Full text of "The Foundations Of Chemical Theory"

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THE FOUNDATIONS OF
CHEMICAL THEORY

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THE

FOUNDATIONS OF
CHEMICAL THEORY

The Elements of
^Physical and General Chemistry

R. M. CAVEN

D.Sc.(London), F.I.C.

Professor of Inorganic and Analytical Chemistry
in the Royal Technical College, Glasgow

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1929

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PREFACE

Modern chemistry presents a multitudinous array of facts. It
could not be otherwise, since the science undertakes to describe
more than eighty separate elements and the compounds they form,
one with another.

The exhibition of these facts, as set forth in a comprehensive
textbook, whether of organic or inorganic chemistry, may well
dazzle and bewilder the student, who will scarcely be likely to
appreciate the beauty of the science, or the glamour of the human
achievement of which it is the monument, if he is required labori-
ously to appropriate the facts without judicious selection and
arrangement.

Indeed, the complaint is often made by students that chemistry
requires too much memory-work, and is therefore not so inspiring
a science as physics, which deals all the while with fundamental
principles, and properties of matter.

Yet the facts of chemistry are the facts of nature, and nature is
not chaotic. Consequently these accumulated facts present to the
mind of man a powerful challenge. They must be systematized.

Chemical theory, however, is worthy of the facts which form
the subject-matter of the science,' for it includes some of the
greatest generalizations of any science. The atomi'c and molecular
theories, the theories of molecular structure, the periodic law, the
conceptions embraced in modern physical chemistry, which con-
stitute, par excellence, the science on its intellectual side, are among
the noblest achievements of the human mind.

Every student of chemistry must become acquainted in some
degree with chemical theory. The question therefore arises how
this theory may best be communicated. Ought it to be inter-
mingled dexterously with chemical facts with the descriptive
part, so much easier to understand and more difficult to remember;

PREFACE

or ought a special course of instruction in the theory of the science
to be devised, apart from the descriptive lectures?

The exigencies of teaching generally lead to the following solu-
tion: that the student learns enough theory during his second year
or intermediate course to serve him for the time being, and that in
the later course for his degree he attends lectures in physical
chemistry, which take for granted a solid foundation of elementary
theory.

A degree course in physical chemistry is a serious affair, and in
the hands of a specialist is likely to make a severe strain on the
mental resources of a student. To profit by such a course a student
should be very sure of his foundations.

The purport of this book may therefore now be stated. As
its title indicates, it is an endeavour to disclose the foundations
of the science, and make them plain and real to the average
student.

The author hopes that the general reader, who wishes to know
what modern chemistry really means, will find within these pages
the information he desires; and that the student to whom chemical
science offers an open field of glowing possibilities will find the
chapters of this book a not unwelcome guide in his earlier ex-
cursions.

Briefly, it is suggested that the book may be read by the
student during or at the end of his second year's course, for the
purpose of knitting together his chemical knowledge in view of the
more advanced studies which lie before him later.

The writer has utilized facts and methods of presentation con-
tained in former books of which he is joint or sole author; and he
desires to acknowledge his indebtedness to various other works,
particularly to the introductory volume of A Textbook of Inorganic
Chemistry, by Dr. J. Newton Friend. He also wishes to thank
Dr. E. B. 11. Prideaux for kindly reading the manuscript of the
book.

ROYAL TECHNICAL COLLEGE,
GLASGOW, Uctobti ,

PREFACE TO NEW EDITION

When this book was published five years ago it was deemed
wise to omit from it any reference to new ideas in chemistry con-
nected with the constitution of the atom, and all that is entailed in
recent knowledge concerning the structure of matter. So far as he
is aware, the author was not criticized for the omission of teaching
which at that time seemed appropriate for advanced but not for
elementary students of chemistry.

It was not long, however, before the author himself began to
regret this omission, and to look forward to the time when the call
for a new edition might enable him to include some of the new
teaching which now undoubtedly forms an important part of the
Foundations of Chemical Theory.

