the
chlorides of molybdenum.

The dichloride, WClj, is obtained as the lowest reduction product
of the hexachloride in a current of hydrogen. The tetrachloride,
WCI4, is formed by heating the pentachloride in an indifferent gas.
The pentachloride, WCI5, is formed by distilling the hexachloride,
which gradually loses chlorine when volatilized. The hexachloride,
WClc is formed by heating the metal in a current of chlorine.

Such a series of compounds of an element with chlorine finds few
parallels in the whole field of chemistry. Molybdenum, as we have
just seen, forms, however, four compounds with chloriuQ,

Tungsten also forms the oxychlorides WOCI4 and WOjCV

Tungstic Acid, H2WO4. — Tungsten forms a dioxide, WOj, which, #
at elevated temperatures, readily takes up oxygen and forms tung-
sten trioxide, WOg. This is the anhydride of tungstic acid. When
a tungstate is treated with an acid at an elevated temperature, the
anhydride WO3 is formed. At lower temperatures tungstic acid,
H2WO4, is thrown down.

As we have seen, tungsten occurs in nature mainly in the form
of salts of this acid. When the oxide is dissolved in strong alkalies,
the corresponding salt is formed — M2WO4. A colloidal solution of
tungstic acid is obtained by treating sodium tungstate with hydro-
chloric acid, and dialyzing the mixture. The tungstic acid remains
behind in the form of a colloidal solution, which, on evaporation,

456 PRINCIPLES OF INORGANIC CHEMISTRY

forma a gummy mass. Tungsten also forms polytungstic CLcidSj b}
the polymerization of tungstic acid and the loss of one or more
molecules of water. Salts of the acid l^^fiy^ are known. Thij
is obtained from four molecules of tungstic acid by loss of thre<
molecules of water : —

4H,W04 = 3H,0 + HjW^Ou.

Tungsten trioxide combines with other acids like molybdenum
oxide, forming complex acids. The best known of these are the
compounds with arsenic, phosphoric, and iodic acids. These arc
known as phosphotungstates^ arsenitungatatesy iodatungstcUeSy etc.

URANIUM (At. Wt. = 238.5)

Uranium is characterized by the unusual variety of its com-
pounds, and by the great number of spectrum lines which it pro-
duces. It has the highest atomic weight of any known element,
and manifests the highest valency, being octivalent This degree
of valency is reached by only one or two other elements. It also
manifests a great variety of valencies, ranging all the way from
three to eight.

Uranium occurs in nature chiefly as the mineral uranite or pUchr
bleiide. This cousists chiefly of the oxide UjOg.

The metal can be prepared by heating this oxide with carbon in
an electric furnace ; also by reducing the chloride with sodium or
aluminium.

The metal has the color of silver, and the specific gravity 18.7.
Finely divided uranium combines with oxygen at about 200®, burning
readily in the gas.

Oxides of Uranium. — Uranium forms a number of compounds
with oxygen and hydrogen. These frequently show basic properties
towards strong acids, as well as acid properties towards strong bases.

Uranous oxide, UO,, forms salts with acids in which the uranium
is quadrivalent. Thus, we have uranous sulphate, U(S04)„ oxalate,
U(C\04)2, etc. When these salts are treated with an alkali the
compound U(0H)4 is precipitated.

When uranous oxide is heated it passes over into the compound
U,Os.

When the compound UjOg is treated with nitric acid, uranyl
nitrate, U02(NOs)2, is formed. When uranyl nitrate is heated, the
trioxide UOg is formed.

The compound U02(OH)2= II2UO4 acts as a base towards strong

MOLYBDENUM, TUNGSTEN, AND URANIUM 457

acids — the group UOj, known as the uranyl group, playing the part
pf a bivalent metal. Thus, we have uranyl sulphate, UO2SO4, uranyl
nitrate; U02(NO,)8, etc.

The higher oxidation products of uranium also have acid proper-
ties towards strong bases. The compound H,U04, or its derivatives,
dissolve i*eadily in strong bases, forming uranatea.

The uranates are prepared by adding a base to a solution of a
uranyl compound. These, however, are not derived from normal
uranic add, H2UO4, but from pyrouranic acid, H^jd^, which is
obviously obtained from two molecules of the normal acid by loss
of one molecule of water. The sodium salt NajUjOy is used as a
pigment for coloring glass under the name of uranium yellow.

Chlorides of TTraniom. — Uranium forms three chlorides: the
trichloride, UCI3, the tetrachloride, UCI4, and the pentachloride, UClj.
The trichloride is obtained by reducing the tetrachloride with hydro-
gen. The tetrachloride is obtained from the pentachloride, which is
formed by heating a mixture of uranium oxide and charcoal in a
current of chlorine.

TTraniom Eadiation. — Compounds of uranium when exposed to
light have the property of emitting an invisible radiation, which
traverses many substances impervious to light, such as black paper,
thin sheets of many metals such as aluminium, copper, etc. This
property is possessed by metallic uranium to from three to four
times the extent that it is manifested by the salts of this metal.

This is entirely different from the phosphorescence shown by
salts of uranium, since the latter disappears very quickly, while the
power of emitting this invisible radiation persists for years.

If a piece of uranium or of one of its salts is placed above a photo-
graphic plate covered with black paper or aluminium leaf, and
various substances are interposed between the uranium and the
plate, after several hours "radiographs" are obtained upon the
plate. These rays were also supposed for a time to be capable of
polarization by means of tourmalines. These phenomena would
suggest properties analogous to those possessed by light, and led
Stokes to conclude that the Becquerel rays occupy a position inter-
mediate between the Rontgen rays and light. He regarded the
Rontgen ray as made up of a great number in independent pulses.
In the Becquerel ray he thought that there was still irregularity,
but some regularity was beginning to manifest itself.

Later experiments, however, have shown that the uranium radia-
tion undergoes neither reflection, refraction, nor polarization.

This radiation is transmitted differently through screens of dif-

458 PRINCIPLES OF IXORGANIC CHEMISTRY

ferent substances, depending upon the angle in which they are
simultaneously placed in the path of the radiation. This would
indicate that the radiation is not homogeneous.

The uranium radiation discharges positive and negative charges
with equal speed, and its power to render a gas a conductor has been
shown by Rutherford to be due to an ionization of the gas. The
above and similar phenomena have been characterized as radio-
activity.

Other Eadioaotive Substances. — The discovery was made in 1898
by G. C. Schmidt that thorium, like uranium and its compounds,
can send out rays which are similar to the Rontgen rays. A little
later (1898) M. and Mme. Curie observed that certain uranium
minemls, such as pitchblende, were radioactive to a much greater
degree than metallic uranium or thorium. The conclusion was
drawn that in such minerals there are other radioactive substances
than uranium, and an attempt was made to isolate such substances.
Pitchblende was dissolved in acid, and hydrogen sulphide passed
into the solution. The sulphide of the active substance is insoluble
in ammonium sulphide, and was partially separated from the other
sulphides insoluble in this substance. Further, when the mixed
sulphides from pitchblende are heated to 700*, the active substance
sublimes into the cooler portion of the tube. The substance obtained
in this way was 400 times as active as uranium. This was further
purified by removing the bismuth until a much greater radioactivity
was shown. This substance was called poloniuMy after the native
country of Mme. Curie.

M. and Mme. Curie discovered a second radioactive substance in
pitchblende. This substance is obtained with the barium, from
which it is imj)ossible to eifect a complete separation. This sub-
stance is not precipitated by hydrogen sulphide nor ammonium sul-
phide. By dissolving the chloride in water and precipitating with
alcohol, a substance was obtained which had a radioactivity 17,000
times that of uranium. This substance they termed radium. The
spectrum was determined by Demarcay, and new lines were dis-
covered.

More recently Dabieme claims to have discovered a third radio-
active substance in pitchblende, which is closely allied to titanium
in its properties.

The rays from radium are much more intense than those from
polonium, uranium, or thorium. Rays from radium and polonium
produce fluorescence in barium platinocyanide, while those from
thorium and uranium are not sufficiently intense to excite this

MOLYBDENUM, TUNGSTEN, AND URANIUM 459

fluorescence. The radiation from polonium is much less penetrative
than that from radium.

Some of the rays from certain radioactive substances are deviated
by a magnetic field. Of these, a part are deviated the one way and
a part the other, showing that some are charged positively and some
negatively. The former ai-e known as a rays, the latter as fi rays.
Certain rays from radium are not deviated by the magnetic field.
These are much more penetrative than the deviable rays, and are
known as y rays.

Certain Eemarkable Properties of Badium. — The most recent
determination of the atomic weight of radium by physical means
gives the value 257.8. This shows that the radium atom has the
largest mass of any known atom. This would be expected, since the
other well-known radioactive substances — uranium, thorium, and
lead — have large atomic weights. Radium, being the most radio-
active, would be expected to have the highest atomic weight.

A very remarkable property of radium is that it maintains itself
at a temperature higher than that of the surrounding medium. This
constant development of heat energy has been shown by Rutherford
to come largely from the emanation^ which can be driven out of the
salts of radium by heat, and can be condenseH in glass tubes sur-
rounded by liquid air. The amount of heat evolved by radium, in
a given time, has been measured by Curie and Dewar by allowing
it to boil liquid hydrogen, and measuring the amount of gas set free.
They have shown that radium sets free enough heat to melt its own
weight of ice every hour.

The most remarkable property of radium, however, is that dis-
covered by Ramsay and Soddy. Helium is constantly being produced
from radium salts. This has since been confirmed by Curie and
Deslandres, and is now established beyond question.

This is the first authentic case on record of the transformation of
one elementary substance into another.

CHAPTER XXXVIII

COPPER, SILVER, GOLD
COPPER (At Wt. = 63.6)

There still remain three elements in the first group of the Periodic
System which have not thus far been studied. It is a defect in this
system that these elements fall in the first group, since they are not
closely allied to the remaining members from the chemical stand-
point. There are, to be sure, ceiiain analogies between copper, silver
and gold, and the alkalies, but there are analogies between almost
any two chemical elements. Were it not for the Periodic System
we should never think of dealing with the above three elements in
the same connection with sodium and potassium. When the Periodic
System leads us to connect elements as unlike as copper and sodium,
it is distinctly harmful, and detracts from, rather than adds to, our
scientific knowledge of these elements.

Occurrence and Preparation of Copper Copper occurs in con-
siderable quantity in the free condition. This is true on Lake
Superior, in Siberia, Japan, and elsewhere. Copper occurs in large
quantities as cuprous oxide, CuO, or cuprite^ cupric oxide, azurite,
and malachite the blue and green basic carbonates, as chalcocUe^ Cu^,
and chalcopyritey CuFeS^ This is also known as copper pyrites.

Copper is prepared from the oxides very simply by heating
with charcoal. From the sulphide it is much more difficult to
obtain pure copper. The sulphide of copper usually contains iron
sulphide and other impurities, and this still further complicates the
problem. The sulphides are roasted until the iron and a part of the
copper are converted into the oxide. When the roasted ore is heated
with sand and charcoal the iron oxide is partly reduced, forms
ferrous silicate with the sand, and disappears in the slag. The coi>-
per is for the most part in the form of the sulphide, but there is still
some iron sulphide present. This is known as matte.

The matte is again roasted, converting more of the iron sulphide
into oxide. It is again fused with sand, and this process repeated
until the iron is removed.

460

COPPER 461

The copper, which is now in the form of the sulphide, is partially
converted into the oxide, when the following reaction between the
sulphide and oxide takes place: —

Cu^ + 2CuO = S0, + 4Cu.

The copper thus prepared may be again heated with sand and char-
coal, again reduced, and so on until a fairly pure copper is obtained.

Copper is finally purified by means of electrolysis. The impure cop-
per obtained by the process just described is moulded into the form
of large, thick plates, known as the " anode plates." These are sus-
pended in a large bath of copper sulphate and connected with the
positive pole of a dynamo. Between these plates are alternately sus-
pended thin sheets of pure copper, which are the cathodes^ and these
are connected with the negative pole of the dynamo. When the
current is passed copper is deposited upon each of the cathodes, and
dissolves from each of the anodes. The action of the current is
really to carry the copper from the anode to the cathode opposite to
it, and deposit it upon the cathode.

Under these conditions the impurities are not deposited with the
copper, but either remain in solution or are deposited in the form of
a viscous mass on the bottom of the copper sulphate bath. These
anode "slimes," as this material is termed, are worked over for
various substances, and especially for gold and silver, which are often
present in considerable quantity.

Properties of Copper. — Copper differs in color from all other
metals, being a peculiar shade of red known as copper-red. Copper
is quite resistant to chemical reagents. In contact with moist air it
becomes covered with a green basic carbonate. When heated in the
air it forms the oxide. Copper is readily acted on by nitric acid,
but is not readily attacked by hydrochloric or sulphuric acid unless
it is hot. Copper is easily attacked by sulphur compounds, forming
the sulphide.

Copper does not decompose water until a white heat is reached,
and then only slowly. Consumed in appreciable quantities, copper
ions are poisonous.

On account of its physical properties, copper is one of the most
valuable of the metals. It can be readily hammered into thin
sheets or drawn into wire, and is very strong. Copper is not very
heavy, having a specific gravity of 8.9. It melts at 1057**, and
can, therefore, be easily cast. Next to silver copper is the best con-
ductor of electricity, and is extensively used in this capacity in con-
nection with telegraphy and telephony, and especially in connection

462 PRINCIPLES OF INORGANIC CHEMISTRY

with electric lighting and electrotraction, where large amounts of
electrical energy must be transported. This is one of the most
important uses of the element copper.

Alloys of Copper. — Copper forms a number of alloys with the
metals, which are very valuable. One of the best known is brass,
which is an alloy of copper and zinc, containing generally about
twice as much copper as zinc; but this varies greatly from one
specimen to another. Oerman silver or argentan, as we have seen, is
an alloy of copper, nickel, and zinc, while China silver is argentan
containing some silver. Copper also forms alloys with nickel and
silver. These are frequently used for coins. The silver coins usu-
ally contain about ten per cent of copper.

Among the best-known alloys of copper are the bronzes. The
ordinary bronzes are alloys of copper and tin, containing from ten
to thirty per cent of tin. Bronzes used for statues also contain zinc
Among the alloys of copper and tin are bell metal, spiegel bronze, etc.
Manganese bronze is an alloy of copper and zinc, to which manga-
nese is added. Phosphorus bronzes are ordinary bronzes containing
phosphorus.

Aluminium bronze is an alloy of copper, containing from six to
eight per cent of aluminium. It is of a yellow color, resembling
gold in appearance.

An alloy of copper and tin, containing about ten per cent of tin
and ninety of copper, is known as gun-metaL

Oxides of Copper. — Copper forms two compounds with oxygen ;
cuprous oxide, Cu^O, and cupric oxide, CuO. These are types of
the two classes of copper compounds — the cuprous compounds in
which copper is univalent, and the cupric compounds in which the
copper is bivalent. Copper, therefore, forms two kinds of ions; the

cuprous ion, Cu, and the cupric ion, Cu. Of these the cupric condi-
tion, in which the cop[)er carries two electrical charges, is the more
stable.

Cuprous oxide can be readily obtained by reducing an alkaline
solution of a cupric salt with a mild reducing agent, such as cane-
sugar. Cuprous oxide is a yellowish-red powder. When treated
with an acid, like sulphuric acid, it forms cupric ions and metallic
copper : —

CujO -f- H2SO4 = H,0 -f- Cu + Cu, SO4.

Cupric oxide is formed by oxidizing copper in the air or in a
stream of oxygen. Also by decomposing a cupric salt by heat. It is
a black powder, which readily gives up its oxygen to reducing agents.

COPPER 468

When cupric oxide is heated in a current of hydrogen it is readily
reduced to metallic copper, the hydrogen being oxidized to water.

Cnprio Hydroxide, Cti(OH)s, is formed by treating a cupric salt
with an alkaline hydroxide : —

CuClj 4- 2 KOH = 2 KCl + Cu (OH)^

Cupric hydroxide is light blue in color, easily passing over into
cupric oxide. When the liquid around the cupric hydroxide is
heated this transformation takes place — the blue hydroxide becom-
ing black in color, due to the formation of the oxide : —

Cu(OH)2 = H,0 + CuO.

Cupric hydroxide is a very weak base, forming salts with acids
which are strongly hydrolyzed. The hydroxide does not form
normal salts with weak acids, but basic salts. Although cupric
hydroxide is a weak base towards acids, it is not an acid towards
bases, as is generally the case. It does not dissolve in alkalies.

Chlorides of Copper. — Both the cuprous and cupric ions combine
with chlorine, and we have cuprous chloride, CuCl, and cupric chlo-
ride, CuClj.

Cuprous chloride is formed by reducing a cupric salt. When
copper sulphate is mixed with sodium chloride and sulphur dioxide
conducted into the mixture, cuprous chloride is formed. This
appears as a white, crystalline compound when the above solution
is poured into water. Cuprous chloride is also formed when hydro-
chloric acid is treated with an excess of copper, and the resulting
solution poured into water.

It readily combines with oxygen, passing into the cupric con-
dition, which is shown by the appearance of the blue color that is

characteristic of the cupric ion, Cu. Cuprous chloride boils at about
1000**, and a determination of the density of its vapor gives a molecu-
lar weight corresponding to the double formula CuaClj.

Cuprous chloride not only absorbs oxygen, but also carbon monox-
ide, forming the compound CujCl2.C0.2H,0.

Cupric Chloride, CnCls.2H20, is formed when cuprous chloride is
treated with chlorine, or when cupric hydroxide is dissolved in
hydrochloric acid : —

Cu(OH)j + 2 HCl = CuCl, 4- 2 H,0.

Cupric chloride crystallizes in blue needles containing two mole-
cules of water. When the water is driven off the anhydrous salt is
yellow or yellowish-brown.

464 PRINCIPLES OF INORGANIC CHEMISTRY

Hie solution of cupric diloride presents certain interesting cohr
phenomena. Dilute solutions are blue, like the dilute solutions of
all cupric salts. This is the color of the cupric ion. More concen-
trated solutions are green. This is due to the mixture of blue cupric
ions and yellow, undissociated molecules of cupric chloride. If con-
centrated hydrochloric acid is added to the green solution of cupric
chloride its color changes to yellowish-brown. By thus adding an
excess of chlorine ions the dissociation of the cupric chloride is
greatly driven back, according to the law of mass action, and the
color of the undissociated molecules of the salt makes its appearance.
Heat has the same influence as an excess of chlorine ions, driving
back the dissociation. When a blue solution of cupric chloride is
heated it therefore becomes green, and a green solution more and
more yellow. These color phenomena are of interest in connection
with the theory of electrolytic dissociation.

Cupric chloride combines with ammonia, forming complex com-
pounds, such as CUCI2.2NH8, CuClj.6NH„ etc.

Sulphides of Copper. — Cuprous sulphide, CujS, occurs in nature.
It is known as copper-glance. It is formed by reducing cupric sul-
phide in a current of hydrogen.

Cupric sulphide, CuS, is formed when hydrogen sulphide is
passed into a solution of a copper salt: —

CuClj + H,S = 2 HCl + CuS.

Copper sulphide is a black, amorphous powder, which is insolu-
ble in dilute acids. It is therefore precipitated from a solution of
a neutral copper salt by hydrogen sulphide. It readily takes up
oxygen, forming copper sulphate.

Copper Sulphate, CaS04.5H20. — This is the best known of all the
compounds of copper, and one of the best-known substances. Cop-
per sulphate or blue vUn'oly as it is called, is formed when sulphuric
acid acts on metallic copper. The hydrogen is not set free but acts
on more sulphuric acid, reducing it to sulphur dioxide: —

Cu 4- H2SO4 = CUSO4 4- H2,
II2SO4 -f Ha = 2 H,0 + SO2.