How to incorporate the new knowledge with the old how
safely to pour new wine into old wine-skins was a difficult
problem. Indeed it is a problem which at the present time is
encountered by every teacher of chemistry who desires to keep
his science up-to-date. The danger is that old doctrine may be
utterly forsaken for new, and that instability and confusion may
in consequence be introduced into chemistry. A wise conservatism
is to be commended with regard to all new teaching in science
festina lente must be our motto. Moreover, the historic method of
study is always wise. Therefore this method has been adopted in
dealing with the new material. The earlier chapters have been
revised in the light of what was to follow, and so have been made
a pathway of approach, clear of stumbling-blocks, to the new
territory.

Thus we have the older atomic and molecular theories, the
older views of valency and chemical constitution, and the periodic

vii

PREFACE TO NEW EDITION

law according to Mendeteeff, as a preparation for the modern view
of the atom, and the modern view of the molecule.

So it is respectfully suggested that teachers will find this a safe
and helpful way of leading their students into possession of the
new domain.

The author wishes to tender his grateful thanks to Dr. J. A.
Cranston for the help he has given in reading and criticizing the
new matter for this edition.

E. M. C.

ROYAL TECHNICAL COLLEGE,
GLASGOW, October, 1925.

NOTE TO THIRD EDITION

Advantage has been taken of this new edition to make some
corrections and improvements in the text, and to add in an
Appendix brief accounts of Chemical Nomenclature and Symbols,
Hydrion Concentration and pH Value, Amphoteric Hydroxides,
and Ionic and Electronic Equations.

R M. C.

ROYAL TECHNICAL COLLEGE,
GLASGOW, February, 1929.

CONTENTS

PART I ATOMS AND MOLECULES

CHAP. Page

I. THE OLDER ATOMIC AND MOLECULAR THEORIES J

Composition of Matter The Elements Laws of Chemical Com-
binationThe Atomic Theory The Molecular Theory Law of
Volumes Avogadro's Theory.

II. EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS - - - 18

Equivalent and Atomic Weights Standard for Equivalent and
Atomic Weights Determination of Equivalent Weights
Methods of Determining Atomic Weights: (a) Method of
Vapour Density and Avogadro's Theory Vapour Densities
by the Methods of Dumas, Hofmann, Victor Meyer;
(6) Method of Chemical Displacement; (c) Method of Specific
Heats (Dulong and Petit's Law); (d) Method of Isomorphism
(Mitscherlich's Law) ; (e) Method of the Periodic Law Illustra-
tion: The Atomic Weight of Carbon Molecular Weights in
Solution: Cryoscopic and Ebulliscopic Methods Molecular
Complexity Molecular Compositions of Compound Gases.

III. OLDER VIEWS OF VALENCY AND CHEMICAL CONSTITUTION 56

Historical Definition of Valency Bonds and Graphic Formulae
Variability of Valency Double Bond in Carbon Compounds
1 * Chemistry in Space "Criterion of Valency Nature of Valency.

IV. CLASSIFICATION OF THE ELEMENTS 63

The Periodic Law according to Mendeleeff Law of Octaves
Development of the Periodic Law Periodicity of Physical
Properties Periodicity of Chemical Properties Periodicity
of Valency Uses of the Periodic Law Correction of Atomic
Weight Values Suggestiveness of the Periodic Law Objec-
tions to the Periodic Law.

V. THE MODERN VIEW OF THE ATOM

Electrolysis and the Electron Radioactivity The Electron
and Valency Effect of Radioactive Change Isobares and
Isotopes Atomic Number The Modern Statement of the
Periodic Law Positive Ray Analysis and Mass Spectra Iso-
topes of Lead Atomic Structure Valency Theories of Lewis
and Langmuir Theory of Bohr.

x CONTENTS

CHAP. Page

VL THE MODERN VIEW OF THE MOLECULE 118

Electrochemical Theory of Berzelius Dualistic and Unitary
Systems Electrolytic Dissociation Theory of Arrhenius
Electrovalency and Covalency The Octet Theory Molecular
Structures Space Lattices Crystal Units and Molecules of
Solids Formula, Old and New.

PART II STATES OF MATTER

VII. THE PROPERTIES OF GASES 134

States of Matter Gas Laws of Boyle and Charles The Gas
Equation Diffusion of Gases; Effusion; Atmolysis Dalton's
Law of Partial Pressures Deviations from Boyle's Law Pro-
perties of Carbon Dioxide Andrews's Experiment Critical
State Liquefaction of Gases Methods of Liquefaction: I
Simple Compression; II Cascade Method of Cooling; III
Method of Adiabatic Expansion ; IV Method of Self-intensive
Refrigeration Liquid Air Liquefaction of Hydrogen and
Helium Practical Applications of Liquefied Gases : Ammonia,
Chlorine, Carbon dioxide, Sulphur dioxide, Air.

VIII. THE PROPERTIES OF LIQUIDS 158

Mobility Viscosity Density Specific and Molecular Volume
Vapour Pressure and Boiling-point Distillation under Atmos-
pheric and Reduced Pressure Relation between Boiling-points
of Liquids in Homologous Series.

IX. THE PROPERTIES OF SOLIDS 166

Formation of Solids Solidification of Vapours; Sublimation
Solidification of Liquids Melting-points of Solids Melting-
points of the Elements Melting-points of Carbon Compounds
Formation of Solids from Solution Solidification of Mixtures ;
Cryohydrates and Eutec tics Crystals and Crystallography -
Crystallization Polymorphism and Allotropy.

X. SOLUTIONS 189

Solvent, Solute, Solution Solutions of Gases in Liquids
Solubilities of Gaseous Mixtures Solutions of Liquids in Liquids
Distillation of Mixed Liquids ; Fractional Distillation Solu-
tions of Solids in Liquids Solubility The Process of Solution
Influence of Temperature on Solubility; Solubility Curves
Relation between Chemical Composition and Solubility.

XL THE PROPERTIES OF DILUTE SOLUTIONS 208

Electrolytes Electrolytic Dissociation Laws and Process of
Electrolysis Solution Pressure Chemical Reactions in Solu-
tion Colours of Salt Solutions Complex Ions Theory of
Indicators.

CONTENTS xi

CHAP. p age

XII. THE COLLOIDAL STATE 223

Solution and Suspension Crystalloids and Colloids Hydrosol
and Hydrogel- Typical Colloids The Ultra-microscope Size
of Colloidal Particles Gradations between Suspension and
Solution Coagulation of Colloids Cataphoresis Peptization
Protective Colloids Phases and Media Range of Colloidal
Phenomena.

PART III CHEMICAL COMPOUNDS AND
CHEMICAL CHANGES

XIII. TYPES OF CHEMICAL COMPOUNDS 233

Hydrides Oxides and Hydroxides Bases, Acids, and Salts
Hali des Sulphides Oxy -salts Hydrated Salts Double and
Complex Salts.

XIV. CHEMICAL CHANGE IN GENERAL 254

A Typical Reaction Reversible Reactions Chemical Equili-
brium Thermochemistry Rate and Limits of Chemical Change
Catalysis.

XV. CHEMICAL CHANGES CLASSIFIED 273

Chemical Action of Heat on Compounds Thermal Dissociation
Thermal Decomposition Chemical Interaction of Water with
Elements and Compounds Hydrolysis Chemical Interaction of
Acids with Metals Interaction of Nitric Acid and Metals-
Effect of Solubility and Volatility on Chemical Change Oxida-
tion and Reduction.

XVI. EQUATION-BUILDING 305

The Meaning of a Chemical Equation Construction of Chemical
Equations : Typical Examples.

APPENDIX :

Chemical Nomenclature and Symbols 315

Hydrion Concentration and pH. Value 318

Amphoteric Hydroxides 316

Ionic and Electronic Equations 320

INDEX 325

THE FOUNDATIONS OF
CHEMICAL THEORY

PART I ATOMS AND MOLECULES

CHAPTER I
THE OLDER ATOMIC AND MOLECULAR THEORIES

i. The Composition of Matter

The term matter at first suggests to the unsophisticated mind
such qualities as bulk, shape, colour, hardness, weight. It is quickly
recognized, however, that some of these qualities belong only to
some kinds of matter, and are absent from others. Consider, for
example, a log of wood floating on water. It will be admitted that
the wood but not the water possesses hardness, though both possess
weight; that the log but not the water has a permanent shape,
though both have bulk. Above the water is the air, and the air
is something, for it blows in a man's face, and raises ripples on
water. By a proper instrument it can be shown that air possesses
weight. So air is matter, though it is without form, permanent
bulk, apparent colour, or hardness.