Copper sulphate crystallizes from the solution in the form of beauti-
ful blue crystals containing five molecules of water — CuSO^.SHjO.
When the sulphate is heated it loses four molecules of water at
100°, but the fifth is retained until 200° is reached. When the
last molecule of water is driven off the salt also loses some sul-
phuric acid.

COPPER 465

When in contact with a sulphate which crystallizes with seven
molecules of water, the salt, CUSO4.7H2O separates. Copper sul-
phate in solution is blue, and this is due to the color of the cupric
ions. The salt with water of crystallization is blue, probably due
to a slight dissociation of the salt in its water of crystallization.
Anhydrous copper sulphate is white, and this is the color of the
molecules of the salt When the white, anhydrous salt is dissolved
in water it is dissociated to a greater or less extent, and the blue
color of the copper ion appears.

Copper sulphate combines with ammonia, forming the compounds
CUSO4.4NH3, and CUSO4.2NH3. Copper sulphate in solution is
a good conductor of the electric current, and from such a solution
the copper is easily deposited electrohitically. This is a convenient
method of determining copper. We have seen that copper. sulphate
is used as the electrolyte in purifying copper by the electrolytic
method.

Copper Carbonate. — A soluble carbonate added to a solution of a
copper salt does not precipitate normal copper carbonate, but a basic
carbonate — Cu2(OH)2(C03). This is the composition of the mineral
maladiite, which has a beautiful green color and is used for orna-
mental objects. It occurs in large quantities, especially in Siberia.
Another less basic carbonate is known as azurite, having a beautifully
blue color. It has the composition Cu3(OH)2(C03)2. Azurite' is
also useful in making ornamental objects.

Other Copper Salts. — The acetate of copper has the composition
Cu(CH3C00)j, and is known as verdigris. The term verdigris has
also been applied to the basic acetate of copper. It is prepared by
the action of acetic acid on copper in the presence of the air. Cop-
per acetylency CU2C2, is the cuprous salt of acetylene, in which the
hydrogen ions are replaced by cuprous copper. It is prepared by
passing acetylene into ammoniacal cuprous oxide. It is an explosive
compound.

Copper ferrocyanide, Cu2Fe(CN)e, has already been referred to in
connection with the preparation of semi-permeable membranes for
measuring osmotic pressure. It is formed by adding a soluble cupric
salt to a solution of potassium ferrocyanide : —

2 CUSO4 + K4Fe(CN)3 = 2 K2SO4 + Cu2Fe(CN)3.

It is a reddish-brown, gelatinous mass, which, in color and general
appearance, resembles ferric hydroxide. As we have seen, it is the
best substance known with which to prepare semi-permeable mem-
branes for demonstrating and measuring osmotic pressure.
2h

466 PRINCIPLES OF INORGANIC CHEMISTRY

Precipitation of Copper by Ziao. — When a bar of zinc is immersed
in a solution of a copper salt, copper is precipitated upon the zinc,
and the zinc dissolves. This is due to the high solution-tension of
the zinc, and the low solution-tension of the copper. Zinc, having
a high solution- tension, sends ions into the solution, while copper
with its very low solution-tension is forced out of the solution.
There is a general principle involved here. A meted precipitates fix>m
their salta tlioae metals whicli stand beloio it in the tension-series, and is
2irecipitaJted by those metals which stand above it in the same series. If
the metals stand too close together in this series, they cannot precipi-
tate the lower member from its salts. There must be a considerable
diflFerence in position in the series in order to have pre?cipitation.
This principle is important to bear in mind, since it enables us to say
at once just what will happen when any metal is immersed in a salt
of any other metal. The simplest method of recalling the principle
is to remember that those metals with great solution-tension, pre-
cipitate from their salts the metals with small solution-tension.

Another Method of Ion Formation. — The precii)itation of one
metal from its salts by another metal furnishes us with a means of
forming ions. When zinc replaces copper from its salts the zinc
atom takes the charge from the copper ion, becoming itself an ion,
while the copper is converted into an atom : —

Zn -f- Cu, SO4 = Cu -f- Zn, s"04.

This is the third method of ion formation with which we have
had to deal (111 and 425). It can be formulated thus: An atom
takes the electrical charge from an existing ion^ becoming itself an ion^
tchile the fonner ion is converted into an atom.

There still remains one other method by which ions can be
formed. This will he taken up under gold.

CHAPTER XXXIX

SILVER AND GOLD

SILVER (At. Wt = 107.93)

We now come to the so-called "noble metals '* or "precious
metals." The well-known elements silver and gold will now be
studied.

Silver is not among the rare elements. It occurs in nature in
considerable abundance and in a number of compounds. The most
important of these is the sulphide, Ag^S, argentite, the double sul-
phide of silver and antimony, AggSbS,, pyrargyrUej the double
sulphide with copper, CuAgS, stromeyei'ite, and the chloride of
silver, AgCl, hornsUver. Silver also occurs in nature in the free
condition in certain localities in the United States, Norway, etc.

Preparation of Silver. — A large number of methods have been
devised and used for obtaining pure silver. The method employed
depends upon the nature of the silver ore which is being used.

If the chloride is employed, the silver is precipitated by means
of iron or lead.

If the sulphide is used, this is either roasted in the air and con-
verted into the sulphate, and the silver precipitated by metallic iron,
or it is converted into the chloride and the silver precipitated by
iron.

When silver is set free along with many other substances, it is
frequently dissolved in mercury and the mercury then distilled off.
This is known as the amalgamation process.

Lead ores, especially galena, usually contain silver, and silver is
frequently obtained mixed with lead. The solution of silver in lead
is concentrated by allowing it to crystallize. Pure lead separates at
first, and the remaining solution becomes richer and richer in silver.
A concentration of the silver in the lead is finally attained, where the
crystals which separate have the same concentration of silver as the
remaining solution.

When this stage is reached it is impossible to effect further separ
ration by crystallization. The above process of fractional crystalli-

467

468

PRINCIPLES OF INORGANIC CHEMISTRY

zation, known as the Pattinson process^ can be continued until tl
solution contains about one per cent of silver.

To effect further separation the lead solution of silver is heate
on the air. The lead is oxidized to litharge, and allowed to flo
away or be absorbed by the porous walls of the cupel in which tl
oxidation takes place. This process is known as ciipellation.

Another method of separating silver from lead is to fuse tl
latter with zinc (Parke's method). Silver dissolves readily in zin
forming a solid solution, and the zinc solution of the silver floats o
the lead and can be removed mechanically. The zinc can be di
solved out by means of dilute acids, or oxidized hot by means <
steam, leaving the silver behind.

Silver is prepared in pure condition by one of two processes
either by dissolving it in concentrated sulphuric acid and precipita
ing the metal by iron, or by the electrolytic process. This consisi
in casting the impure silver in anodes and using pure sheet silver a
the cathodes. The electrolyte is a nitric acid solution of silve
nitrate, to which copper nitrate is added to increase its conductivity
The silver separates upon the cathode in the form of beautifu
crystals.

Properties of Silver. — Silver is a white metal with a high lustre
It has the specific gravity 10.57 when distilled, and melts at abou
1000". It is not as hard as copper, and of all the metals is the bes
conductor of heat and electricity. It can be easily drawn into win
or hammered into thin foil. When melted in the presence of the aii
silver absorbs large volumes of oxygen ; indeed, as much as fifteen oi
twenty times its own volume. When the metal cools the oxygen is
given out, and this is known as the spitting of silver.

Silver is not easily attacked by chemical reagents. It is no1
attacked by the strongest alkalies even when hot, nor by dilute acids
with the exception of nitric acid. Concentrated nitric acid easilj
converts it into the nitrate, and concentrated sulphuric acid into the
sulphate, the acid being reduced to sulphur dioxide as with copper : —

2 Ag 4- 2 H,S04 = 2 H,0 + SO, 4- Ag^SO^.

Silver combines directly with sulphur, forming silver sulphide
This can be seen by holding a moist silver coin in a current of hydro-
gen sulphide. It also combines directly with the halogens at ordinary
temjKH'atures. Silver dissolves readily in a solution of potassium
cyanide.

Colloidal Silver. — A number of solutions of silver in water have
been described, which have all of the properties of colloidal solutions

SILVER AND GOLD 469

These have been prepared by Lea and others by reducing silver salts.
The citrate heated in a current of hydrogen, or reduced by ferrous
sulphate, yields a colloidal solution of silver in water. Lea prepared
solutions of silver which have very different colors and somewhat
different properties. From these solutions ordinary silver is easily
obtained. Another method of preparing colloidal solutions of silver
has recently been devised by Bredig. Two bars of silver are im-
mersed in water and their lower ends placed close together. An
electric current is passed between the bars, when metallic silver is
torn off in such a fine state of division in the water, that a drop of
the solution appears homogeneous under the microscope.

From such a colloidal solution of silver the metal is obtained by
boiling with hydrochloric acid. Such solutions have remarkable
catalytic action, as we shall see when we come to study platinum.

Alloys of Silver. — Silver forms a number of valuable alloys.
The alloy with mercury or the amalgam occurs in nature, and can
also be readily prepared by bringing the two metals into contact.
The alloy with copper, which is used in making coins, has already
been referred to. The alloy with aluminium can be used for solder-
ing aluminium.

Silvering Silver, as we have seen, is quite resistant to ordinary

chemical agents. It is, consequently, used for making utensils and
objects of ornament. These are, however, expensive, and silver-
plcUed wares are much used in their stead. These consist of brass,
copper, or other metallic objects covered completely with metallic
silver. They, therefore, have the properties of silver objects.

Silver-plating finds extensive application. The silver is deposited
by a number of methods ; the silver is reduced directly upon the
object to be plated, or it is applied mechanically and pressed upon
the object while hot. Another method is to apply the silver in the
form of an amalgam and then distil off the mercury; while still
another method which is extensively used is the electrolytic. Silver
is deposited electrolytically from a solution of the cyanide dissolved
in potassium cyanide.

Silver is now being extensively deposited upon glass in the con-
struction of mirrors. The glass surface must be entirely freed from
grease and all other impurities, and is then treated with ammoniacal
silver nitrate to which some caustic soda and a mild reducing agent
have been added. The reducing agents generally employed are alde-
hyde ammonia, or sugar of milk. The silver is slowly thrown out of
its salt and deposited uniformly upon the glass surface in the form
of a coherent layer with a bright surface. The silvering of glass is

470 PRINCIPLES OF INORGANIC CHEMISTRY

important in connection with optical apparatus, since such surfaces

are good reflectors of light. It is also finding increasing application

in connection with the preparation of ordinary mirrors, taking the

place of mercury, which is very poisonous.

Oxides and Hydroxide of Silver. — Silver forms three comx)ounds

with oxygen : the suboxide, Ag40, the normal oxide, Ag^O, and the

superoxide, AgO. It, however, forms only one hydroxide — AgOH,

which is stable only at very low temperatures. The hydroxide is

thrown down when an alcoholic solution of caustic potash is added

to an alcoholic solution of silver nitrate at a low temperature ( — 40**).

At all ordinary temperatures silver hydroxide loses water, forming

black silver oxide : —

2AgOH = H,0 4-Ag,0.

Silver hydroxide is a strong base, its salts not being hydrolyzed
by water. In this respect it resembles the alkalies. Silver oxide
is sufficiently soluble to give a strongly alkaline reaction. Such a
solution must contain silver and hydroxyl ions.

The silcer ion is univalent, combining with the anions of acids
and foruiing salts of the general type: AgCl, AgNO,, Ag^SO^, etc.
It is of interest to note that when atoms of silver pass over into
ions of silver a large amount of heat is absorbed — the reaction is
endothermic. The silver ion is especially a reagent for the halogen
ions. It combines with them, forming stable, insoluble compounds,
which we shall now study.

Silver Chloride, AgCl, is formed whenever a silver ion comes in
contact with a chlorine ion : —

Ag 4-01 = AgCl.

It is a white precipitate which quickly darkens when exposed to the
light. The darkening is due to the formation of a subchloride of
silver, AgoCl, or Ag4Cl.

Silver chloride is practically insoluble in water, and is conse-
quently used to determine quantitatively both silver and chlorine.
It is readily soluble in aqueous ammonia, and is thus distinguished
from the bromide and iodide.

Silver Bromide, AgBr. — Silver bromide is precipitated when sil-
ver ions come in contact with bromine ions : —

Ag -h Br = AgBr.

Silver bromide is white, with a slightly yellowish tint; is soluble
with diffioulty in ammonia and almost insoluble in water. Silver

SILVER AND GOLD 471

bromide is even more sensitive to light than silver chloride, and
upon this fact is based its use in photography.

Photography. — The science of photography is based almost
exclusively upon the action of light on silver bromide. The " sensi-
tive film" is prepared by adding ammonium bromide to gelatine,
and then adding silver nitrate in the dark. The following reaction
takes place : — ^^^^^ ^ ^^^^^ ^ NH^NOa -h AgBr.

The silver bromide is distributed through the gelatine in a very fine
state of division. The mass is then warmed until the precipitate
has become sufficiently coarse-grained — the coarser the grains the
more sensitive to the action of light. The soluble salts, ammonium
nitrate and silver nitrate, are removed by washing with water, and
the fused mass is then poured upon the surface of glass plates, to
which it adheres in the form of a thin film.

The plate containing the film is then exposed to the action of the
light from the object which it is desired to photograph. The time
of the exposure depends upon the intensity of the light and the sen-
sitiveness of the film. It may vary from several seconds to a hun-
dredth of a second, or even less. The action of the light is probably
to reduce the silver bromide to a sub-bromide of silver, although
this is not proved.

The exposed plate is now treated with a " developer," which con-
sists of some reducing agent, such as pyrogallic acid, ferrous sul-
phate, etc. The object of the developer is to reduce the silver
bromide depositing metallic silver, and the whole science of pho-
tography depends upon the fact that the silver bromide which is
most strongly illuminated, is most readily reduced by the developer.
Where the object was brightest, the plate is covered with a deeper
film of metallic silver, which becomes less and less dense as the
illumination is less and less. The result is a photograph of the
object with the light parts dark and the dark parts light. This is
the so-called " negative." The negative thus obtained still contains
unreduced silver bromide, since the portion of the salt which was
not exposed to the light is not reduced by the developer. If the
negative is exposed to the light in this condition, the silver bromide
would be acted upon, and the original picture would be destroyed by
superimposed images. To avoid this the negative must be fixed,
i.e. treated with a solution of a substance which will dissolve the
unreduced silver bromide. The fixing agent usually employed is a
solution of sodium thiosulphate, Na^SsOa, known technically as
" hyposulphite," or even as " hypo." This acts upon the silver

472 PRINCIPLES OF INORGANIC CHEMISTRY

bromide, forming the double salt, AgfifOi.2'NsigSJ^9, which is quite
soluble in water and is easily removed when the plate is washed
with running water. The negative is now " fixed," and ready to be
used in making " prints " or " positives." A " positive picture," or a
photograph proper, is obtained by placing the negative above paper
covered with the sensitive film and exposing it to light. The dark
parts of the negative cut off the light and appear bright on the posi-
tive picture, and, conversely, the light parts appear dark, since much
light passes through and acts upon the sensitive film upon the paper.
The positive, therefore, represents the lights and shades in the order
in which they occur in the object, and is a true picture in metallic
silver of that object.

In some cases the print is immersed in a bath containing a gold
or platinum salt, when the silver precipitates the gold or platinum,
itself passing into solution. Such photographs have, then, the soft
brown color of finely divided gold, or the harder, steel-gray tint of
finely divided platinum.

Photography not only finds extensive application in the arts,
but is of fundamental importance in scientific investigations. Many
epoch-making discoveriesL in physics, chemistry, astronomy, and biol-
ogy could have never been made without the use of the camera. As
an example, the whole science of spectrum analysis, as we know it
to-day, depends for its existence largely upon photography.

Silver Iodide, Agl, is formed whenever silver and iodine ions
come in contact. It is a yellow solid, insoluble in water and in
ammonia. Silver iodide, like the bromide and chloride, is sensitive
to light, and was formerly used in connection with photography.
Indeed, the earliest method of preparing photographs, devised by
Daguore, made use of silver iodide. A plate of silver was exposed
to the vapors of iodine, when it became covered with a layer of silver
iodide. It was then exposed to the light reflected from the object
to be photographed. The silver iodide was reduced, and strongest
where the illumination was greatest. The plate was then exposed
to the vapors of mercury. The mercury combined with the silver,
and appeared bright where the reduction was the greatest. By this
means a daguerreotype was produced, which was bright where the
illumination was greatest and dark where it was least.

Silver iodide has been practically abandoned in the preparatioki
of sensitive films, silver bromide being used almost exclusively.

Silver Nitrate, AgNOg, is formed when silver is dissolved in con-
centrated nitric acid. It crystallizes in beautiful, colorless plates,
melting at 200^ It is, therefore, frequently moulded into thin

SILVER AND GOLD 473

cylinders, and thus comes on the market under the name of lunar
caustic. When brought in contact with organic matter silver nitrate
is readily reduced, metallic silver being deposited. It also forms
insoluble compounds with albuminoids. We can now understand
why the hands are blackened by contact with silver nitrate, and
why it is used to cauterize small wounds and stop the flow of blood.

Silver Sulphide, Agjiy is precipitated as a black powder by the
action of hydrogen sulphide on a solution of a silver salt. It is in-
soluble in water and in dilute acids, and, therefore, silver belongs in
the class of elements whose sulphides are precipitated from neutral
salts by hydrogen sulphide. Moist silver combines directly with
sulphur and forms the sulphide, and hydrogen sulphide acts upon a
moist silver coin, producing the black sulphide. Silver is, there-
fore, frequently used to detect the presence of sulphur.

Silver Sulphate, Ag^O^ is iormed by dissolving silver in hot,
concentrated, sulphuric acid. It is only slightly soluble in water.

Silver Carbonate, Ag^COs, is formed whenever silver ions are

brought in contact with carbonic ions, CO3 : —

2 AgNOa + Na,COa = 2 NaNO, -f Ag,COa.

It is a yellow solid insoluble in water.

Other Compounds of Silver. — The silver salt of triazoic acid,
silver triaaoate, AgN,, is formed by adding the acid to a soluble silver

«alt : — AgNOa -l-HNa = HNOa + AgN,.

The salt resembles silver chloride in appearance, but is unstable and
very explosive.

Silver cyanide, AgCN, is formed as a white solid by the action of
silver ions on cyanogen anions : —

Ag, NOa-fK, CN = k, NO, + AgCN.

Silver cyanide readily dissolves in an excess of potassium cyanide, •
forming a double cyanide, which is soluble. This has the composi-
tion KAg(CN)2. This double salt is used for electroplating objects
with silver. In solution it breaks down into potassium and silver
cations, and cyanogen anions, and the silver is deposited as a uni-
form, coherent layer on the object which it is desired to cover with
silver.

Silver Sulphocyanate, AgSCN, is precipitated as a white solid

when silver and sulphocyanogen (SCN) ions come in contact : —
Ag, N0a4-NH„ SCN=NH4, NO, 4- AgSCN.

\

474 PRINCIPLES OF INORGANIC C5EMISTRY

Silver sulphocjanate is made use of in determining silver quan-
titatively. A standard solution of ammonium sulphocjanate is
added, drop by drop, to a solution of the silver salt containing some
ferric ions ^ (ammonium iron alum) and nitric acid. As soon as all
the silver is precipitated the sulphocyanogen ion reacts with the
ferric ion, forming iron sulphocyanate, which is characterized hy its
blood-red color. The appearance of this color shows that all of the
silver has been precipitated. Knowing the volume of the solutions
of ammonium sulphocyanate used, and its strength, we have all the
data necessary for calculating the amount of silver present. Such a
quantitative method is known as a volumetric method, and from its
discoverer as Volhard^a method.

Silver chromate^ A<2:,Cr04. is also useful in quantitative analysis
on account of its red color. When a solution of silver nitrate is
added to a solution of a chloride containing chromate ions (potassium
chromate), the color of the silver chromate will appear as soon as all
of the chloride is precipitated. In this way either the amount of
the silver ions or that of the chlorine ions can be determined by
having a standard solution of the other ion. This is known from
its discoverer as the MoJir method of determining silver.

GOLD (At Wt. = 197.2)

The oleiuent gold is one of the "noble metals," and is frequently
classed with platinum and allied elements. On the whole, however,
it seems best to study gold in connection with copper and silver.

Gold occurs in nature chiefly in the uncombined condition, in
the form of nuggets or grains in quartzite rocks or in sands. It
also occurs combined with the element tellurium, as the telluride.
When it occurs native it is by no means pure, containing silver,
copper, etc.

The Metallurgy of Gold. — Gold occurs usually in very small
quantities, and widely distributed through a large mass of rock or
sand. It must be obtained free from large quantities of foreign sub-
sttauces. This is ac^complished by placer mining and by vein mining.
In placer mining the earth or sand is washed with water, the light
materials being carried away and the gold left behind with the
heavier substances. In hydranlic mining the gold-bearing earth and
sands are washed down from the hills by water under pressure, and

1 A very larpre excess of the ammonium iron alum is added to make the
reaction more sensitive. This is an excellent example of the effect o/m<issas
utilized ill quanlitative analysis.

SILVER AND GOLD 476

the heavy gold collected by dissolving it in mercury. The gold is
obtained from the amalgam by distilling off the mercury. When
the gold occurs in veins in the quartz, this is finely powdered and
then treated with mercury. Gold amalgam is formed, and the gold
obtained by distilling off the mercury.

It not infrequently happens that the amalgamation process does
not work satisfactorily on account of the nature of the impurities
present with the gold. If arsenic is present the chlorination process
is used. This consists in treating the gold ore with chlorine, bleach-
iug-powder, and sulphui'ic acid ; the gold chloride formed being dis-
solved in water. Gold is obtained from the chloride by reduction
with ferrous sulphate or carbon, or by precipitation with hydrogen
sulphide as the sulphide, and heating the sulphide.

If other impurities are present, especially tellurium, the cyanide
process is used. This consists in treating the gold ore with potas-
sium cyanide, in which finely divided gold readily dissolves. The
gold is precipitated from the cyanide solution by means of metallic
zinc, or electrolytically, and then subjected to cupellation. Gold
thus obtained is impure and must be purified. The silver can be
removed by dissolving it out in nitric acid or concentrated sulphuric
acid. If the amount of gold in the alloy exceeds twenty-five per
cent, this does not work satisfactorily. In such cases the alloy is
fused with enough silver to dilute the gold to not more than one-
fourth. The process is therefore called quartaXion.

Another method of separating silver from gold is to dissolve the
alloy in aqua regia, and to treat the solution, after evaporating the
nitric acid, with a reducing agent such as ferrous chloride : —

3 FeClj -h AuClj = 3 FeClg -f Au.

Properties of Gold. — Gold is a soft, yellow solid, melting at 1064®
and forming a greenish liquid. It has a very high specific gravity —
19.3. Gold is extremely malleable, and can be hammered into leaves
not more than two-millionth s of a millimetre in thickness. Such
gold leaf is translucent and has a green color.

Gold is very resistant chemically, being attacked by comparar
tively few substances. Gold is not attacked by any of the strong
mineral acids. It dissolves in chlorine water, aqua regia, caustic
alkalies, nitrates, and cyanides.

The solution of gold in chlorine water is of special interest, since
it represents a fourth and the laM mode of ion formation. Gold has a
very low solution-tension, and, therefore, sends very few ions into

476 PRINCIPLES OF INORGANIC CHEMISTRY

solution. Chlorine water does not conduct the electric current^ an
therefore, the chlorine is not ionized. When the molecules of go]
come in contact with the molecules of chlorine, the former becon
cations and the latter anions : —

Au -h CI 4- CI + a = AiT, ci, ci, cl.

A molecule of a substance ichich can form cations comes in conta
with a molecule of a substance which can form anions, and both ai
ionized.

A colloidal solution of gold is readily prepared by reducing a d
lute alkaline solution of the chloride with formic aldehyde, and p
moving the crystalloids formed by dialysis. It is readily prepare
by the method of Bredig, to be described under platinum. Two bai
of gold are brought close together under water, and a considerabl
electric current passed between them through the water. The gol
is torn off in a very fine state of division, and there results a colloids
solution of the metal. The properties of such solutions will be d<
scribed more fully under platinum. A mixture of colloidal ^old an
colloidal stannic acid is known as the purple of Cassius.

Gold forms alloys with a number of the metals. The best know
and most important are the alloys with copper and silver. Pure gol
is too soft for use either as coin or as ornamental objects. To mak
it harder and more durable, copper is added. This gives to the gol
a deep-red color. The alloy containing ten per cent of copper is fn
quently used. The purity of the gold is expressed in carats, pur
gold being 24 carats. The number of carats means the number o
parts of gold in 24 parts of the alloy. Thus, 18-carat gold means a:
alloy containing 18 parts gold and 6 parts copper.

The alloy of gold and silver is extensively used instead of pur
gold, being more resistant to abrasion and more durable.

Gold Plating. — Metal objects are covered with gold in the sam
manner and for the same purpose that they are covered with silvei
Gold plating has been effected by a number of methods, but thes
have practically all given place to the electrical. The object to b
electroplated with gold is made the cathode, and a piece of pure goh
the anode, the double cyanide of gold and potassium being th<
electrolyte.

The object of plating the ordinary metals, such as copper, brass
etc., is twofold. The gold-plated metal has the appearance of solic
gold, with all of its attractive features. Further, such objects are re
sistant to chemical reagents, the covering of gold protecting the lesj
resistant metal beneath.

SILVER AND GOLD 477

Oxides and Hydroxides of Gold. — Gold forms the two oxides
Au,0 and Au^Oj, which are typical of the univalent and trivalent
compounds of gold. It also forms the corresponding hydroxides,
Au(OH) and Au(0H)8. Although these compounds are weak bases,
combining with the anions of certain acids and forming salts, the
auric hydroxide also has acid properties. Auric oxide and hydroxide
dissolve in caustic alkalies, forming aurates. These are salts of the
acid HAuOj, which is Au(0H)3 minus water : —

Au(OH), = H,0 4- H AuOj,
HAuO, -f NaOH = H,0 + NaAuOj^.

Salts of Gold. — Gold forms the aurous ion Au, and the auric ion An.
One of these carries one electrical charge, or is univalent, and the
other carries three electrical charges, or is trivalent These ions can
form salts with the anions of certain acids, and a few of these will
be considered. ^

The aurous ion, Au, combines with chlorine and forms auroua
chloride, AuCl. This compound is prepared by carefully heating
auric chloride to 180*. This decomposes into aurous chloride and
chlorine. Aurous chloride combines with the chlorides of the alka-
lies, forming double chlorides of the composition MAuCl^.

Auric chloride, AuCl^, is prepared by dissolving gold in aqua regia
and gently heating the resulting product to remove hydrochloric acid.
The compound formed by the action of aqua regia on gold has the
composition HAUCI4, and is known as hydrochlorauric acid. This
compound can be easily isolated in the form of yellow crystals, and
many salts of this substance are known. Thus, we have KAUCI4,
KaAuCli, etc. These have been regarded as double chlorides of gold
and the alkalies, but are well-defined salts of a well-characterized
acid, which can be readily obtained.

Gold forms two compounds with sulphur, aurous and auric sul-
phides. AuroiLS stdphide, Au^S, is formed by the action of hydrogen
sulphide on a hot solution of a salt of gold. It is light gray in
color. When the solution of the gold salt is cold, the compound
AujSj is precipitated as a black powder. Aurous sulphide dissolves
in alkali sulphides, forming compounds of the type MAuS, which are
soluble in water. The compound AUjSg dissolves in yellow ammo-
nium sulphide, forming the compound NH4AUSJ, which is soluble in
water. This is important in connection with the detection and
separation of gold.

CHAPTER XL
I.BAD, Tnr

LEAD (AW Wt. = 206i))

There remain two elements in group IV which have not been
studied. These are tin and lead. Although the atomic weight of
tin i$ less than that of lead, its chemistry is more complex, and it
will be considered after lead.

Oeevrace. Prgparatioa, and Pn^ertiat of Lead. — Liead occurs in
nature in a number of compounds. The most important is the sul-
phide PUS, or galena. It also occurs as the carbonate PbCO,, ceru^
fUe: the ohn>mate PR^rO^ crocoisite; the moljbdate PbMo04, ««^-
fiHiie : and in other forms.

Galena, being the principal ore of lead, is the one from which
most of the lead of commerce is prepared. Several methods are
emploveil to obtain the lead from the sulphide. One method is to
roast the sulphide, converting it into the oxide, and then reduce the
oxide with carbon : —

2PbS + 30, = 2PbO + 2SO„
PbO + C = CO + Pb.

A second niethoil is to roast the sulphide until a part of it is con-
verted into the oxide, and then heat the oxide with the sulphide :

PbS + 2PbO = 3Pb + SO^

In the above process a part of the lead is converted into the sulphate.
This also yields metallic lead when heated with the sulphide : —

PbS04 4- PbS = 2 SO, + 2 Pb.

The separation of silver from lead has been considered under
silver.

Lead is a very soft, bluish-white metal, which melts at 335**. It
has a fairly high specific gravity, 11.4. It becomes coated with a
layer of oxide when exposed to the air at ordinary temperatures, and
readily combines with oxygen when melted in the presence of the

478

LEAD, TIN 479

air. The action of water upon lead is of hygienic importance.
Pure water dissolves lead much more readily than ordinary, impure
water, such as that in springs, rivers, etc. The impurities react with
the lead and form a coating of carbonate, sulphate, etc., which pro-
tects the metal from the further action of the water. If the water
contains free carbon dioxide or organic acids, it acts upon the lead,
converting it into the salt of the acid in question. The hygienic
question is, whether drinking water should be conducted through
lead pipes. All things considered, it is much safer not to use them,
since pipes of other metals are practically unacted upon by water.

Lead does not dissolve appreciably in hydrochloric acid, since its
surface becomes quickly covered with a layer of insoluble chloride.
It dissolves to some extent in concentrated sulphuric acid, and the
sulphate is precipitated when the acid is diluted. Concentrated sul-
phuric acid which has come in contact with lead during its prepara-
tion, almost always contains some lead sulphate in solution. Nitric
acid and certain organic acids readily dissolve lead, converting it
into the corresponding salt.

Lead is readily/ precipitated from its salts by a number of metals.
This is especially the case with zinc and iron. When a bar of zinc
is suspended in a solution of a lead salt, the lead is thrown out and
the zinc dissolves. What takes place is a transfer of the electrical
charge from the lead ion to the zinc atom, converting the former
into an atom and the latter into an ion : —

Pb, NO3, NO3 -f Zn = Zn, NO3, NO3 -|- Pb.

The reason that this takes place is that zinc has an enormous solu-
tion-tension and lead a very small solution-tension (see p. 393). The
zinc atoms take the charge from the lead ions, becoming themselves
ions, and the lead, having lost its charge, is converted into atoms,
which are insoluble in water, and the lead is precipitated.

The lead frequently acquires very beautiful, tree-like forms when
it separates upon the zinc. This is known as Arbor Satumi, or the
lead tree. Lead forms a few alloys which are of value. Type metal
consists of 12 parts of lead, 3 parts of tin, and 5 parts of antimony.
Pewter is an alloy containing 4 parts of lead and 1 part of tin.

Oxides of Lead. — Lead forms a number of compounds with oxy-
gen. Lead suboxide, Pb^O, is formed as the first product of the
oxidation of lead, and by decomposing the oxalate at a low tem-
perature.

Ijead oxide, PbO, is formed by heating salts of lead, especially
the nitrate and carbonate. It is also known as litharge or moMicoL

480 PRINCIPLES OF INORGANIC CHEMISTRY

It is reddish-yellow or brown, depending npon the method of forma-
tion. It dissolves slightly in water, forming the corresponding
hydroxide, which will be considered a little later.

Miniumj Vbfi^ is formed by gently heating lead oxide on the air
(to d00''-400*'). On account of its bright red color it is known as
red lead. When minium is highly heated, it breaks down into
litharge and oxygen. When treated with nitric acid, minium partly
dissolves and partly remains behind as lead dioxide. It is regarded
as a mixture of 2 PbO and PbOf

Lead sesquioxidey Pb^Oj, is formed by oxidizing an alkaline solu-
tion of lead oxide with sodium hypochlorite.

Lead dioxide, PbO„ is formed by oxidizing the lower oxides of
lead by the action of nitric acid upon minium, as we have seen ; but
very readily by treating a lead salt with bleaching-powder. The
reaction with lead nitrate is: —

2 Pb(NO,), + Ca(OCl), -h 2 H,0 = CaCl, -f 4 HNO, + 2 PbO^

Lead dioxide is also formed when a solution of lead nitrate is electro-
ii lyzed. The dioxide separates at the anode.

Like dioxides in general, lead dioxide readily gives up its oxygen,
and is, therefore, an excellent oxidizing agent. When boiled with
hydrochloric acid, chlorine is evolved as with manganese dioxide.
When acidified and treated with hydrogen dioxide, oxygen is evolved
from both dioxides : —

PbO, + HA + 2 HNO, = 2 H,0 4- PbCNO,), 4- O^

Lead dioxide is the anhydride of an acid. When treated with
a strong alkali it dissolves, forming salts of the general type M,PbO,.

Hydroxides of Lead. — When a lead salt is treated with an alkali,
lead hydroxide is precipitated as a white, amorphous mass : —

PbCNOg), -f 2 KOH = 2 KNO3 + Pb(OH)^

It has basic properties forming salts with strong acids, and also
acid properties dissolving readily in strong bases. The salts have
the composition MjPbO^ and are known as jAumhites,

The hydroxide corresponding to lead dioxide, Pb(0H)4, also has

II acid properties, and is known as normal or orthoplumbic acid. The

ii metaplumbic acid is obtained from the ortho acid by the loss of one

i molecule of water : —

'' Pb(0H)4 = H,0 + HjPbOs.

' Salts of both of these acids are known as plumbates. We have

the calcium and potassium salts Ca,PbOi, KjPbOa, etc. The lead

LEAD, TIN 481

salts of these acids, PbjPbOi and PbPbOn, are the well-known oxides
PbsO^ and Pb,Os.

Chlorides of Lead. — Lead generally forms the bivalent ion Pb,
which readily combines with the anions of acids, forming salts that
are beautifully crystallized. It also forms the tetravalent ion Pb,
which can combine with certain anions of acids and form salts ; but
these are unstable. This is what we would expect, since we have
just seen that tetravalent lead and even bivalent lead can manifest
acid properties.

Lead chloride, PbCl^ is readily formed by bringing together
++ —

lead ions Pb and chlorine ions CI : —

Pb, NOa, NO3 + CI, Na -|- CI, Na = Na, NOa + Na, NOs -f PbCV

Lead chloride is a white, crystalline substance, somewhat soluble
in hot water, but only slightly soluble in cold water. Lead forms a
tetrachloride y PbCl4, but not by any direct method. When lead diox-
ide is dissolved in the most concentrated hydrochloric acid at a low
temperature, and ammonium chloride added to the solution, a yel-
low salt is obtained having the composition (NH4)2PbCla. When
this salt is treated with concentrated sulphuric acid, the following
reaction takes place : —

(NH,)3PbCla -h H^04 = (NH4)^04 4- H,PbCla.

The hydrochlorplumbic acid decomposes at once into hydrochloric
acid and lead tetrachloride : —

H,PbCl«=:2HCl + PbCl4.

Lead tetrachloride is very unstable, breaking down into lead chlo-
ride and chlorine. It solidifies at — 15** and is strongly hydrolized
by water, forming lead dioxide and hydrochloric acid.

Iodide of Lead, Pblj, is an especially beautiful substance. It is
formed by the action of a soluble iodide on a soluble lead salt : —

Pb(N08), + 2 KI = 2 KNO, -f Pbl^

It crystallizes from hot water and acetic acid in unusually beau-
tiful, glistening, yellow plates.

Lead Nitrate, PbCNOg),, is formed by the action of dilute nitric
acid on lead. It is not very soluble in strong nitric acid, and, there-
fore, the strong acid acts less vigorously upon the lead than the
weak. It is one of the few lead salts which are readily soluble in
water.

2i

482 PRINCIPLES OF INORGANIC CHEMISTRY

Lead Sulphide, PbS, occurs in nature, as we have already seen,
as galena. It is formed whenever bivalent lead ions come in con-
tact with sulphur ions — whenever a soluble lead salt is treated with
hydrogen sulphide : —

PbCNOa), 4- Hj^ = 2 HNO3 + PbS.

Lead sulphide is a black solid only slightly soluble in dilute acids.
Lead, therefore, belongs to those metals whose sulphides are precip-
itated from neutral salts by hydrogen sulphide.

Lead Sulphate, PbS04, occurs in nature as anglesite or lead vitriol
It is isomorphous with heavy spar or barium sulphate. It is formed

whenever lead, Pb, ions come in contact with sulphuric, SO4, ions. It
is very insoluble in water, and is, therefore, formed whenever a sol-
uble sulphate is added to a soluble lead salt : —

Pb, NO3 NOa + K, K, SO4 = K, NO3 -f K, NOa -h PbSO^.

Lead Persulphate, Pb(S04)2, or the lead salt of persulphuric acid,
is formed by electrolyzing strong sulphuric acid between electrodes
of lead.

Lead Carbonate, PbCOg. — The carbonate of lead is an important
compound, and especially the basic carbonates. The normal carbon-
ate occurs in nature as cerussite, and is isomorphous with aragonite,
a form of calcium carbonate. The normal carbonate is formed when
ammonium carbonate is added to a solution of lead nitrate: —

Pb(N03)2 + (^n,)jCO^ = 2 NH^NOa -f PbCO,.

If any other alkaline carbonate is used, such as sodium carbonate, a
basic lead carbonate is precipitated, and this is extensively used as a
pigment under the name of toliite had.

The old Dutch method of making white lead consists in placing
sheet lead, rolled into spirals, in porcelain pots containing a little
vinegar. The latter did not touch the lead. The vessels were then
placed in horse-manure, which decomposed and furnished the neces-
sary amount of carbon dioxide. The lead plates became covered in
time with a layer of basic lead carbonate, which was removed
mechanically.

This method, which consumed much time, has now given place to
some extent to quicker processes. Normal lead acetate is shaken
with litharge and water, when the basic acetate is formed. This is
then treated with carbon dioxide, when basic lead carbonate is
formed. The ordinary white lead which is used as a pigment is a

LEAD, TIN 483

mixture of basic carbonates of different composition. It must never
be used in painting objects which are exposed to the action of hydro-
gen sulphide, since it will turn black, due to the formation of lead
sulphide. ^^

Lead CSiromate, TbCr04, is formed whenever lead ions, Pb, and

chromic ions, Cr04, come together : —

Pb, KO3, NO3 -f K, K, CtO, = K, NO3 -f K, NO3 + PbCrO^.

The yellow lead chromate is difficultly soluble in water, and is
used as a pigment — chrome yellow. The basic chromate is yellowish
red and is known as chrome orange.