At the conclusions suggested by these thoughts the Ancients
arrived after their own fashion, Their fundamental classification
of natural things included the three categories: earth, water, air,
together with a fourth fire. These four were the elements, accord-
ing to Aristotle, but their names really stood for qualities rather
than separate species of matter. For instance, earth meant dryness
and coldness, water wetness and coldness, and so on. Nevertheless,
the first three terms at least suggest an outlook on the world which
was essentially true, since they stand for the three fundamental
forms of matter: solid, liquid, gas.

(D60) 1 2

2 CHEMICAL THEOKY

This conclusion as to the threefold constitution of the world
is reached by an extensive outlook upon nature; an intensive, an
introspective view, such as the following illustration furnishes,
leads to another conclusion.

Sea-water is distinguished from fresh water by its saltness, that
is, by its special taste. What proportion of salt water mixed with
fresh water could be so distinguished depends upon the sensitive-
ness of the human palate; but such a test would assuredly fail
when the salt water was highly diluted. The addition of silver
nitrate to the much diluted salt water would, however, serve to
detect the presence of salt after the test of taste had failed, because
of the turbidity or opalescence which the silver nitrate produces
with even very small quantities of chlorides in solution. Would
this test fail in its turn when the utmost delicacy was required, or
would it detect the minutest quantity of salt? It might fail by
reason of defective human vision, were it not that an instrument
has been made on purpose to detect the slightest cloudiness in
a liquid; but it must fail at last for quite another reason. For the
test depends on the insolubility of silver chloride in water; but
silver chloride is not quite insoluble in water, and on this account
will not be precipitated when the salt solution is excessively
dilute.

So salt may perhaps be present in, and diffused through, water
in quantity too minute to be detected by any test whatsoever.
How far, then, may the dilution be carried; will salt still be present
after infinite dilution? Or, more generally, is matter infinitely
divisible? This is the question which arises directly out of an
experiment which any novice in chemistry can perform. The
same question presented itself to the alert minds of the ancient
peoples of the East; not, it is true, by reason of experimental
investigation, but because of meditation on the nature of the
material world. Thus the question was: Is matter infinitely divis-
ible or not? To believe the latter is the easier and more satisfying
philosophy; this was the philosophy of Plato and Democritus. So
matter was supposed to consist ultimately of hard, indivisible, and
indestructible particles separated by vacuous interspaces; that is, of
atoms.

Of the two theories of the Ancients, to which the Greeks gave
finished expression the theory of Elements and the theory of
Atoms the former passed through strange vicissitudes, which

THE ATOMIC AND MOLECULAR THEORIES 3

need not here be traced; whilst the latter remained latent until
modern times, when it was found to be in accord with the con-
clusions derived by Dalton from experimental data.

2. The Elements

A theory of the elements should precede a theory of atoms.
So, dismissing the ancient theory of the elements, it may be said
quite briefly that an element, as generally understood, is an ultimate
species of matter; or, to adopt the more usual and explicit definition:

An element is a substance which hitherto has not been resolved
into two or more dissimilar kinds of matter.

We owe this idea of an element first of all to Boyle (1678); it
was Lavoisier (1789), however, who realized its provisional nature;
and, indeed, some of Lavoisier's elements, such as lime and the
alkalis, are now proved to be compounds. If this definition merely
marked the present state of progressive human achievement it
would not be a scientific definition. An assurance is necessary that
some kind of finality has been or may be reached in the decom-
position of substances; that at least the "elements" are equally
elementary; this assurance may be given with every confidence.
That the elements are absolutely undecomposable has, however, never
been a settled belief of the chemist; on the contrary, he has held the
opinion from time to time that they are derived from, and so are
resolvable into, a common primordial substance. The phenomena
of radioactivity now furnish evidence of the spontaneous and
perpetual disintegration, into simpler forms of matter, of the atoms
of certain of the elements; and new theories of matter which are now
taking firm root in chemistry represent the atom as a complex struc-
ture; further, the fact has now been established that the elements
are not truly homogeneous, but consist of atoms of differing relative
weights, which, however, are indistinguishable and inseparable by
ordinary chemical means.