Lead chromate is also formed when a soluble lead salt is treated
with a soluble dichromate, since the chromate of lead is insoluble and
the dichromate soluble : —

2 Pb(N08), + KjCrA + H,0 = 2 PbCrO^ -f 2 KNO3 + 2 HNO,.

Lead Acetate, Pb(CH3C00)s 3 H2O, is an important soluble salt of
lead. It is formed by the action of acetic acid on lead oxide or
finely divided lead. On account of its sweet taste it is known as
sugar of lead. As we have seen, the acetate is an important inter-
mediate product in the foimation of basic lead carbonate, the latter
being formed when carbon dioxide is passed through a solution of
the acetate or basic acetate.

Basic acetates are formed when lead oxide is dissolved in normal
acetates. In contradistinction to sugar of lead, these are known as
vinegar of lead.

The Storage Battery or Accamalator. — One of the most impor-
tant uses of lead to-day, is in connection with the form of an electrical
battery know as the storage battery or accumulator. When a plato
of lead and one of lead dioxide are dipped into sulphuric acid, and
connected externally, we have a cell which is capable of furnishing
considerable electrical energy. Such a cell is formed by passing an
electric current between the lead plates covered with lead sulphate,
and immersed in sulphuric acid. The sulphate at the one plate is
reduced to lead, which is deposited upon that plate, and at the other
oxidized to lead dioxide, which is deposited upon the second plate.
We have thus converted electrical energy into chemical or intrinsic
energy. When the " charging current " is interrupted and the cell
closed, an electric current flows in the direction opposite to that used
in charging the cell. While the " discharging current " is flowing,
the lead at the one plate and the lead dioxide at the other are con-

4M PRINXIPLES OF INORGANIC CH£MISTRY

Terted into lead sulphate. The electrical energy produced comes
from the chemical energy which disappears. The chief source of the
electromotive force in a storage or secondary battery is the transfer-

mation of quadrivalent lead ions Vh, into bivalent lead ions Pb.
The quadrivalent ions are furnished continually by the lead dioxide.
They pass into bivalent ions, giving up two electrical charges, and
form with the sulphuric ions, ^4, lead sulphate. By means of these
reciprocal transformations of lead, lead sulphate, and lead dioxide, it
is possible to convert electrical energy into chemical, or " store" it
as we say ; and then to reconvert the chemical or intrinsic energy
into electrical energy at wilL The great weight of the lead plates
is a serious disadvantage in the storage battery, and an attempt is
now being made to use iron and nickel instead of lead.

TIN (At Wt. = 119.0)

An element chemically allied to lead, but differing from it in many
respects, is tin. This metal is useful on account of its chemical
inactivity, and is valuable on account of its properties and because it
does not occur in large quantities. Tin occurs in the uncombined
condition along with gold, but chiefly as tin dioxide, SnO,, known as
tinstone or cassiterite. The chief localities for tin are Cornwall in
England, Siberia, and the East India Islands.

Preparation and Properties of Tin. — The sulphur, arsenic, and
similar impurities are removed from the tin ore by roasting, and the
oxide is then roasted with carbon. The oxide is readily reduced to
the metal, and this is purified by repeated melting ; the molten tin
being poured off from the less easily fusible alloys. The purest tin
comes from the island of Banca, and is known as Banca tin. Another
comparatively pure form is known as block-tin.

Tin is light in color and quite soft. It can be readily ham-
mered or rolled into thin sheets known as tin-foil, which is used
for covering objects to protect them from the action of air, moisture,
etc. It is crystalline, and the movement of the crystals over one
another produces a crackling noise known as the cry of tin. Tin melts
at 233° and volatilizes at 1500°. It has a specific gravity of 7.3.
While tin is malleable at ordinary temperatures it becomes brittle
above 200°. Tin is not very readily attacked by reagents. It is not
acted on chemically by the air or moisture. It dissolves in concen-
trated hydrochloric and sulphuric acids, forming the corresponding
salts. Concentrated nitric acid oxidizes tin to metastannic acid.
Tin dissolves in caustic alkalies, forming salts of stannic acid. On

LEAD, TIN 485

account of its resistance to ordinary reagents tin is frequently used
for covering objects which are to be used in the kitchen, or vessels
for holding fruit or water. Indeed, in distilling water block-tin con-
densers are frequently employed, since water-vapor is practically
without influence upon them. Iron objects are plated with tin by
first cleansing them by washing with an acid, and then dipping them
into molten tin.

Allotropic Forms of Tin. — Tin occurs in allotropic modifications.
At ordinary temperatures the white modification with which we are
so familiar is the stable form. Cohen showed that below 20^ the
white modification is unstable, and this passes over slowly into a
gray modification which is the stable form at low temperatures.
The transformation temperature is 20^ Just below this tempera-
ture the gray modification is formed slowly ; the velocity of the trans-
formation into the gray modification increasing until a temperature
of — 48'' is reached. The velocity of the transformation into gray tin
at temperatures slightly below 20° is increased by the presence of a
little gi-ay tin or pink salt, SnCl4.2NH4Cl. Gray tin is brittle,
readily crumbling to powder, and has a smaller specific gravity
(5.8) than white tiu. This explains the crumbling of tin known as
tin-peat, which occurs in tin organ pipes and other tin objects in cold
countries.

The brittle modification of tin which exists at elevated tempera-
tures is probably another allotropic form.

Alloys of Tin. — Molten tin dissolves readily in most of the other
metals in the molten condition, and forms a large number of alloys
with them. Some of these are very important substances. Soft
solder is an alloy of tin and lead. The bronzes, as has already been
stated, are alloys of tin. Mannlieim gold is an alloy of zinc, copper,
and tin. Britannia metal is an alloy of tin and antimony, containing
one part of antimony to nine of tin. Tin readily forms an amalgam
with mercury, and this is used for plating glass mirrors. Tin-foil is
coated with mercury and a glass plate placed upon it. The tin amal-
gam which is formed adheres tightly to the glass. This method of
making mirrors is being rapidly abandoned and silvered surfaces
used instead, as has already been mentioned.

The Tin Ions. — The element tin forms two kinds of ions, those
++ ++-»■+

which are bivalent, Sn, and those which are quadrivalent, Sn. Of

these the quadrivalent ion is the more stable, the bivalent tending to

pass over into it. No other ion of tin is known. The stannous ions

passing over into the stannic, take up readily two electrical charges,

or are good reducing agents, as we say.

486 PRINCIPLES OF INORGANIC CHEMISTRY

StannooB (SnO) and Stannio (SnOj) Oxides. — The two oxid^ cor-
responding to the stannous and stannic conditions are known.
Stamiou8 oxide, SnO, formed by heating stannous hydroxide in a
current of carbon dioxide, is comparatively unstable, readily com-
bining with oxygen and forming stannic oxide, SnO,- Stannic oxide,
SnO^, occurs in nature as tinstone. It is also readily prepared by
burning tin in the air. Tin dioxide can be obtained in a number of
crystalline forms.

StannooB (Sn(0H)2) and Stannio (Sn(0H)4) Hydroxides. — Stan-
nous hydroxide is precipitated from stannous salts by the addition
of an alkali : —

SnCla + 2 NaOH = 2 NaCl + Sn(OH)j|.

Stannous hydroxide readily dissolves in an excess of the alkali,
forming salts. These, however, are unstable, readily passing over
into salts where the tin is tetravalent, metallic tin separating from
the solution.

Stannic Hydroxide, Sn(0H)4, is formed by treating tin with nitric
acid of medium concentration. It readily loses water, forming meta-
stannic acid, HjSnOg, which is soluble in alkalies forming metcLstan-
nates, Metastannic acid is insoluble in water, and also in acids.

Another compound having the composition HjSnOa is known,
which is soluble in acids. It is formed by boiling stannic chloride
with water, or by treating stannic chloride with ammonia. It is also
formed by treating the potassium salt with an equivalent of an
acid : —

KjSnOs + 2 HCl = 2 KCl + H^SnO,.

It dissolves readily in alkalies forming salts, which are known as
stannates. To distinguish it from the isomeric metastannic acid, it
is known as stannic acid. The latter passes over slowly into the
former.

Stannous Chloride, SnCls.2H20. — The best-known stannous salt
is the chloride, SiiCl^. It is formed by dissolving tin in hydrochloric
acid, also by heating tin in a current of dry hydrochloric acid gas.
Stannous chloride crystallizes with two molecules of water. It
melts at 250° and boils at 610°. On account of the marked ten-
dency of the stannous ion to pass over into the stannic, it readily
undergoes superficial oxidation. On account of this same tendency,
stannous chloride is an excellent reducing agent, and is frequently
used where mild reductions are desired, especially in organic chem-
istry. It combines readily with chlorine, as we would expect, form-
ing stannic chloride. So great is this power that it removes the

LEAD, TIN 487

chlorine from the chlorides of mercury, precipitating metallic
mercury : — ^^^^ _^ ^^^^^ = Hg + SnCl^,

2 HgCl + SnClj = 2 Hg -h SnCl^.

Stannous chloride combines directly with free chlorine, forming
stannic chloride : —

Sn, Ci, CI + Clj J'Sn^ CI, Ci, CI, CI.

This is another example of that mode of ion formation where a
cation takes on more positive electricity, at the same time converting
an atom into an anion.

Stannous chloride combines also with hydrochloric acid, forming
hydrochlorstannou8 acids. The salts of two such acids are well
known, and show that the acids have the compositions HSnClj and
H2SnCl4. Stannous chloride is known commercially as tUi-^alt.

Stannic Chloride, SnGl4. — The tetiachloride is formed, as we have
seen, by the action of mercuric chloride on stannous chloride or on
metallic tin. It is also formed by the direct action of chlorine gas
on tin. It is a liquid boiling at 114°, and fuming in contact with
the air. It is known as spiritus fumans Libavii, When brought in
contact with a little water, it forms a viscous mass known as tin-
butter, having the composition SnCl^.SHjO. With much water
stannic chloride is readily hydrolyzed, forming the hydroxide and
hydrochloric acid. The hydrochloric acid set free as the result of
the comparatively slow hydrolysis, produces the rapidly increasing
conductivity which is observed when stannic chloride is dissolved in
water. Stannic chloride combines directly with hydrochloric acid,
forming hydrochlorstannic acid, HgSnCle.GHjO, which can be
obtained in the free condition. Salts of this acid can be obtained
by bringing alkaline chlorides in contact with stannic chloride. The
ammonium salt, (NH4)2SnCl6, is known in commerce as pink salt.

Sulphides of Tin. — Stannous sulphide, SnS, is formed by con-
ducting hydrogen sulphide into a solution of stannous chloride : —

SnCla + HjS = 2 HCl + SnS.

Stannous sulphide is a brownish-black powder, insoluble in
dilute acids, but soluble in ammonium poly sulphide, forming am-
monium stdphostannate, in which the tin is in the quadrivalent

condition : — ^ ^ ^^^ ^ ^ ^^„ ^ ^ ^

SnS -h (NH4),S, = (NH4)jSnSj^

Stannous sulphide is also formed by melting tin with sulphur.

Ul.li

488 PRINCIPLES OF INORGANIC CHEIOSTRT

Stannic sulphide is precipitated as a light yellow powder wl
hydrogen sulphide is passed into a solution of a stannic salt o
stannate : —

SnCl4 + 2 H,S = 4 HCl + SnS,.

It may also be prepared by heating together tin, sulphur, and anu
nium chloride. Prepared by the latter method it is crystalline, ]
has a golden-yellow color ; it is known as mosaic gold and used a
pigment and for gilding.

Stannic sulphide dissolves readily in ammonium sulphide, fo
ing sulphostannates or thiostannatea : —

SnS, -h (NH^)^ = (NHOtSnS,.

This reaction is of importance in the separation of tin from n
of the elements. It will be remembered that in addition to tin,
sulphides of only arsenic, antimony, gold, and platinum are solu
in yellow ammonium sulphide. By this reaction arsenic, antimo
tin, gold, and platinum are separated from the remaining elemei

CHAPTER XLI

I. RX7THENIX7M RHODIUM PALLADIUM
n. OSBfflUM IRIDIUM PLATINUM

Although classed together these two groups of three elements
each have certain properties which are markedly different The
most striking is the atomic weights. The iirst three have atomic
weights which are not widely removed from one hundred, while the
atomic weights of the last three are not widely removed from two
hundred. Similarly, the specific gravity of the first three elements
is close to twelve, while the specific gravity of the last three is
about twenty-two. We shall take up first the lighter elements, —
ruthenium, rhodium, and palladium.

RUTHENIUM (At. Wt. = 101.7)

Ruthenium occurs in nature with platinum. When the platinum
is dissolved in aqua regia^ ruthenium remains behind undissolved.

Ruthenium was discovered in 1845 by the German chemist Claus.
The metal is steel-gray in color and has a specific gravity of 12.2.
It fuses only at very high temperatures, certainly not below 1800®.

Ruthenium combines with oxygen, forming the compounds RuO,
RujOs, RuOj, RuOj, and RUO4. The higher oxides are anhydrides of
acids, which will be referred to a little later.

Ruthenium is quite resistant to aqua regia, but combines with
chlorine, forming three compounds, — RuClj, RuCla, and RUCI4.

The trichloride combines with hydrochloric acid, forming HjRuCls,
and the tetrachloride with hydrochloric acid, forming HjRuCV

Ruthenium dissolves in fused caustic potash, forming two classes
of salts. When fused with potash and potassium nitrate, potassium
mthenaie, K2RUO4, is produced. This salt forms dark crystals which
dissolve, and give a deep orange-<5olored solution. When the ruthe-
nate is treated with a dilute acid it passes over into the perruthenate
KRUO4. In the ruthenates and perruthenates we have obviously
the analogues of the manganates and permanganates.

489

490

PRINCIPLES OF INORGANIC CHEMISTRY

RHODIUM (At. Wt.= 103.0)

The element rhodium occurs in very small quantities, and usi
with platinum and gold. It has unusually valuable properties, :
ing higher than platinum and being unacted upon by acids and
by aqua regia. When platinum is alloyed with a considerable c
tity of rhodium, it becomes much more resistant to chemical reag
An alloy of platinum and rhodium, containing from 30 to 40
cent of rhodium, is not acted upon by aqua regia^ Such an j
fuses higher than platinum, and is more valuable than pure platii

Rhodium acts mainly as a trivalent element, forming such
pounds as Rh(0H)8, ^hCls, etc. Rhodium also forms cbaractei
double chlorides of the type MsRhCle, where M is an alkali met

PALLADIUM (At. Wt. = 106.5)

Palladium occurs with platinum, but the chief source of p
dium is a gold ore in Brazil. The ore is fused with silver and
treated with nitric acid. The silver and palladium dissolve, and
silver is separated from the palladium by precipitation, as the <
ride. The palladium can then be precipitated as the cyanide, w
is ignited or thrown out of the solution directly with metallic zi

To chemical reagents, palladium is the least resistant of all
platinum metals. It dissolves readily in nearly all of the sti
acids. It is white like silver and fuses at about 1500**, whici
lower than any of the other platinum metals.

Palladium is best known chemically in connection with
remarkable power to absorb hydrogen gas. Metallic palladiun
100° can absorb more than 800 volumes of hydrogen gas. It
been supposed for a long time that there was a compound fon
having the composition PdjH and known as lyalladium hydride. S^
doubt has recently been thrown upon, the existence of such a c
pound by studying the absorption of hydrogen by palladium in te:
of the phase rule.

Palladium absorbs more than 600 volumes of hydrogen, wi
would correspond to the formation of this compound, but this i:
be simply a solution of hydrogen in the supposed compound Pd
Van't Hoff has shown that if hydrogen under increased pressur
brought in contact with palladium hydride, the hydrogen is absorl
in terms of Henry's law, i.e. proportional to the pressure of the {
This he has cited as an excellent example of a solid solution o
gas in a solid, — hydrogen in palladium hydride.

OSMIUM 491

Palladium hydride readily gives up its hydrogen to substances
which can be reduced, and is, therefore, an excellent reducing agent.
The halides are reduced by it to the corresponding acids.

Palladium hydride gives up all of its hydrogen when heated for a
time to a red heat.

Palladium forms the oxides PdO and PdO^ It also forms the
chlorides PdCl, and PdCl4. These are typical of the palladous and
palladic compounds. In the former the palladium is bivalent and
in the latter tetravalent.

Palladic chloride may be regarded as combining with hydro-
chloric acid and forming hydrocklorpalladic acid, HjPdCl«. This
compound is formed when palladium is dissolved in an excess of
aqua regia. Salts of this acid are obtained by treating palladic
chloride with the chloride of an alkali : —

PdCl4-h2KCl=KsPdCl«.

The potassium and also the ammonium palladic chlorides are only
slightly soluble in water. In this respect, as well as in their compo-
sition, they resemble the corresponding compounds of platinum, as
we shall iee.

^ OSMIUM (At Wt. = 191.0)

We now come to the three heavy metals of the platinum group.
The first of these, osmium, has the specific gravity 22.5, and is,
therefore, the heaviest of all known elements. It occurs in platinum
ores as an alloy with iridium. When these ores are treated with
aqua regia, neither the iridium nor the osmium dissolves. Osmium
like ruthenium forms a volatile oxide, OSO4, and this property is util-
ized to separate osmium from iridium.

Osmium is characterized by its high melting-point. It does not
melt at the highest temperatures thus far produced, and attempts
have been made to substitute osmium wire for the carbon filament
in the electric light The alloy of osmium with iridium finds appli-
cations on account of its resistance to chemical reagents, and on
account of its hardness and resistance to mechanical abrasion.

Heated in contact with the air, osmium combines with oxygen,
forming osmium tetroxide (OSO4), which is easily volatile at 100^ It
is easily reduced by organic matter to metallic osmium, and is, there-
fore, used in the preparation of microscopic sections of tissues. For
the same reason it is very poisonous when inhaled, producing great
irritation of the mucous membrane of the eyes, nose, and throat
The aqueous solution of osmium tetroxide is perfectly neutral, and
the compound has been erroneously called osmic acid. Osmum forms

I;'

492

PRINCIPLES OF INORGANIC CHEMISTRY

salts of the composition MjOsOi, but the corresponding oxide, <
is not known.

Osmium combines with chlorine, forming the dichloride, C
and the tetracMoride, O8CI4. Osmium also forms salts of the gei
composition MgOsCle.

IRIDIUM (At. Wt. 193.0)

Iridium occurs, as already mentioned, along with osmiur
osniium-iridium in platinum ores. It is also combined, in part,
the platinum as platinum-iridium. The separation of osmium ;
iridium has already been referred to. Iridium is separated :
platinum by dissolving the two in aqua regia, and adding ammoi
chloride. A double chloride of iridium and ammonium, (NH4)J
is formed, which is readily soluble in water, but not in water ,
rated with ammonium chloride. Metallic iridium is light in c
and does not melt until the temperature of the flame of the
hydrogen blowpipe is reached. It has a specific gravity of 22.4,
is, therefore, next to osmium, the heaviest of all known element;

Iridium is more resistant to chemical reagents than platii
and alloys of iridium and platinum are used, instead of pure platii
for making chemical apparatus such as crucibles, dishes, etc.
standard metre in Paris consists of 90 per cent platinum and IC
cent iridium ; and this particular alloy has been found, on the vr]
to be the best in preparing measuring apparatus.

Iridium is insoluble in even the strongest mineral acids,
dissolves slowly in aqua rcgia only when in the finely div
condition.

Iridium acts toward oxygen as a trivalent and quadriva
element, forming the oxides IrjOs and Ir02. The correspond
hydroxides, Ir(()ll)3, and lr(0H)4, are known. Iridium combines 1
chlorine, forming the compounds IrCl2, IrClg, and IrCl4. Towj
chlorine it, therefore, acts as a bivalent, trivalent, and quadriva
element. Like the other members of the platinum group, irid
forms chloro acids' and alkali salts of these acids. Thus we h

MJrClc, in which the ion, IrClg, is bivalent, and also MjIrCl^

which the ion, IilHo, is trivalent. Ammonia forms with iridiun
with the other platinum metals complex compounds.