These considerations, however, do not affect the chemist's working
theory of the elements. He knows that the nearly ninety different
kinds of matter into which he has resolved the many substances
found in nature, and out of which he can elaborate a vast number
of compounds to which nature has no counterpart, pass unchanged
through the crucible of his everyday operations. So he habitu-
ally regards the catalogue of the elements that hangs in his

4 CHEMICAL THEORY

laboratory as a permanent record, not only of human skill, but
of Nature's handiwork as well.

3. The Atomic Theory

The atoms of Greek philosophy were indestructible; indeed, the
indestructibility of matter has probably always been an axiom
of science, notwithstanding the surprising and fantastic changes
matter was supposed to undergo in the hands of the alchemists
of the Middle Ages. This principle was first clearly illustrated,
however, by Lavoisier in his application of quantitative methods to
chemistry, and was subsequently demonstrated, within the limits of
the most accurate experimental research, by Stas, Landolt, and
others. 1 In 1770 Lavoisier gave an account of experiments he had
performed to test the supposition that water is transformed into
earth by boiling. A weighed quantity of water was boiled for
101 days in a weighed and sealed glass vessel; and at the end
of that time it was found that while "earth" appeared in the
vessel the water weighed the same as at first, and the weight
of the "earth" was equal to the loss in weight which the glass
vessel had incurred. Thus it was shown that the "earth" came
from the glass and not from the water, and that water is not
transformed into earth by boiling. In these experiments the use
of the balance played an essential part; but this was a novelty in
chemistry. The scientific achievements of such men as Boyle,
Black, Cavendish, Priestley, Scheele, notwithstanding their great
value, were chiefly of a qualitative nature. Henceforth, however,
chemistry was concerned with weighing things, and a new era
began.

It was now but a step to the quantitative analysis of chemical
substances. Lavoisier took this step in his investigation of mer-
curic oxide, or the calx of mercury, which Priestley and Scheele
had decomposed into mercury and oxygen. Soon there arose an
important question, the answer to which could be found only by
quantitative analysis. This was the question: Is a chemical com-
pound necessarily constant in composition, or may its composition
vary within certain limits according to the way in which it is
prepared ?

It may appear to be a truism that the same compound must

*For an account of these researches see The Study of Chemical Composition, by I. Freund.

THE ATOMIC AND MOLECULAR THEORIES 5

always have the same composition, so that it is better to state the
problem in this way: Can the products of different chemical re-
actions, designed to produce the same compound, really differ
slightly in composition? /Berthollet was of opinion that they
could; that the composition of a compound might vary within
certain limits according to the way in which it was prepared;
indeed, that the conditions of its genesis are the overruling factors
of its composition. Barium sulphate was cited as an example. All
known specimens of this compound were found to be identical in
composition, but this identity was due, not to any inherent pro-
perty of the constituent elements of the compound, but to the fact
that by uniting in such proportions these elements produced a
compound of maximum insolubility in water. 'It was fair to suppose,
therefore, that if the salt could be precipitated from some other
medium than water it would have a different composition accom-
modated to a new requirement of maximum insolubility. Such an
idea was gravely erroneous, and was quite foreign to the principles
on which the atomic theory was soon to be founded. Yet the idea
appeared to have experimental support; and, indeed, it contained
the germ of an important truth. In support of his belief, Berthollet
showed that when nitric acid reacted with mercury or with tin the
composition of the nitrate of mercury or oxide of tin produced
varied within certain limits according to the concentration of the
acid employed. Proust, on the other hand, maintained that " be-
tween pole and pole compounds are identical in composition; their
appearance may vary owing to their manner of aggregation, but
their properties never". After a controversy carried on with
BertKollet over a period of eight years (1800-8), Proust fully
established his proposition, and showed that the variable products
obtained by Berthollet were variable mixtures of invariable com-
pounds. Thus was established the first law of chemical combina-
tion the law of definite or fixed proportions:

The same chemical compound always contains the same elements
united together in the same proportions; or the proportions between
the constituent elements of a chemical compound bear an unalter-
able relation to each other, and to the proportion of compound
formed.