PLATINUM (At. Wt. = lOl.P)

Platinum, one of the most valuable elements from the standp^
of chemistry on account of its power to resist chemical reage:

PLATINUM 498

occurs fairly widely distributed, but not in large quantities. The
countries in which platinum occurs in the greatest quantity are
California, Borneo, and the Ural Mountains. It occurs in company
with the other platinum metals, and is separated from them by
methods with which we are now more or less familiar.

Properties of Platiniun. — Platinum is light in color, has a specific
gravity of 21.4, and melts at about 1770° in the flame of the oxy-
hydrogen blowpipe. It is both ductile and malleable — can be drawn
into thin wire and hammered into thin sheets.

Platinum can exist in a number of physical conditions. In addi-
tion to ordinary white platinum, which is very compact and metallic
in all of its properties, we know platinum in the form of a spongy
mass which is known as platinum sponge. Platinum is obtained in
this condition when the chloride of ammonium and platinum, which
will be referred to a little later, is heated. Platinum sponge has
a gray color, and differs fundamentally from the black variety of
platinum known as platinum black. This is obtained by reduction
of platinum compounds. When both of these varieties are highly
heated and subjected to pressure, they pass back into ordinary white
platinum. Finely divided platinum has a remarkable power to
absorb oxygen. It can absorb several hundred volumes of oxygen,
and the oxygen in this condition is very active chemically, readily
effecting oxidations. In a similar manner it can absorb consider-
able quantities of hydrogen, and the hydrogen under these condi-
tions has strongly reducing properties.

This power of platinum to absorb gases, and to produce chemical
combinations between them, can be readily illustrated by means of
a piece of platinum foil and a gas-jet. Ignite the jet and heat the
platinum to redness. Then extinguish the jet until the platinum
ceases to glow. If now the gas-jet is turned on again and the
gas allowed to flow over the surface of the hot, but not incan-
descent platinum, it will be ignited again. This experiment
can be repeated as often as desired with the same piece of
platinum.

The power of platinum to absorb gases and produce chemical
reaction catalytically is also illustrated by a form of lamp devised by
Davy, An ordinary spirit lamp is filled with a mixture of alcohol
and ether and ignited. A spiral of platinum wire is suspended in
the flame and heated to incandescence. The flame is then extin-
guished uiitil the spiral ceases to glow. If the vapors of alcohol
and ether are now allowed to fall again on the hot platinum, the
latter will again become incandescent and ignite the lamp.

4d4 PRINCIPLES OF INORGAXIC CHEMISTRY

The Ddbereimer lamp, constmcted early in the nineteenth century,
is based upon the same principle. A current of hydrogen is allowed
to flow over spongy platinum, when it combines so rapidly with
oxygen that the platinum is heated to incandescence and ignites the
hydrogen-

Chemically, platinum is very resistant to reagents, and upon this
fact its value chiefly depends. It is not attacked appreciably by
any of the strong mineral acids, but dissolves in aqua regia. It,
however, dissolves in the fused alkalies, cyanides, etc. Platinum
must not be heated in contact with carbon, since some of the carbon
dissolves and makes the platinum brittle. Special precaution must
be taken not to heat a mixture of carbon and saud in platinum ves-
sels, since the platinum combines with the silicon, forming a brittle
substance. Similarly, phosphates must not be heated in platinum
vessels along with a reducing agent, since platinum readily com-
bines with phosphorus. Platinum readily forms alloys writh the
heavy metals, such as lead, mercury, etc. Such substances must
never be heated in platinum apparatus.

Uses of Platinum. — Platinum is especiaUy useful to the chemist
on account of its resistance to chemical reagents. It is not attacked
appreciably by the strongest acids, not even by hydrofluoric acid.
It is also not attacked by the fused alkaline carbonates. Platinum
vessels can thus be used where even porcelain could not be employed.
On account of its high melting-point platinum is used for holding
substances which are to be heated to a high temperature. Since it
does not volatilize in the flame of the Bunsen burner it does not
impart any color to the flame, and is therefore useful in spectrum
analysis to hold the substance whose spectrum it is desired to study.
A platinum wire is made into a loop and dipped intt) the substance
in question. It is then inserted into the flame, when only the
spectrum which is characteristic of the substance appears.

Platinum does not dissolve in acids on account of its low solution-
tension, and on account of this same property it does not pass into
solution when made the pole of an electric current. Platinum is,
therefore, used as electrodes in affecting electrolysis where it is
desired to separate the metals quantitatively and determine the
amounts which are present.

Platinum finds extensive use in almost every phase of gravi-
metric analysis. On account of its high fusion- and boiling-points
it has no appreciable vapor-tension, even at the temperature of the
blaHt-lami>. When it is desired to heat precipitates to a high tem-
perature platinum crucibles are frequently employed.

PLATINUM 495

Platinum finds to4ay extensive applications in the arts. Plati-
num has nearly the same temperature coefficient of expansion as
glass, and almost exactly the same as red and blue fusion glass. If
it is desired to seal an electrical connection through glass a platinum
wire is the most convenient means. This fact is utilized not only in
constructing standard cells, but also incandescent electric lights.
Large amounts of platinum are used in this way.

Platinum is also used in the form of fine wire to draw together
the tops of the mantles in the Welsbach burner. It is also employed
in constructing stills for concentrating sulphuric acid, and as a
catalyzing agent in the new method of making sulphuric acid. In
contact with platinized asbestos, sulphur dioxide combines with
oxygen at the elevated temperature forming sulphur trioxide.

Colloidal Solution of Platiniun. — A pseudosolution or colloidal
solution of platinum can readily be prepared by a method recently
discovered by Bredig. When two bars of platinum are dipped into
water, so that their lower ends are close together, but do not touch,
as shown in Fig. 43, and
an electric current of 8-12
amperes and 30-40 volts
passed through the circuit,
the platinum is torn off of
the bars in such a fine
state of division that the
solution appears perfectly
homogeneous, when ex-
amined under the most ^iq. 43.
powerful microscope. That

the platinum is not in a state of true solution, is shown by the
fact that it does not lower the freezing-point or vapor-tension of
the solvent. Such solutions, which have been prepared not only
of platinum, but of gold, silver, cadmium, iridium, etc., have been
studied extensively by Bredig, especially in the cases of platinum,
gold, and silver. These pseudosolutions of the heavy metals have
been found to have some remarkable properties in connection with
their catalytic action, especially as effecting the decomposition of
hydrogen dioxide. Analogies between the action of these solutions
and organic enzymes have been pointed out and established experi-
mentally, which are undoubtedly deep-seated. Without going into
detail in this subject, it may be said that infinitesimal quantities of
the same substances which poison the organic enzymes, also "poison"
the colloidal solutions of the metals, preventing or greatly hindering

496 PRINCIPLES OF INORGANIC CHEMISTRY

their catalytic activity. A number of other analogies have been
shown to exist.

Oxides and Hydroxides of Platinnm. — Platinum forms two oxides,
PtO and PtOj, which are derived from the corresponding hydroxides,
Pt(OH), and Pt(0H)4, by careful heating. The hydroxides are ob-
tained from the platinous and platinic chlorides, or from the double
chlorides with the alkalies, by the action of a base. All of these
substances are unstable at elevated temperatures, breaJkiug down
and yielding metallic platinum.

Chlorides of Platinum. — Platinum forms two chlorides — platinous

chloride, PtClj, in which the platinum is bivalent, Pt, and platinic

++++
chloride, in which the platinum is tetravalent, Pt.

Platinous chloride is obtained by heating platinum sponge to

250° in chlorine gas. It is a green powder which is insoluble in water.

When treated with hydrochloric acid it dissolves forming hydro-

chlorplatinous acid, HjPtCl^. When this compound is heated to 300*

it breaks down into hydrochloric acid and platinous chloride. This

acid is also produced by reducing hydrochlorplatinic acid, HjPtCl^

the latter losing two atoms of chlorine and passing over into the

former. Salts of hydrochlorplatinous acid are known. These have

the composition M2PtCl4. The chlorplatinous ion PtCl4 is, there-
fore, bivalent.

Platinic chloride, PtCl4, is obtained by carefully heating hydro-
chlorplatinic acid to 350° in the presence of chlorine. It is not
obtained by dissolving platinum in aqua regia, Platinic chloride,
which is quite soluble in water, readily combines with hydrochloric
acid forming hydrocMorplatinic acid, HgPtClg. This same com-
pound is formed when platinum is dissolved in aqua regia.

Many salts of this acid are known. We have already seen
that the potassium and ammonium chlorplatinates, K,PtCl« and
(NH4)2PtCl6, are difficultly soluble in water, while the sodium salt,
NajPtClfl, is readily soluble. This reagent thus enables us to sepa-
rate sodium from potassium. It is also very important in connection
with the quantitative determination of potassium, since potassium
chlorplatinate is almost completely insoluble in a mixture of water
and alcohol. The calcium and rubidium salts of this acid are even
less soluble than the potassium and ammonium compounds.

In this acid the chlorplatinic ion, PtClg, is bivalent, and this is
another example of a metal forming part of an anion. When a
solution of this acid is electrolyzed the platinum passes to the anode
and not to the cathode. We have met with a number of similar cases.

PLATINUM 497

Platinum in the bivalent condition forms a double nitrite with the
alkaline nitrites, which is analogous in composition to the alkali
platinous chlorides. Thus, when potassium platinous chloride is
allowed to remain in contact for a time with potassium nitrite, the
following reaction takes place : —

KjPtCl^ + 4 KNOa = KsPtCNOs)^ + 4 KCl.

The compound K2Pt(N02)4 is known as potassium platinonitrite.

Sulphides of Platiniim. — Platinum forms the two sulphides, PtS
and PtSj, as we would expect. These are obtained by precipitating
platinous and platinic solutions by means of hydrogen sulphide.
When heated these compounds lose sulphur, and metallic platinum
is formed.

Double Cyanides of Platinum. — Platinum forms a number of
double cyanides, which are characterized by their unusually beau-
tiful color and fluorescence. They can all be regarded as compounds
of platinous cyanide, Pt(CN)2, with the cyanide of the element in
question. The acid of which these substances are salts, H2Pt(CN)4,
can be obtained by treating the barium salt with sulphuric acid.
The potassium salt, K2Pt(CN)4.3H20 is formed by adding platinous
chloride to potassium cyanide.

Two compounds which are characterized by their unusual beauty
are the double cyanides of barium and magnesium. Bariimi plati-
nous cyanide, BaPt(CN)4.4H20, is formed by passing hydrocyanic
acid into water containing a mixture of barium carbonate and
platinous chloride. It is light yellow in color, with a beautiful play
of greenish-violet light over the surface as the crystals are turned.
This compound which converts ultra-violet radiation into visible
light, and is, therefore, fluorescent, has come into prominence
recently in connection with the study of Kdntgen rays. It trans-
forms these, and also the radiation given off by uranium, into visible
radiations, and has been used in the preparation of screens upon
which Rontgen rays are allowed to fall that their presence might be
detected by the eye.

The magnesium platinous cyanide, MgPt(CN)4, is remarkable for
its color. It can be formed by treating an aqueous solution of the
barium salt with a solution of magnesium sulphate. It crystallizes
in red prisms. It has a bright green lustre, so that when the sides
are observed in reflected light they appear green, while the ends
are blue.

It is questionable whether another compound of equal beauty is
known in the whole field of chemistry.
2k

INDEX

AbM Nollet, demonstration of osmotic
pressure, 100.

Absolute boiling-point of a gas, 284.

Absolute zero, can it be realized experi-
mentally ?, 44.

Absolute zero of temperature, determi-
nation of, 26.

Absorption of certain constituents from
the soil by plants, 303.

Accumulator, 483.

Accumulators, 193.

Acetylene, effect on luminosity of
flames, 203.

Acetylene hydrocarbons, 278.

Acetylene light, 296.

Acid and basic properties, 144.

Acid, definition of, 124.

Acid dibasic, 212.

Acidic indicator, 213.

Acid means hydrogen ions, 124.

Acid, monobasic, 212.

Acid salts, 216.

Acids, halogen, comparison of the, 169.

Acids, hydrogen present in all, 40.

Acids, neutralization of by bases, 210,
217.

Acid sodium carbonates, 327.

Acid sodium sulphate, 325.

Acids of silicon, 300.

Acids, organic, 288.

Acid sulphides, .181.

Acids, weak, neutralization with weak
bases, 219.

Acid, tribasic, 212.

Agate, 300.

Air, 235.

Air, a mixture or compound, 237.

Air, liquid, 238.

Air, physical properties of, 237.

Ahr-slaked lime, 365.

Alba, magnesia, 386.

Albite, 414.

Alcohols, 284.

Aldehydes, 287.

Alkali metals, characteristics of, 361.

Alkaline earths, detection of, 381.

Alkaline earths, relations between, 38L

Allotropic forms of carbon, 272.

AUotropic modification of oxygen, 29.

Allotropy, 29.

Alloys, 314, 476.

Alloys of aluminium, 408.

Alloys of manganese, 436.

Alloys of silver, 469.

Alum, burnt, 413.

Aluminate of sodium, 410.

Aluminates, 410.

Aluminite, 413.

Aluminium, alloys of, 408.

Aluminium amalgam, 409.

Aluminium bronze, 408, 462.

Aluminium carbide and carbonate, 414.

Aluminium chloride, 411.

Aluminium, detection of, 416.

Aluminium fluoride, cryolite, 412.

Aluminium, occurrence and prepara-
tion, 407.

Aluminium oxide and hydroxide, 409.

Aluminium, properties of, 408.

Aluminium silicates, 414.

Aluminium silicates, applications of,
415.

Aluminium sulphate, 412.

Aluminium sulphide, 412.

Alums, 333, 413.

Alums, chromium, 449.

Amalgamation process, 467, 475.

Amalgams, 207, 388, 399, 409, 485.

Amblygonite, 352.

Amethyst, 300.

499

500

INDEX

Amethyst, oriental, 400.

Ammonia, 202.

Ammonia, action of on cobalt salts, 433.

Ammonia, action of on mercury salts,
405.

Ammonia, chemical properties of, 203.

Ammonia, composition of, 204.

Ammoniac, sal, 203.

Ammonia liquid, high specific heat of,
206.

Ammonia, or Solvay process of pre-
paring ammonium carbonate, 327.

Ammonia, physical properties of, 205.

Ammonium, 204, 207.

Ammonium acid carbonate, 360.

Ammonium amalgam, 207.

Ammonium carbonate, 360.

Ammonium chloride, 356.

Ammonium chloride dissociated by
heat, 367.

Ammonium chloride, dissociation by
heat diminished by an excess of
either product of dissociation, 357.

Ammonium chloride, dry, dissociation
of by heat, 367.

Ammonium hydrazoate, or triazoate,
358.

Ammonium hydrosulphide, 369.

Ammonium hydroxide, 221, 365.

Ammonium hydroxide, dissociation of,
206,222.

Ammonium hydroxide, measurement
of the dissociation of, 222.

Ammonium ion, characteristic reaction

. of, 362.

Ammonium, salts of, 368-360.

Ammonium sodium phosphate, 329.

Amorphous forms of carbon, 273.

Amphoteric reaction, 328.

Analogues, atomic, 148.

Andrews, discovered critical tempera-
ture and pressure, 27.

Andrews, on critical constants, 284.

Anglesite, 482.

Anhydrite, 129, 368.

Anion, 64, 111.

Annealed glass, 375.

Anode, 50.

"Antichlor," 194,326.

Antimonic acid, 264.

Antimonic acid, meta-, 204.

Antimonic acid, pyro-, 264.

Antimonious acid, sulph-, 266.

Antimonite, metasulph-, 266.

Antimony, acids of, 264.

Antimony butter, 266.

Antimony, compounds with hydrogen
and oxygen, 262.

Antimony, compounds with the halo-
gens, 264, 265.

Antimonyl group, 263.

Antimony, occurrence and prepazatioD,
261.

Antimony, oxides of, 262.

Antimony, properties of, 261.

Antimony, sulphides of, 265.

Apatite, 242, 372.

Aqua regia, 232.

Aragonite, 369.

Arbor Saturnl, 479.

Argentan, 462.

Argentite, 467.

Argon, 239.

Argon, number of atoD:i8 in the mole-
cule of, 239,

Argyrodite, 306.

Arrhenius, work on the origin of the
theory of electrolytic dissociation,
109.

Arsenate, sulphometa-, 260.

Arsenic acids, 268.

Arsenic, compounds with halogens, 258.

Arsenic, compounds with hydrogen
and oxygen, 256.

Arsenic, compounds with sulphur, 259.

Arsenic, occurrence and preparation,
265.

Arsenic, properties of, 255.

Arsenic pyrites, 255. ,

Arsenic acid, sulph- or thlo-, 200.

Arsenic, sulpho-salts of, 259.

Arsenic, white, 257.

Arsenious acid, 267.

Arsenious acid, sulph- or thio-, 260.

Arsine, 256.

Artificial preparation of ice, 206.

Asbestos, 384.

Atmosphere, composition of the, 235.

INDEX

601

Atmospheric air, 235.

Atmospheric air and rare elements oc-
curring in it, 235.

Atomic analogues, 148.

Atomic theory, 12.

Atomic volumes, 14.5.

Atomic volumes, curve of, 146.

Atomic weights and chemical proper-
ties, 142.

Atomic weights and combining num-
bers, 69.

Atomic weights and physical proper-
ties, 145.

Atomic weights corrected by the Peri-
odic System, 147.

Atomic weights, determination of, 60.

Atomic weights from molecular weights,

• 73.

Atomic weights from specific heats, 75.

Atomic weights, isomorphism an aid in
determining, 77.

Atomic weights, most accurate method
of determining, 79.

Atomic weights of the elements, table
of, 81.

Aurates, 477.

Available phosphoric acid, 374.

Avogadro's hypothesis, 71.

Avogadro^s hypothesis and molecular
weights, 72.

Avogadro's law, apparent exceptions
to, 87.

Avogadro^s law applied to the osmotic
pressure of solutions, 106.

Avogadro's law, explanation of appar-
ent exceptions to, 88.

Azurite, 460, 465.

Baeyer prepares permonosulphoric acid,
195.

Banca tin, 484.

Barite, 378.

Barium chromate, 449.

Barium, detection of, 381.

Barium dioxide, 378.

Barium dioxide used in preparing oxy-
gen, 16.

Barium hydroxide, 379.

Barium, occurrence, 378.

Barium, other insoluble compounds ol.
381.

Barium, oxides of, 378.

Barium platinocyanide, 497.

Barium, salts of, 379-381.

Barium sulphate transformed into ba-
rium carbonate, 380.

Baryta water, 879.

Base, diacid, 211.

Base-forming elements, 361.

Base, monacid, 211.

Bases are hydroxy I compounds, 210.

Bases, neutralization of by acids, 210,
217.

Bases weak, neutralization with weak
acids, 219.

Base, triacid, 211.

Basic and acid properties, 144.

Basic indicators, 215.

Basic lining in Thomas-Gilchrist con-
verter, 422.

Basic salts, 216.

Batteries, use of zinc in, 391.

Bauxite, 305, 407, 409.

Beckmann, boiling-point apparatus, 97.

Beckmann, freezing-point apparatus,
94.

Beckmann thermometer, 96.

Becquerel rays, 457.

Bell metal, 462.

Benzene hydrocarbons, 278.

Berlin blue, 427.

Berthollet, on law of constant propor-
tion, 9.

Beryl, 383.

Bessemer converter, 422.

Bessemer steel, 422.

Binary electrolyte, 112.

Bismuth chloride, 268.

Bismuth hydroxide, 268.

Bismuth, occurrence and properties,
267.

Bismuth, oxides of, 268.

Bismuth sulphide, 268.

Bismuthyl group, 268.

Bivalent, 134.

Blanc, I^, method of preparing sodium
carbonate, 326.

Bleaching-powder, 866.

602

INDEX

Bleaching-powder used in making chlo-
rine, 116.

Block-tin, 484.

Blowing of glass, 376.

Blowpipe, 293.

Blowpipe, oxyhydrogen, 87.

Blue vitriol, 464.

Bog-iron ore, 410.

Boiling-point, absolute, of a gas, 284.

Boiling-point apparatus of Beckmann,
06.

Boiling-point apparatus of Jones, 08.

Boiling-point method of determining
molecular weights, 05.

Bone-ash, 372.

Bone-black, 274.

Boracite, 307.

Borax, 307, 330.

Borax bead, 330.

Borax octohedral, 330.

Borax, prismatic, 330.

Boric acids, 308.

Borocalcite, 307.

Boron, 307.

Boron, compounds with the halogens,
309.

Boron nitride, 308.

Boron, occurrence, preparation and
properties, 307.

Boron, specific heat varies with the
temperature, 307.

Boron trioxide, 308. _

Boyle and Gay-Lussac, combined ex-
pression of the laws of, 26.

Boyle's law for gases, 24.

Boyle's law for osmotic pressure, 105.

Brass, 388, 462.

Brauner's Periodic System, 141.

Braunite, 436.

Brick or fire-brick, 415.

Brimstone, crude, 171.

Britannia metal, 485.

Bromic acid, 157.

Bromine, 152.

Bromine atoms and bromine ions, 154.

Bromine, chemical properties of, 153.

Bromine, compound of with chlorine,
158.

Bromine, compound with iodine, 165.

Bromine, compounds with oxygen and
hydrogen, 157.

Bromine, detection of, 153.

Bromine, occurrence and preparation,
152.

Bromine, physical properties of, 154.

Bromine water, 164.

Bromoform, silicon, 303.

Bronzes, 408, 436, 462, 485.

Bunsen and Kirchhofif, discovery of
rubidium, 354.

Bunsen and Kirchhoft^s law of light
absorption, 45.

Bunsen burner, 293.

Bunsen, method of determining densi-
ties of gases, 87.

Bunsen photometer, 297.

Burning in oxygen, explanation of, 18.

Burnt alum, 413.

Caesium, occurrence, compounds, 355.

Cadmium, 396.

Cadmium molecule consists of one

atom, 397.
Cadmium, salts of, 307.
Cailletet, liquefaction of oxygen, 27.
Calamine, 387.
Calcite, 368.

Calcium acid carbonate, 371.
Calcium carbide, 368.
Calcium carbonate, 368.
Calcium ca,rbouate, decomposition

curves, 370.
Calcium carbonate, primary, 391.
Calcium, compounds with the halogens,

305.
Calcium, detection of, 876.
Calcium hydride, 363.
Calcium hydroxide, 364.
Calcium hypochlorite, 129, 366.
Calcium manganite, 438.
Calcium, occurrence, preparation, and

properties, 353.
Calcium oxalate, 375.
Calcium oxide, 363.
Calcium phosphates, 372.
Calcium silicate, 374.
Calcium sulphate, 367.
Calcium sulphides, 307.

INDEX

503

Calomel, 401.

Calorie, 22.

Calorimelerf 22.

Candle, 291.

Carats, 476.

Carbamate of ammoniam, 860.

Carbide of calcium, 368.

Carbides, 2^.

Carbon, allotropic forms of, 272.

Carbon, amorphous forms of, 273.

Carbonates, formed from silicates, 302.

Carbonates, hydrolyzed by water, 282,

328.
Carbon bums in oxygen, 17.
Carbon, compounds with halogens, 288.
Carbon, compounds with hydrogen, 277.
Carbon, compound with silicon, — car-
borundum, 304.
Carbon, different forms contain dif-
ferent amounts of energy, 275.
Carbon dioxide, 2HO-283.
Carbon dbxHe reriuced by plants, 282.
Carbon dlsiulphMe, 289.
Carbonic acid, thio-, 280.
Carbonic acid, trithio-, 280.
Carbon monoxide, 278.
Carbon monoxide, thermochemistry of,

279.
Carbon, r61e of, in producing light, 291.
Carbon, specific heat of, 276.
Carbon subpjcide, 287.
Carbon tetrnchlortde, 288.
Carbonyl compounds, 429.
Carbonyl, nickel, 434.
Carborundum v?04,
Camallite, 1 10, 333, 337, 354, 384.
Carols liquid 105.
Carr^ ice machines^ 206.
Casfiiterite, 484.
CflflsiiK^ purple of, 476.
Cast-iron, gray, white, 421.
Catalytic decomposition of hydrogen

dioxide, 67.
Catalytic reactions and catalyzers, 36,

187.
Cathode, 50.
Cation, 64, 111.
Caustic, lunar, 473.
Caustic potash, 336.

Cavendish discovered hydrogen, 33.

Celestite, 376, 377.

Cells, storage, 193.

Cement, 416.

Cement, hydraulic, 370.

Cement, Portland, 370.

Cements, hydraulic, 410.

Cerite, 417.

Cerium, compounds of, 300.

Cenjssite, 478, 482.

Cbal copy rile, 460.

Chalcosite, 460.

Chalk, 363, 369.

Chamber acid, 189.

Chamber crystals, 234.

Chamber, leaden, 188.

Chameleon, mineral, 441.

Charcoal, 273.

Chemical action at a distance, 425.

Chemical combination, 6.

Chemical elements, 4.

Chemical equation, 6.

Chemical properties and atomic weights,
142.

Chemical reactions reversible, 179.

Chemistry and Physics, relations be-
tween, 1.

Chemistry, science of, 7.

Chili saltpetre, 158, 312.

China silver, 462.

Chloramide, mercurous, 406.

Chlorates, 131.

Chlorates, ion of chlorine, and the ion
of, 131.

Chlorctiromic acid, 451.

Chlor-ekctrolytie gas, 121.

Chloric acid, 130.

Chlori nation process, 475.

Chlorine, 115.

Chlorine a bleaching agent, 118.

Chlorine, action on hydrogen, 118.

Chlorine, action on organic compounds,
118.

Chlorine, action on water, 118.

Chlorine a diisinfectant, 118.

Chlorine and hydrogen, volume rela-
tions in which they combine, 121.

Chlorine, an element or compound, 115.

Chlorine, a strong oxidizing agent, 118.

504

INDEX

Chlorine, chemical properties of, 117.

Chlorine, combuBtion in, 117.

Chlorine, coinpooud of with bromine,
158.

Chlorine, compounds with iodine, 105.

Chlorine, compounds with oxygen and
hydrogen, 128.

Chlorine, compounds with sulphur, 105.

Chlorine litiipniitm;^ ^ii», 121.

Chlorine dioxide^ 33.

Chlorine, dry inactJvtiy of, 121.

Chlorine, electrolytic method of pre-
paring, 117.

Chlorine hydrate, 110.

Chlorine ion and the ion of chlorates,
131.

Chlorine, liquefaction of, 120.

Chlorine, occurrence and preparation,
115.

Chlorine, power of to combine with
oxygen, 128.

Chlorine, prepared by methods of
Deacon and VVeldon, 116.

Chlorine wattr 118.

Chloroform, ailipoiij 303.

Chlorous acid, 133.

Cblort%alphuric acid, 196.

Chromateii, Aiih

Chrome orange, 483.

Chrome yellow, 450, 483.

Chromic acid, 44(5, 449.

Chromic acid, chlor-, 451.

Chromic acid, chlorides of, 451.

Chromic acid, Morse's method of pre-
paring, 449.

Chromic acid, per-, 451.

Chromic chloride, 448.

Chromic salts, 447.

Chromites, 445, 449.

Chromium, 445.

Chromium alums, 449.

Chromium, detection of, 452.

Chromium, hydroxides of, 446.

Chromium ions, valence and properties
of, 446.

Chromium, oxides of, 445.

Chromium, pt?riodical or rhythmic solu-
tion in acids, 445.

Chromous salts, 447.

Chiomyl chloride, 451.

Chrjsoberjl, .^-^ :. 410.

Cmnabar, 7], 308, 404.

Citmte soluble pho6{^iate, 373.

Clark element, 406.

Claosius's electro-chemical Uieory, llOl

aay, 407, 414.

Coal-gas, 202.

Coal or stone-coal, 276.

Cobalt, 431.

Cobalt bead, 432.

Cobalt cyanides, 432.

Cobalt, double nitrite of, 433-

Cobaltic compounds, 431.

Cobaltite, 431.

CobiikitetJ, 432.

Cobaltous compounds, 431.

Cobaltous salts, 432.

Cobalt, oxides and hydroxides, 431.

Cobalt salts, action of ammonia on,
433.

Coins, 462.

Coke, 274.

Colloidal silver, 468.

Colloidal solution, 424.

Colloidal solution of gold, 476.

Colloidal solution of platinum, 496.

Colloidal solution of tungstic acid, 466.

Colloids, 301.

Color, change in, with change in elec-
trical charge, 420.

Color of pentianganates, 443.

Columbium, 270.

Combination, chemical, 6.

Combining numbers and atomic
weights, 69.

Combining numberi, chemical methods
of dutenjiining» 00.

Combining weights, law of, 11.

Combustionn 18.

CombnsrJnn in diloHne, 117.

Combustion, increase of weight in, 19.

Combustion, measurement of the beat
of, 21.

Combustion, oxygen used up in, 20.

Combustion, phlogiston theory of, 18.

Commercial fertilizer, 373.

Commercial phosphate, analysis of,
873.

INDEX

605

Compoflition changed in chemical re-
action, 2.

Composition of the earth, 6.

Compoeition of water, 60.

Compounds and elements, 8.

Compounds with oxygen, names of, 23.

Concentration element, 304.

Condensation of steam, heat of, and of
solidification of water, 56.

Conductivities, Kohlrausch method of
measuring, 112.

Conductivities, molecular, 112.

Conductivity measurements, dissocia-
tion calculated from, 114.

Conductivity method, 112.

Conductivity of water, 63.

Conservation and correlation of energy,
13.

Conservation of energy, importance of
for the science of chemistry, 13.

Conservation of mass, law of, 8.

Constant proportion, law of, 0.

Continuity of passage from liquid to
gas, 285.

Cooling, effect of on flame, 295.

Copper acetylene, 465.

Copper, alloys of, 462.

Copper ferrocyanide, 428, 465.

Copper, occurrence and preparation,
460.

Copper, oxides of, 462.

Copper purified by electrolysis, 461,
465.

Copper pyrites, 171, 460.

Copper, salts of, 403-465.

Correlation and conservation of energy,
13.

Corrosive sublimate, 401.

Corundum, 407, 409.

Crafts-Friedel reaction, 411.

Critical density, 284.

Critical temperature and pressure, 60,
283.

Critical volume, 284.

Crocoisite, 445, 478.

Crookes, discoverer of thallium, 417.

Cry of tin, 484.

Cryolite, 168, 313, 407, 412.

Crystallization, fractional, 157.

Crystallization or freezing, 50.

Crystallization, water of, 46.

Crystalloids, 301.

CupellatiQn, 468.

Cupric compounds, 462.

Cuprite, 460.

Cuprous compounds, 462.

Curie, M. and Mme., on XBdioactive

substances, 458.
Cyanic acid, 290.
Cyanic acid, sulpho-, 290.
Cyanide process, 475.
Cyanine, 215.
Cyanogen, 289.

Daguerre, 472.

Daguerreotype, 472.

Dal ton, on the atomic theory, 12.

Daniell cell, 395.

Davy prepares potassium, 388.

Davy prepares sodium, 313.

Davy safety lamp, 295, 493.

Deacon's process for making chlorine,
116.

Decomposition, double, 328.

Decomposition, heat of, and of forma-
tion, 22.

Decrepitation of sodium chloride, 317.

Deliquescent, 322.

Densities and molecular weights of
gases, 82.

Densities of gases, determination of
molecular weights from, 70.

Density, critical, 284.

Detonating gas, 85.

Detonating gas, chlorine and hydrogen,
121.

" Developer," 471.

Dewar, liquefaction of fluorine, 167.

Dewar, liquefaction of oxygen, 28.

Dewar liquefied and froze hydrogen, 48.

Diacid base, 211.

Diagram of sulphur, temperature-press-
ure, 174.

Diagram of water, temperature-pressure,
58.

Dialysis, 301, 424.

Dialyzer, 301.

Diamond, 272.

506

INDEX

m\.

Diamorphism, 172, 900.

Diaspare, 409.

Dibasic acid, 212.

Dichromates, 450.

Dielectric constant of water, 61.

Diffusion of gases, law of, 4f.

Disinfectant, 66.

Disinfectant, chlorine, 118.

Dissociating power of water, 63.

Dissociation by heat, 162.

Dissociation by heat, of ammonium

chloride, 357.
Dissociation calcolated from condnc-

tiYity measurements, 114.
Dissociation electrolytic, 64.
Dissociation electrolytic, measurement

of. 111.
Dissociation, hydndytic, 249.
Dissociation, natore of conditioned by

the presence of another sobstance,

389.
Dissociation of ammonium hydroxide,

222.
Dissociation of a weak base^ measure-
ment of, 222.
Dissociation of copper sulphate in its

water of crystallization, 465.
Dissociation of diy ammonium chloride,

357.
Dissociation of hydrogen sulphide, 181.
Dissociation of nitrates, 232.
Dissociation of nitric acid, 232.
Dissociation of sulphuric acid, 191.
Dissociation of vapors diminished by

an excess of one of the products of

dissociation. 91.
Dissociation, origin of the theory of

electrolytic, 109.
Dissolved substances affect the proper-
ties of water, 63.
Dissolved substances, determination of

the molecular weights of, 82.
Dissolved substances, molecular weights

of, determined by the boiling-point

method, 95.
Dissolved substances, molecular weights

of, determined by the freezing-point

method, 93.
Distance, chemical action at a, 425.

Disthene, 414.

Distillation, 49.

Distillation fractional, dS.

Disulphuric acid, 104.

Ddbereiner lamp, 494.

D5bereiner triads, 130.

Dog*s grotto, 280.

Dolomite, 381, 386.

Double decomposition, 323.

Drummond light, 364.

Dry ammonia does not act oi

hydrochloric acid, 203.
Dry chlorine, inactivity of, 121.
Dry hydrochloric acid does not a

dry ammonia, 203.
Dry hydrogen will not oombiiie

dry oxygen, 38.
Dulong and Petit's law, 75.
Dulong and Petit, specific hemt o

bon, 276.
Dumas, method of, 82.
Dynamic condition of eqoilibriain

Earth, composition of, 6.
Earthenware or stoneware, 415w
Effervescent water, 48.
Efflorescent, 324, 326.
Ekaaluminhim, 149, 419.
Ekaboion, 149, 416.
Ekasilicon, 149.
Electric furnace, 273.
Electricity, amount of, 361.
Electricity, potential of, 361.
Electric light, 296.
Electro-chemical theories of da

and Williamson, 110.
Electrolysis, 60.
Electrolysis of hydrogen, 45.
Electrolysis of potassium hydro;

theory of, 335.
Electrolysis of water, 50.
Electrc^ytes, 63.
Electrolytic dissociation, 64.
Electrolytic dissociation, me&sorei

of. 111.
Electrolytic dissociation, origin of

theory of, 109.
Electrol^-tic dissociation, theory

110.

INDEX

507

Electrolytic dissociation, theory of and
osmotic pressure, 100.

Electrolytic gas, 35.

Electrolytic gas, chlorine and hydro-
gen, 121.

Electrolytic method of preparing chlo-
rine, 117.

Electrolytic process, 468.

Electrolytic purification of copper, 466.

Electro-plating, 4S4.

Elements and compounds, 3.

Elements and compounds, number of, 3.

Elements, chemical, 4.

Elements predicted by means of the
Periodic System, 147.

Enantiotropic, 172.

Enantiotropic mercuric iodide, 403.

Endothermic reaction, 225.

Energy, different amounts in oxygen
and ozone, 31.

Epsom salt, 385.

Equation, chemical, 6.

Equilibrium, 180.

Equilibrium, dynamic condition of, 180.

Equilibrium metastable, 60.

Equivalent normal solution, 211.

Esters, 288.

Ether, 288.

Ethylene hydrocarbons, 277.

Eudiometric method, 236.

Eutectic alloy, 267.

Euxenite, 416, 417.

Faraday, liquefaction of chlorine from

chlorine hydrate, 110.
Faraday's law the basis of chemical

valency, 134.
Feldspars, 407, 414.
Ferrates, 430.
Ferric acid, 430.
Ferric compoimds, 423-430.
Ferrocyanic acid, hydro-, 427, 428.
Ferrocyanide of copper, 428.
Ferrocyanide of potassium, 427, 428.
Ferrous compounds, 423-426.
Filtration, 48.

'* Fixing" agent in photography, 325.
Flame, effect of cooling, 205.
Flame, oxidizing, 204.

Flame, reducing, 203.

Flames and their luminosity, 202.

Flint, 300.

Flowers of sulphur, 171.

Fluorescence, 365.

Fluorine, 165.

Fluorine, compound with iodine, 160.

Fluorine, liquefaction of, 167.

Fluorine, occurrence and preparation,
165.

Fluor spar, 165, 363.

Flux, 330.

Formation, heat of, and of decomposi-
tion, 22.

Fractional crystallization, 157.

Fractional distillation, 65.

Franklinite, 387.

Franklin, work on liquid ammonia, 206.

Freezing of water, 55.

Freezing, or crystallization, 60.

Freezing-point apparatus of Beckmann,
94.

Freezing-point method of determining
the naolecular weights of dissolved
substances, 03.

Freezing-point of water, 63.

Friedel-Craft^s reaction, 411.

Fulminating mercury, 405.

Fuming nitric acid, 232.

Furnace, electric, 273.

Fusible white precipitate, 406.

Fusion, heat of, of ice, 56.

Gadolinite, 306, 416, 417.

Gahnite, 410.

Galena, 171, 478.

GaUium, 416.

Galvanized zinc, 388.

Garnets, 374, 415.

Gamierite, 433.

Gases, determination of the molecular

weights of, 82.
Gases, law of Boyle for, 24.
Gases, law of Gay-Lussac for, 26.
Gases, permanent, 43.
Gas law as applied to osmotic pressure,

exceptions, 108.
Gas-pressure and osmotic pressure,

causes of, 107.

608

INDEX

Gas-pressure and osmotic pressure,
equality of, 107.

Gas-pressure and osmotic pressure, re-
lations between, 104.

Gas to liquid, continuity of passage
from, 285.

Gay-Lussac and Boyle, combined ex-
pression of the laws of, 26.

Gay-Lussac, law of for gases, 26.

Gay-Lussac, method of for determin-
ing molecular weights, 84.

Gay-Lussac, numbers of atoms in equal
volumes of different gases, 71.

Gay-Lussac showed elementary nature
of chlorine, 115.

Gay-Lussac 's law for osmotic pressure,
106.

Gay-Lussac tower, 188.

Generalizations, 7.

Germanium, 306.

Germanium, compounds of, 805.

German silver, 388, 434, 462.

Gersdorffite, 433.

Gibbs, Willard, phase rule of, 69.

Gilchrist-Thomas, converter, 422.

Glacial phosphoric acid, 261.

Glasses, colored, 374.

Glass, different varieties of, 374.

Glauber's salt, 323.

Glazing of porcelain, 415.

Glover tower, 188.

Glucinum, .383.

Gluoinum, compounds of, 383.

Gneisses, 298.

Gold, 474.

Gold, colloidal solution of, 476.

Gold, fooPs, 426.

Gold, metallurgy of, 474.

Gold, oxides and hydroxides of, 477.

Gold plating, 476.

Gold, properties of, 476.

Gold, salts of, 477.

Goldschmidt method of preparing
chromium, 446.

Goldschmidt method of preparing tung-
sten, 455.

Gold telluride, 474.

Granites, 298.

Graphite, 272.

Greenockite, 897.

Grotto, dog's, 280.

Groups, relations within the, 144

Guldberg and Waage's law of

action, 92.
Gun-metal, 462.
Gunpowder, 343.
Gypsum, 363, 367.

Hall method of preparing alumj

408.
Halogen acids, comparison of the
Halogens, compounda with tx>ron
Hard water, 47.
Hard waters, 871.
Hausmannite, 436.
Heat, dissociation by, 162.
Heat energy produced when o:

and hydrogen combine, 37.
Heat of combustion, measoreme]

21.
Heat of condensation of steam, a

solidification of water, 56.
Heat of formation and of decoD

tion, 22.
Heat of fusion of ice, 56.
Heat of neutralization, 217.
Heat of neutralization, explanati

constant, 218.
Heat of vaporiz?ition of water, 54.
Heavy spar, 378.
Heinrich Rose points out that

cates are being converted inU

bonates, 802.
Helium, 240.
Hematite, 419, 423.
Henry's law, 63.
Hess, law of the thermoneutrali

salt solutions, 220.
Hofmann, modification of Gay-Lua

method, 84.
Homologous series of compounds.
Homologous series of silicic acids.
Hornblende, 384, 387.
Horn -silver, 467.
HUbnerite, 465.
Hydrargillite, 409.
Hydration, water of, 46.
Hydraulic cement, 370, 410.

INDEX

609

Hydraulic minin'g, 474.

Hydrazine, 202.

Hydrazine, properties of, 208.

Hydrazoic acid, method of formation,
208.

Hydride of lithium, 353.

Hydride of sodium, 315.

Hydride of tellurium, 199,

Hydrides, 40.

Hydnodic acid, 161.

Hydrobromic acid, 155.

Hydrochlorauric acid, 477.

Hydrochloric acid, 121.

Hydrochlorpalladic acid, 401.

Hydrochlorplatinic acid, 496.

Hydrochlorplatinous acid, 496.

Hydrochlorplumbic acid, 481.

Hydrochlorstannous acid, 487.

Hydrocyanic acid, 290.

Hydroferricyanic acid, 428.

Hydroferrocyanic acid, 427.

Hydrofluoric acid, 168.

Hydrofluosilicic acid, 804.

Hydrofluozirconic acid, 306,

Hydrogen, 33.

Hydrogen, action on chlorine, 118.

Hydrogen and chlorine, volume rela-
tions in which they combine, 121.

Hydrogen and hydroxyl ions combine
when in the presence of one an-
other, 215.

Hydrogen and nitrogen, compounds of,
209.

Hydrogen and nitrogen, volume rela-
tions in which they combine, 202.

Hydrogen, combination with oxygen,
35.

Hydrogen, compounds of with metals,
89.

Hydrogen, compounds with sulphur,
177.

Hydrogen diffuses rapidly, 41.

Hydrogen dioxide, 64.

Hydrogen dioxide a good oxidizing
agent, 66.

Hydrogen dioxide also a reducing
agent, 66.

Hydrogen dioxide and water, relations
between, 67.

Hydrogen dioxide, catalytic decompo-

. sition of, 67.
Hydrogen dioxide, preparation and

purification, 64.
Hydrogen dioxide used in preparing

oxygen, 16.
Hydrogen, discovered by Cavendish,

33.
Hydrogen, dry, will not combine with

dry oxygen, 38.
Hydrogen, electrolysis of, 45.
Hydrogen, liquefaction of, 42.
Hydrogen liquid, Dewar's experiment

with, 44.
Hydrogen nascent, 40.
Hydrogen persulphides, 183.
Hydrogen present in all acids, 40.
Hydrogen, reducing power of, 38.
Hydrogen selenide, 198. •
Hydrogen silicide, 299.
Hydrogen, solidification of, 43.
Hydrogen spectrum, 44.
Hydrogen sulphide, 177.
Hydrogen sulphide a dibasic acid, 181.
Hydrogen sulphide, dissociation of, 181.
Hydrolysis, 213,

Hydrolysis of carbonates, 282, 328.
Hydrolytic dissociation, 240, 323.
Hydroscopic substancCi 223.
Hydrosulphide of calcium, 367.
Hydrosulphides, 181.
Hydroxylamine, 223.
Hydroxyl and hydrogen Ions combine

when in the presence of one another,

215.
Hydroxyl compounds include all 1

210.
*' Hypo," 194, 325, 471.
Hypobromous acid, 157.
Hypochlorite of calcium, 366.
Hypochlorous acid, 128,
Hyponitrous acid, 228.
Hypophosphoric acid, 251.
Hypophosphorous acid, 252.
»* Hyposulphite," 194.

Ice, artificial, preparation of, 206.
Ice, heat of fusion of, 56.
Iceland spar, 368.

610

INDEX

Ice machiues, Carr^f 206.

Illumination, 201.

Imperfections in the Periodic System,

161.
Indicators, 212.
Indicators, acidic, 213.
Indicators, basic, 215.
Indicators, theory of, 212.
Indium, 417.
Inductive capacity, specific, of water,

61.
Infusible white precipitate, 406.
Ink, sympathetic, 482.
Intensities of light, measurement of

relative, 297.
Iodic acid, 163.
Iodide of nitrogen, 233.
Iodine, 158.
Iodine, compounds with oxygen and

hydrogen, 163.
Iodine, compound with bromine, 165.
Iodine, compound with chlorine, 165.
Iodine, compound with fluorine, 169.
Iodine, detection of, 159.
Iodine, detection of in the presence of

bromine and chlorine, 160.
Iodine, molecular weight in the form

of vapor, 161.
Iodine, occuiTence and preparation, 158.
Iodine pentafluoride, 169.
Iodine, tincture of, 161.
Ion, chlorine, and the ion of chlorates,

131.
Ion formatiori, modes of. 111, 425, 466,

476.
Ions, 64, 110.

Ions, chromic and dichromic, 461.
Ions the active agents chemically, 429.
Iridium, 492.
Iridium compounds, 492.
Iridium-osmium, 492.
Iridium-platinum, 492.
Iron bums in oxygen, 17.
Iron chromite, 446.
Iron, impure or commercial, 421.
Iron, occurrence and preparation, 419.
Iron oxides, 423.
Iron pentacarbonyl, 279.
Iron, rusting of, 420.

Iron sesquisulphide, 426.

Iron, Spiegel, 421, 437.

Iron, sulphides of, 426.

Iron vitriol, 426.

Iron, welding of, 420.

Iron, wrought, 421.

Isomeric substances, 227.

Isomorphism among the alums, 4

Isomorphism, an aid in detem

atomic weights, 77.
Isomorphous substances, 267.
Itacolumite, 272.

Jasper, 300.

Jones's boiling-point apparatus, 9

Kainite, 345, 384.

Kaolin, 407, 414.

Kelvin, on the size of atoms, 12.

Ketones, 288.

Kieserite, 384.

Kindling temperature, 295.

Kinetic theory of liquids, 286.

Kipp apparatus, 86, 116.

KirchhofT and Bunsen, discoYe]

rubidium, 354.
Kux:hhofr and Bunsen's law of

absorption, 45.
Kohlrausch, law of, 222.
Kohlrausch, method of measuring

ductivities, 112.
Kopp, molecular heats the sum o

atomic heats, 76.
Krypton, 241.
Kundt^s method of determining

velocity of sound in a gas, 240.

Lamp-black, 274.

Lamp, safety, 295.

Landolt, on the law of the conservj

of mass, 9.
Langbeinite, 385.
Lanthanum, 417.
Lapis lazuli, 416.
Laughing gas, 224.
Lavoisier explains rOle of oxyge;

combustion, 19.
Lead, 478.
Leaden chamber, 188.

INDEX

611

Lead, hard, 266.

Lead, hydroxides of, 480.

Lead, occurrence, preparation, and

properties, 478.
Lead, oxides of, 470.
Lead precipitated by metals, 479.
Lead salts, 481-483.
Lead, sugar of, 483.
Lead tree, 479.
Lead, vinegar of, 483.
Lead vitriol, 482.
Le Blanc, method of preparing sodium

carbonate, 326.
Leclanch^ cell, 439.
Lecoq de Boisbaudran discovered eka-

aluminiuro — gallium, 149,416.
Lepidolite, 352, 354.
Leucite, 354.
Light, electric, 296.
Light, intensities of, measurement of

relative, 297.
Light, Welsbach, 296.
Lime, 363.

Lime, air-slaked, 365.
Lime, soda, 365.
Limestone, 363, 369.
Lime water, 364.

Linde, liquefaction of oxygen, 27.
Liquids, kinetic theory of, 286.
Liquid to gas, continuity of passage

from, 285.
Litharge, 479.

Lithium, compounds of, 352-353.
Lithium, discovery, preparation, and

properties, 352.
Lithium hydride, 353.
Lithium ion, characteristic reaction of,

362.
Litmus, 214.

Lothar Meyer's Periodic System, 137.
Low temperatures, measurement of, 43.
Luminescence, 367.
Luminosity, effect of pressure on, 293.
Luminosity of flames, 292.
Luminosity of flames affected by acety-
lene, 293.
Lunar caustic, 473.
Luray cave, 371.
Luteo compounds, 433.

Lycopodiom powder, 240.
Lye, 333. •

Magnalium, 409.

Magnesia, 384.

Magnesia alba, 386.

Magnesite, 384.

Magnesium, 384.

Magnesium, compounds of, 384-386.

Magnesium nitride, 387.

Magnesium platinocyanide, 497.

Magnesium, separation of from the ele-
ments of the calcium group, 387.

Magnesium silicates, 386.

Magnetite, 419, 423.

Malachite, 460, 465.

Mammoth cave, 371.

Manganates, 437.

Manganese, 436.

Manganese, alloys of, 436.

Manganese blende, 438.

Manganese bronze, 436, 462.

Manganese, hydroxides of, 437.

Manganese, occurrence, preparation,
and properties; 436.

Manganese, oxides of, 436.

Manganese spar, 439.

Manganese tetravalent, 439.

Manganese, valence and properties of,
439.

Manganic acid, 437, 440.

Manganic acid, per-, 437.

Manganic compounds, 439.

Manganites, 437.

Manganous acid, 437, 440.

Manganous salts, 437.

Mannheim gold, 485.

Mantle to Welsbach light, 296.

Marble, 353, 363, 369.

Marignac, on atomic weights, 80.

Marl, 414.

Marsh's method for detecting anti-
mony, 262.

Marsh's method for detecting arsenic,
255.

Martin-Siemens process, 422.

Mass action, 346.

Mass action, law of, 91.

Mass, effect in quantitative analysis, 474.

512

INDEX

Mass, efFect of, on chemical activity, 30.

Massicot, 470.

Mass, law of the conservation of, 8.

Mass on chemical action, effect of, 302.

Mass or size of an atom, 12.

Matches, safety, 330.

Matte, 460.

Membranes, semi-permeable, 428.

Mendel^eff's Periodic System, 137.

Mercuric and mercurous oxides, 400.

Mercuric compounds, 401-406.

Mercuric oxide used in preparing oxy-
gen, 16.

Mercurous compounds, 400-406.

Mercury, fulminating, 405.

Mercury, molecular weights of metals
in, 309.

Mercury molecule is monatomic, 399.

Mercury, properties of, 398.

Mercury, purification of, 398.

Mercury salts, action of ammonia on,
405.

Mercury salts, only slightly dissociated^
401.

Metals, 310.

Metameric compounds, 229.

Metastable equilibrium, 60.

Metathesis, 323.

Methane hydrocarbons, 277.

Methyl orange, 214.

Mica. 374, 407, 415.

Microcosmic salt, 329, 360.

Milk of lime, 364.

Minesite, 345.

Mining, placer, vein, hydraulic, 474.

Minium, 480.

Mirrors, 469.

Mitscherlich, on isomorphism an aid
in determining atomic weights, 77.

Mohr method of determining silver, 474.

Moissan, 272, 296, 414.
.Moissan, apparatus and method for
preparing fluorine, 166.

Molecular conductivities, 112.

Molecular normal solution, 211.

Molecular weights and Avogadro's hy-
pothesis, 72.

Molecular weights, atomic weights
from, 73.

Molecular weights determined from the
densities of gasles, 70.

Molecular weights of dissolved 8ul>-
stances determined by the boiling-
point method, 95.

Molecular weights of dissolved sub-
stances, determined by the freez-
ing-point method, 93.

Molecular weights of gases and of
dissolved substances, 82.

Molecular weights of gases, determina-
tion of, 82.

Molybdenite, 453.

Molybdenum, 453.

Molybdenum, compounds of, 453-
464.

Molybdic acid, 453.

Molybdic acid, phospho-, 454.

Monacid base, 211, 212.

Monazite sand, 306, 417.

Monobasic acid, 212.

Monotropism, 172.

Mordant, 413.

Morley, on the ratio of hydrogen to
oxygen, 80.

Morse's method of preparing chromic
acid, 449.

Morsels method of preparing perman-
ganic acid, 441.

Morse's method of preparing semi-per-
meable membranes, 102.

Mortar, 370.

Mosaic gold, 488.

Muddy water, 48.

Multiple proportions, law of, 10.

Murium, 115.

Nascent hydrogen, 40.

Nature, tlie study of, 1.

Needle valve, 238.

Negative photograph, 471.

Neodymium, 270.

Neon, 241.

Nemst's law of saturation, 318.

Nessler's reagent, 403.

Neutralization, explanation of the con-

sUnt heat of, 218.
Neutralization of acids and bases, 210,

217.

IKDEX

613

Neutralization of weak acids and bases,
219.

New elements predicted by the Periodic
System, 147.

Newland^s octaves, 136.

Niccolite, 433.

Nickel, 433.

Nickel carbonyl or tetracarbonyl, 484.

Nickel, compounds of, 434.

Nickel-plating, 434.

Nickel tetracarbonyl, 279.

Nilson discovered ekaboron — scandium,
149.

Nilson, discoverer of scandium, 416.

Nitric acid, 229.

Nitric oxide, 225.

Nitride of boron, 808.

Nitride of magnesium, 387.

Nitrifying ferment, 341.

Nitrogen, 200.

Nitrogen and hydrogen, compounds of,
209.

Nitrogen and hydrogen, volume rela-
tions in which they combine, 202.

Nitrogen, compounds with chlorine,
bromine, and iodine, 233.

Nitrogen, compounds with oxygen, hy-
drogen, and sulphur, 221, 233.

Nitrogen dioxide, 226.

Nitrogen iodides, 233.

Nitrogen, occurrence and preparation,
200.

Nitrophenol-p, 214.

Nitroprussiate of sodium, 429.

Nitrosulphonic acid, 189, 234.

Nitrosyl chloride, 233.

Nitrosyl-sulphuric acid, 189, 233.

Nitrous acid, 229.

Nitrous oxide, 224.

Nitryl chloride, 233.

Non-electrolyte8» 63.

Normal solution, 112.

Normal solution, equivalent, 211,

Normal solution, molecular, 211.

Octaves of Newlands, 136.
Octivalent, 134.
Oil-lamp, 291.
Olivine, 387.
2l

Olszewski and Wroblewskl, liquefaction

of oxygen, 28.
Onyx, 300.
Opal, 300.
Orthoclase, 414.
Osmium, 491.
Osmium salts, 492.
Osmium tetroxide, 491.
Osmotic pressure, 100.
Osmotic pressure and gas-pressure,

causes of, 107.
Osmotic pressure and gas-pressure,

equality of, 107.
Osmotic pressure and gas-pressure,.

relations between, 104.
Osmotic pressure and the theory of

electrolytic dissociation, 100.
Osmotic pressure, Boyle's law for, 105.
Osmotic preasure, demonstration of, 100.
Osmotic pressure, exceptions to the

applicability of the gas-laws to, 108.
Osmotic pressure, Gay-Lussac's law for,

105.
Osmotic pressure, Morse's method of

measuring, 102.
Osmotic pressure of solutions, Avo-

gadro's law applied to, 106.
Osmotic pressure, Pfeffer's measure-
ments of, 103.
Ostwald, color of permanganates, 443.
Oxidation, slow and rapid, 21.
Oxidizing flame, 294.
Oxygen, 15.

Oxygen, allotropic modification of, 29.
Oxygen and hydrogen, heat energy pro-
duced when they combine, 37.
Oxygen and hydrogen, mixture of

afFected by the presence of certain

substances, 36.
Oxygen and hydrogen, relations by

volume in which they combine, 36.
Oxygen and ozone, different amounts

of energy in, 81.
Oxygen, combination with hydrogen,

35.
Oxygen, difference between, and ozone,

31.
Oxygen, dry, will not combine with dry

hydrogen, 38.

514

INDEX

Oxygen, liquefaction of, 26.
Oxygen means acid-fonner, 124.
Oxygen, names of compounds formed

with, 23.
Oxygen, occurrence in nature, 16.
Oxygen, physical properties of, 24.
Oxygen, power of entering into chemi-
cal combination, 29.
Oxygen, preparation of, 15.
Oxygen, pressure of varies with the

conditions, 24.
Oxygen, properties of liquid, 28.
Oxygen, rdle of in combustion, 10.
Oxygen, substances burn readily in, 16.
Oxygen, transformation of ozone into,

31.
Oxygen used up in combustion, 20.
Oxyhydrogen blowpipe, 37.
Ozone, 29.
Ozone and oxygen, different amounts

of energy in, 31.
Ozone, difference between and oxygen,

31.
Ozone, preparation of, 29.
Ozone, transformation of into oxygen,

31.

Palladic acid, hydrochlor-, 401.

Palladium, 490.

Palladium hydride, 490.

Palladium oxides, 491.

Palmaer demonstration of the solution-
tension of mercury, 391.

Paris, plaster of, 368.

Parke's method, 468.

Passive condition, 421.

Passive stat«, 232.

Pattinson process, 408.

Pebal's experiment, 89.

Periodic System, 130.

Periodic System, imperfections in the,
160.

Periodic System of Brauner, 141.

Periodic System of Mendel^ff and
Lothar Meyer, 137.

Periodic System used to correct atomic
weights, 147.

Periodic System used to predict the
existence of new elements, 147.

Permanent gases, 43.

Permanganates, color of, 443.

Permanganic acid, 437, 441.

Permanganic acid, Morse's method of
preparing, 441.

Permonosulphurio acid, 195.

Petit and Dulong's law, 75.

Pewter, 479.

Pfefifer's measurements of osmotic
pressure, 103.

Phase rule of Willard Gibbs, 59.

Phenolphthaleln, 213.

PhenolphthaleKn cannot be used with
weak acids nor bases, 213.

Philosopher's wool, 388.

Phlogiston theory of combustion, 18,

Phosgene, 279.

Phosphine, 246.

Phosphomolybdic acid, 454.

Phosphonium iodide, 246.

Phosphorescence, 367.

Phosphorite, 242, 372.

Phosphorus, 242.

Phosphorus, acids of, 247.

Phosphorus bronzes, 462.

Phosphorus burns in oxygen, 17,

Phosphorus, compounds with hydro-
gen, 245.

Phosphorus, compounds with oxygen
and hydrogen, 246.

Phosphorus, compounds with the halo-
gens, 253.

Phosphorus, metallic, crystallized, or
black, 244.

Phosphorus oxychloride, 254.

Phosphorus pentachloride, 253.

Phosphorus pentachloride, ▼aix>r de-
composed by heat, 253.

Phosphorus, red, 243.

Phosphorus, strengths of the acids of
252.

Phosphorus, white, 244.

Phosphorus, yellow, 243.

Photochemical reactions, 118.

Photography, 471.

Photometer, Bunsen, 207.

Photophone, 197.

Physical properties and atomic weights,
145.

INDEX

515

Physics and chemistry, relatioiui be-
tween, 1.

Pictet, liquefaction of oxygen, 27.

Pig-iron, 420, 421.

Pink salt, 487.

Pitchblende, 466.

Placer mining, 474.

Plantations, saltpetre, S41.

Plaster of Paris, 368.

Platinic acid, hydrochlor-, 496.

Platinonitrite of potassium, 407.

Platinous acid, hydrochlor-, 496.

Platinum black, 498.

Platinum, chlorides of, 496.

Platinum, colloidal solution of, 495.

Platinum, double cyanides of, 497.

Platinum-iridium, 492.

Platinum, oxides and hydroxides of,
496.

Platinum resistance thermometer,
43.

Platinum sponge, 493.

Platinum, sulphides of, 497.

Platinum, uses of, 494.

Pleonast, 410.

Plumbago, 273.

Plumbates, 480.

Plumbic acid, hydrochlor-, 481.

Plumbic acid, meta-, 480.

Plumbic acid, normal, 480.

Plumbites, 480.

Point, transition, 172.

Pollux, 35').

Polonium, 458.

Polymeric substances, 227.

Polymorphism, 172.

Porcelain, 415.

Portland cement, 370.

Positive photograph, 472.

Potash, 336.

Potash caustic, 336.

Potash, red prussiate of, 428.

Potash, yellow prussiate of, 427.

Potassium amide, 340.

Potassium bromide, 339.

Potassium carbide, 340.

Potassium chlorate, used in preparing
oxygen, 16.

Potassium, detection of, 361.

Potassium does not act with dry oxy-
gen, 334.

Potassium ferricyanide, 428.

Potassium ferrocyanlde, 427.

Potassium hydrazoate, 340.

Potassium hydride, 336.

Potassium hydroxide, 335.

Potassium hydroxide prepared by the
electrolysis of the chloride, 335.

Potassium ion, characteristic reaction
of, 362.

Potassium nitrate solubility curves, 342.

Potassium, occurrence and preparation,
333.

Potassium permanganate, 441.

Potassium peroxide, 336.

Potassium persulphate, 345.

Potassium phosphorus, 348.

Potassium platinonitrite, 497.

Potassium, preparation of, 334.

Potassium pyroantimoniate, 349.

Potassium, salts of, 337-350.

Potassium silicates, 249.

Potassium silicofluoride, 349.

Potassium silver cyanide, 473.

Potassium subchloride, 338.

Potassium tetroxalate, 351.

l^otential of electricity, 361.

I'raseo compounds, 433.

Praseodymium, 270.

Pressure, critical, 60, 283.

Pressure, effect of on luminosity, 293.

Pressure-temperature diagram of sul-
phur, 174.

Pressure-temperature diagram of
water, 58.

Priestley and Scheele discover oxygen,
19.

Prince Rupert drops, 376.

Principle of Soret, 106.

Proportion, law of constant, 9.

Proportions, law of multiple, 10.

Proust, on the law of constant propor-
tion, 10.

Prout's hypothesis, 136.

Prussian blue, 427.

Prussiate, red, of potash, 428.

Prussiate, yellow, of potash, 427.

Pnissic acid, 290.

616

INDEX

Pseudosolutions, 301.
Purple of CassiuB, 476.
l^lrpureo compouirds, 433.
Pyrargyrlte, 467.
Pyrites, 171, 410.
l^rites, arsenical, 256.
Pyrites, copper, 171.
Pyrolusite, 436.

Quadrivalent, 134.
Quartation, 475.
Quartz, 298.
Quinquivalent, 134.

Radiation, uranium, 457.

Radioactive substances, 458.

Radioactivity, 468.

Radiograplis, 467.

Radium, 468.

Ramsay determines the number of

atoms in the molecule of argon, 240.
Ramsay discovers argon, 239.
Ramsay places rare elements in the

atmosphere, in the Periodic System,

151.
Rapid and slow oxidation, 21.
Reaction, catalytic, 187.
Reactions, reversible, 30.
Reactions, reversible chemical, 170.
Heaction, yield of, 180.
Reducing flame, 203.
Reducing power of hydrogen, 38.
Regnault, molecular heats the sum of

the atomic heats, 76.
Resistance of water, 63.
Resistance thermometer of platinum,

43.
Reversible chemical reactions, 30, 170,

348.
Revert phosphoric acid, 373.
Rhodium, 4iK).
Rhodocroisite, 436.
Richards, on atomic weights, 80.
Rock crystal, 300.
Rock salt, 116.
Rj,Ontgen rays, 457.
Rose, on mass action, 346.
Rose's fusible metal, 267.
Rotatory power of sodium chlorate, 321.

Rouge, 423.

Rubidium dioxide, 364.

Rubidium hydroxide, 354.

Rubidium, occurrence, preparation,

properties, 364.
Rubidium, salts of, 354.
Ruby, 407, 400.
Rupert drops, 376.
Rusting of iron, 420.
Ruthenate of potassium, 489.
Ruthenium, 480.

Rutherford, on radioactiylty, 458.
Rutile, 305.

Safety lamp, 205.

Safety matches, 330.

Sal ammoniac, 203, 357.

Saltpetre, 333.

Saltpetre bacteria, 341.

Saltpetre, Chili, 312.

Saltpetre, India crude, d41.

Saltpetre plantations, 341.

Saltpetre, soda, 168.

Salt«, 215.

Salts, acid, 216.

Salts, basic, 216.

Salts, formed from anion of acid and

cation of base, 215.
Salts, nomenclature of, 215.
Salt solutions, explanation of the law

of the thermoneutrality of, 220.
Samarium, 418.
Samarskite, 306.
Sand, 298.
Sandstone, 300.
Sapphire, 407, 400.
Saturated solutions, 61.
Saturation, law of, 318.
Scandium, 416.
Scheele and Priestley discover oxygen,

19.
Scheelite, 455.
Schlippe's salt, 266, 331.
Sch(5nite, 385.
Selcnic acid, 108.
Selenide of hydrogen, 108.
Selenious acid, 108.
Selenium, 197.
Selenium, compounds of, 198.

D^DEX

617

Seleniam, metallic, 107.

Semi-permeable membranes, Morse's
method of preparing, 102, 428.

Senarmontite, 262.

Septivalent, 134.

Serpentine, 384, 386.

Shapleigh, Waldron, 271.

Siderite, 410.

Siemens-Martin process, 422.

Silicates converted into carbonates, 302.

Silicates, double, 414.

Silicon, 208.

Silicon, acids of, 300-301.

Silicon, amorphous, 200.

Silicon bromoform, 303.

Silicon chloroform, 303.

Silicon, compounds with the halogens,
303.

Silicon, compound with carbon — carbo-
rundum, 304.

Silicon crystalline, 208.

Silicon dioi^ide, 200.

Silicon hydrides, 200.

Silicon iodoform, 303.

Silicon, preparation of, 208.

Silver, 467.

Silver, alloys of, 460.

Silver, colloidal, 468.

Silver, German, 388.

Silvering, 460.

Silvering of glass, 460.

Silver ion, 470.

Silver oxides and hydroxid, 470.

Silver-plated wares, 460.

Silver-plating, 460.

Silver, preparation of, 467.

Silver, salts of, 470-474.

Silver, spitting of, 468.

Silver triazoate, 473.

Size or mass of an atom, 12.

Slaked lime, 304.

Slow and rapid oxidation, 21.

Smalt, 432.

Smaltite, 431.

Smithsonite, 387.

Soapstone, 387.

Soda-lime, 366.

Soda saltpetre, 158.

Sodium, 312.

Sodium aluminate, 410.

Sodium amide, 321.

Sodium ammonium phosphate, 320.

Sodium chlorate rotates plane of polar-
ized light, 321.

Sodium, compounds with the halogens,
317.

Sodium dry, does not act on dry sul-
phuric acid, 314.

Sodium halides, solubilities of, 320.

Sodium hydride, 316.

Sodium hydroxide, 316.

Sodium ion, characteristic reactions of,
362.

Sodium ion, the active agent, 316.

Sodium nitroprussiate, 420.

Sodium, occurrence of, 312.

Sodium, preparation of, 313.

Sodium pyroantimoniate, 331.

Sodium, salts of, 310-331.

Sodium subchloride, 310.

Sodium sulphate, solubility curves of,
326.

Sodium sulphite, hydrolytic dissociation
of, 323.

Sodium triazoate, 321.

Sodium water-glass, 331.

Soft solder, 486.

Solder, 3:30.

Solder, soft, 486.

Solfatara, 171.

Solidification of water, heat of, 66.

Solubility, limited and unlimited, 62.

Solubilities of sodium halides, 320.

Solution, equivalent normal, 211.

Solution, molecular normal, 211.

Solutions, pseudo-, 301.

Solutions, supersaturated, methods .of
preparations, 62.

Solutions, unsaturated, saturated and
supersaturated, 61.

Solution-tension of metals and pri-
mary cells, 303.

Solution-tension of metals, demonstra-
tion of, 301.

Solution-tension of some of the more
common metals, the relative, 303.

Solvay or ammonia process of prepar-
ing ammonium carbonate, 827.

618

INDEX

Solvent power of water, 61.

Soot, 274.

Soret principle, 106.

Spec t tic heat of boron varies with the

temperalure, 307.
Specific heat of water, 60.
Specific beats, ut^jmic weights from, 76.
Specific inductive capacity of water, 61.
Spectroscope, 46.
Spiegel bronze, 462.
Spiegel iron, 421, 437.
Spinel, 410.

Spiritus fumans Libavii, 487.
Spitting of silver, 468.
Spoduinene, 352.
Sponge, platinum, 493.
SUnoates, 480.
Siannate, sulpho-, 487.
Stannic acid^ 480*
Staanoaa acid» hydropblor-, 487.
Stannous and stannic hydroxides, 486.
Stannous and stannic oxides, 486.
Stannous chloride, 486.
. Stalactites, S71.
Stalagina, 371.
Stalagmites, 371.
Stas, on atomic weights, 80.
Steam, heat of condensation of, 66.
Steel, 422.

Steel, Bessemer, 422.
StLTL (list nil erism^ 229.
Stibine, 261.
Stibnite, 171.
Still, 49.

Stokes, on Rontgen rays, 467.
Stolzite, 465.

Stoneware or earthenware, 416.
Storage battery, 483.
Storage cells, 193.
Strass, 374.
Stromeyerite, 467.
Strontianite, 376.
Strontium, detection of, 377.
Strontium, occurrence, preparation,

propertiew, 376.
Strontium, salts of, 377.
Study of nature, 1.
Sublimation of iodine, 161.
Substitution, 119.

Sugar of lead, 483.
Sulphide, hydrogen^ 177.
Sulphides, 172.
Sulphides, acid, 181.
Sulphoeyanic acid, 290.
Sill phonic acid, nitro-^ 2S4.

Sulpliur hiiius in oxygen, 17.
Sulphur, compounds with chlorine, 19&
Sulphur, compounds with hydrogen,

177.
Sulphur, compounds with hydrogen and

ojcygen, liig.
Sulphur^ com pounds with oxygen and

hydrogen, 184,
Sulpiiur, compounds with the halogens

and sulphur 105,
Sulphur, difference between the poly-
morphous formfl, 172.
Sulphur, flowers of, 171.
Sulphuric acid, 187.
Sulphuric acid, chlor-, 106.
Sulphuric acid, di-, 194.
Sulphuric acid, dissociation of, 191.
Sulphuric acid dry, does not act on

dry sodium, 314.
Sulphuric acid, nitrosyl-, 189, 233.
Sulphuric acid, per-, 194.
Sulphuric acid, permono-, 195.
Sulphuric acid, pyro-, 194.
Sulphuric acid, thio-, 194.
Sulphur, occurrence and purification

171.
Sulphurous acid, 186.
Sulplmrous acid, hydro-, 194.
Sulphurous acid, strength of, 186.
Sulphur, roll or slick, 171.
Sulphur, temperature- prtBsur© diagram

of, 174.
Sulphur, vapor-density of, 174.
Sulphur water, 48.
Sulphuryl chloride, 196.
Sulphuryl chloride, pyro-, 196.
Supercooling and superheating of water

66.
Supercooling, removal by solid phase

of same substance, 67.
Superheating and supercooling of water

66. '

INDEX

519

Supersatarated solutions, 61.
Supersaturated solutions, methods of

preparation, 62.
Sylvine, 116, 837.
Sylvite, 333.
Sympathetic ink, 432.

Talc, 384, 386.

Tantalum, 271.

Tartar emetic, 263.

Telluric acid, 100.

Tellurious acid, 100.

Tellurium, 108.

Tellurium, compounds of, 100.

Tellurium hydride, 100.

Temperature, critical, 60, 283.

Temperature, determination of the ab-
solute, 26.

Temperature, kindling, 205.

Temperature-pressure diagram of sul-
phur, 173.

Temperature-pressure diagram of water,
68.

Tempering of steel, 423.

Ternary electrolyte, li2.

Thallium, 417.

Thanhs experiment, 00.

Thenardite, 323.

Thermometer of Beckmann, 05.

Thermometer, platinum resistance, 43.

Thermoneutrality of salt solutions, ex-
planation of the law of, 220.

Thilorier*8 mixture, 283.

Thionic acids, poly-, 105.

Thomas-Gilchrist converter, 422.

Thomas slag, 422.

Thomson, J. J., electrolyzed hydrogen,
45.

Thorite, 306, 418.

Thorium, 306.

Tin, 484.

Tin, allotropic forms of, 485.

Tin, alloys of, 485.

Tin-butter, 487.

Tin, cry of, 484.

Tincture of iodine, 161.

Tin ions, 485.

Tin-pest, 485.

Tin, preparation and properties of, 484.

Tin-salt, 487.

Tinstone, 484.

Tin, sulphides of, 487.

Titanic acid, 305.

Titanite, 305.

Titanium, 305.

Titanium, compounds of, 905.

Topaz, oriental, 400.

Tourmaline, 352, 387.

Towers, Glover and Gay-Lussac, 188.

Transition point, 172.

Traube, preparation of semi-permeable

membranes, 100.
Tree, lead, 470.
Triacid base, 211.
Triads of Dobereiner, 136.
Triazoic acid, methods of formation,

208.
Tribasic acid, 212.
Triphylite, 352.
Trivalent, 134.
Tungsten, 454.
Tungsten, chlorides of, 455.
Tungstic acid, 455.

Tungstic acid, colloidal solution of, 455.
Tungstic acids, poly-, 456.
TurnbulPs blue, 428.
Type metal, 470.

Ultramarine, 415.
Ultramarine, yellow, 450.
Univalent, 134.
Univalent ions, 361.
Unsaturated solutions, 61.
Uranates, 457.
Uranic acids, 457.
Uranite or pitchblende, 456.
Uranium, 456.
Uranium, chlorides of, 457.
Uranium, oxides of, 456.
Uranium radiation, 457.
Uranium yellow, 457.
Uranyl group, 457.

Valence, variable, 406.

Valency, 134.

Valency, Faraday's law the basis of, 134

Valve, needle, 288.

Vanadium, 270.

520

INDEX

Van*! Hoff on osmotic preflsnre, 105.
Vapor-densities, abnormal, 87.
Vapor- denaitiea, abnormal, explanation

of, 88.
Vapor-density measurements, results

of, 87.
Vapor-density of sulphur, 174.
Vaporization of water, heat of, M.
Vapors, dissociation of, diminished by

an excess of one of the products of

dissociation, 91.
Vapor-tension of water in its different

states of aggregation, 58.
Variable valence, 400.
Vein mining, 474.
Velocity of a reaction, 101.
Verdigris, 405.
Vermilion, 404.
Victor Meyer, gas-di^lacement method

of, 85.
Victor Meyer, on the vapor-density of

iodine, 101.
Vinegar of lead, 483.
Vitriol, blue, 404.
Vitriol, green, 420.
Volatile compound, whenever it can be

formed it is formed, 191.
Volhard^s method of determining silver,

474.
Volume, critical, 284.
Volume relations in which hydrogen

and chlorine combine, 121.
Volume relations in which hydrogen

and oxygen combine, 30.
Volumes, atomic, 145.
Volumes, atomic, curve of, 140.
Volumetric method, 474.

Waage and Guldberg's law of mass
action, 92.

Waldeii*8 work on liquid sulphur di-
oxide, 185.

Wanklyn and Robinson's experiment,
IK).

AVater, analytical method of determin-
ing the composition of, 51.

Water and hydrogen dioxide, 40.

Water and hydrogen dioxide, relations
between, 07.

Water m it oocurs in niUiire is imp

46.
Water, a stable compound, 52.
Water, boiling-point of, 54, 63.
Water, composition of, 50.
Water, conductivity of, 63.
Water, dielectric constant of, 61.
Water, dissociating power of, 63.
Water, effervescent, 48.
Water, electrolysis of, 50.
Water, freezing of, 55.
Water, freezing-point of, 63.
Water-gas, 279, 292.
Water-glass, 831.
Water, heat of condensation of aU

and solidification of, 56.
Water, heat of vaporization, 54.
Water, muddy, 48.
Water, not an element, bat a o

pound, 50.
Water of crystallization, 46.
Water of hydration, 46.
Water, physical properties of, 54, 6<
Water, properties affected by diasoi

substances, 03.
Water, purification of, 48.
Waters, hard. 47, 371.
Waters, mineral, 47.
Water, solvent power of, 61.
Water, specific heat of, 60.
Water, specific inductive capacity

01.
Waters, permanently hard, 372.
Waters, temporarily hard, 371.
Water, sulphur, 48.
Water, superheating and supercool

of, 50.
Water, synthetical method of dei

mining the composition of, 52.
Water, temperature-pressure diagi

of, 58.
Water, vapor-tension of in its differ

states of aggi'egation, 58.
Weak acids and bases, explanation

results with, 219.
Weak acids and bases, neutralizat

of, 219.
Weight, increase of in combustion,
Weights, law of combining, 11.

INDEX

521

Welding of iron, 420.

Weldon mud, 440.

Weldon process for making chlorine,
110, 438.

Welsbach light, 290.

Weston cadmium element, 405.

Wheatstone bridge, 113.

White lead, 482.

White, permanent, 380.

Williamson^s electrochemical theory,
110.

Winkler discovered ekasilicon — germa-
nium, 149, 305.

Witlierite, 378.

W5hler, first prepared aluminium, 407.

Wolframite, 455.

WoUastonite, 374.

Wood burns in oxygen, 16.

Wood's fusible metal, 267.

Wool, philosopher's, 388.

Wroblewski and Olszewski, liquefaction
of oxygen, 28.

Wrought-iron, 421.

Wulfenite, 453, 478.

Xenon, 241.

Yellow pruBsiate of potash, 427.
Yellow ultramarine, 450.
Yield of the reaction, 180.
Ytterbium, 417.
Yttrialite, 417.
Yttrium, 417.

Zeolites, 415.

Zero absolute, can it be realized experi-
mentally ?, 44.

Zero absolute, of temperature, deter-
mination of, 26.

Zinc, 387.

Zinc amalgam, 388.

Zinc blende, 171, 387.

Zinc dust, 387.

Zinc, galvanized, 388.

Zinc, granulated, 387.

Zinc oxide and hydroxide, 388.

Zinc precipitated by copper, 466.

Zinc, salts of, 389^90.

Zinc, use of in primary batteries, 391.

Zinc white, 388.

Zircon, 306.

Zircon, compounds of, 306.

Zirconium, 306.

The Theory of Electrolytic Dissociation

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