This was the first foundation of the atomic theory. ?

Nevertheless, it was a pity that the truth in Berthollet's view
was entirely overlooked iri the victory of Proust: the truth that

6 CHEMICAL THEOKY

the proportions or concentrations in which reacting substances are
present may determine the proportions subsisting between the
products of a reaction, although the proportions in which elements
or compounds actually react to form these products are quite be-
yond the influence of external and variable conditions. Thus, in
the case of the action of nitric acid on mercury, studied by Ber-
thollet, the concentration of the acid determines whether mercurous
or mercuric nitrate or a mixture of these two salts is produced,
although it can have no influence on the unalterable chemical
composition of either of the two salts.

The existence of two nitrates of mercury is, however, a note-
worthy fact, which appears the more striking when it is discovered
that in one compound the proportion of mercury to nitrate is
exactly twice what it is in the other. Further examples of this
phenomenon were observed by Dalton, who showed that the pro-
portion of hydrogen to a fixed quantity of carbon is twice as
great in methane as in ethylene, and of oxygen to a fixed quan-
tity of carbon, twice as great in carbonic acid gas as in carbonic
oxide.

Other examples of compounds showing analogous relations are
two of the oxides of lead, one of which contains twice as much
oxygen compared with lead as the other, and the five oxides of
nitrogen, in which the quantities of oxygen combined with a fixed
amount of nitrogen are as 1 : 2 : 3 : 4 : 5. Here was an important
generalization, which was formulated by Dalton as the laiu of
multiple proportions:

When one element combines with another in more than one pro-
portion, these proportions bear a simple relation to one another.

The foregoing facts furnish material enough for the atomic
theory. It is usual, however, to add to the laws of definite and
multiple proportions a third law, the law of reciprocal proportions^
which, however, follows logically from the other two laws.

It was shown by Richter, about 1780, that the ratio between
the quantities of two acids which neutralize a fixed amount of
alkali is the same whatever the alkali may be; and by Berzelius, in
1810-2, that 381 parts of lead combine separately with 58-73 parts
of sulphur and 29 6 parts of oxygen, whilst 58-73 parts of sulphur
combine in turn with 57 45 parts of oxygen. Now, 57 45 = 29 6 x 2
within the limits of the experimental error of the time; and these,
facts may be expressed diagrammatically thus: ^

THE ATOMIC AND MOLECULAR THEORIES

Sulphur -< >- Oxygen

So is illustrated the law of reciprocal proportions: v

The proportions of two elements which separately combine with
a fixed proportion of a third element are also the proportions of
these elements which combine with each other, or else in accord-
ance with the law of multiple proportions they bear a simple ratio
to these proportions.

This law has Within it, especially in the way in which it was
illustrated by Richter, the idea of chemical equivalents; and so it
may be stated in this axiomatic way: v

Quantities of substances which are chemically equivalent to the
same quantity of a third substance, are chemically equivalent to one
another.

Thus it appears that, granted the validity of the idea of chemical
equivalents, which will be examined later, the law of reciprocal
proportions requires no experimental justification.

Although we owe the essence of the modern atomic theory to
Dalton alone, the precise way in which the theory took shape in
the mind of its author has been rather problematical. At the
close of a paper on the absorption of gases by water, Dalton wrote
as follows:

"An inquiry into the relative weights of the ultimate particles of
bodies is a subject, as far as I know, entirely new. I have lately been
prosecuting this inquiry with remarkable success."

i/ No hint is given in the context of the way in which the atomic
values, which follow, were estimated, nor of the precise reason why
such values were believed to exist. The idea that matter consists
of discrete particles was, however, in the air. Apart from the
ancient theory of atoms, a theory of particles had been held more
or less firmly by F. Bacon, Boyle, Higgins, and others; whilst
Newton made the following explicit statement:

11 It seems probable to me, that God in the beginning formed matter

CHEMICAL THEORY

in solid, massy, hard, impenetrable, movable particles, of such sizes and
figures, and with such other properties, and in such proportion to space,
as most conduced to the end for which He formed them; and that these
primitive particles, being solids, are incomparably harder than any porous
body compounded of them, even so very hard as never to wear or break
in pieces; no ordinary power being able to divide what God Himself
made one in the first creation. . . . The changes of corporeal things are
to bo traced only in the various separations and new associations and
motions of these permanent particles." *

It would almost appear from such a pronouncement that
Newton and not Dalton was the author of the atomic theory.
Yet this statement is not a chemical theory: it is a cosmic theory
intimately related to Newton's great discovery of universal gravi-
tation. Dalton, however, was greatly indebted to Newton and
the idea of ubiquitous particles which the theory of gravita-
tion involved; and it appears that he conveyed this idea into
chemistry and employed it to explain the laws of chemical com-
bination.

To understand the atomic theory, therefore, is simply to under-
stand how the theory of particles fits the chemical laws. This is
quite easy.

Let there be three elements, A, B, C, and let the areas of the

squares:

represent the combining

weights of these elements on any arbitrary scale,

being the

quantity of B which is found to combine separately with 1 A j

parts of A and

proportions; whilst

parts of C, according to the law of definite

parts of C also combine with [A I

parts of A, according to the law of reciprocal proportions, so that
the following compounds are formed:

THE ATOMIC AND MOLECULAR THEORIES 9

Then, according to the law of multiple proportions, compounds
such as these may be formed:

1 A

&c. How else can these experimental facts be interpreted than by
the idea of " permanent particles " ? The elements combine accord-
ing to the laws of definite and multiple proportions because they
combine atom by atom: 1 atom of A with 1 atom of B; 1 atom
of A with 2 atoms of B; 2 atoms of A with 1 atom of B; and
so on. That is Dalton's atomic theory, and the theory is expressed
succinctly in the following statements:

1. All matter consists of discrete particles called atoms, which
remain unbroken throughout chemical change.

2. Atoms of the same element are ordinarily supposed to be
similar in all respects.

3. Chemical compounds are formed by the union of the atoms of
different elements in simple numerical proportions.

4. The proportions in which, elements combine to form compounds
are determined by the atomic weights of the elements.

The transition from the laws to the theory is quickly made: it
is taken, so to speak, in a stride; but the boundary line between
them must not be obliterated. The laws of chemical combination
are statements of experimental facts; the theory is an explanation
of these facts which is very probably true, but it does not stand
in the same category as the facts. In science, facts and theory
must always be distinguished as clearly as possible.

When the atomic theory is accepted it at once appears that
the combining weights of the elements represent the combining
weights of their atoms. The atomic theory involves the atomic
weights. No atomic theory previous to that of Dalton involved
atomic weights; these were a novelty, and their introduction
constitutes Dalton's great contribution to chemical science. The
following atomic weights are selected from a list published by
Dalton, the atomic weight of hydrogen being 1.

CHEMICAL THEORY

DALTON'S ATOMIC WEIGHTS

Hydrogen ...

An atom of water or steam,

Azote-

composed of 1 of oxygen

Carbon or Charcrn

+ 1 of hydrogen

Oxygen

T*l 1

Phosphorus

An atom of ammonia, com-

Sulphur
Magnesia ...

posed of 1 of azote + 1
of hydrogen

Lime
Soda

TlYiTI

An atom of carbonic oxide,
composed of 1 of carbon

JLILHl . .

Potash

+ 1 of oxygen

Zinc...

A 4- e v. 'i

Copper
Silver

56
100

An atom ot caroomc aciclj
1 carbon + 2 oxygen

Gold

140

An atom of sulphuric acid,

Mercury ...

167

1 sulphur + 3 oxygen

If the student compares these atomic weights with those in
use at the present day, he will see that they differ widely from
the modern figures. Inaccuracies in Dalton's values are to be
expected, but it is not experimental error which attributes, for
example, an atomic weight of 7 to oxygen, instead of 16. As a
matter of fact these combining weights are not atomic weights at
all, but are approximately what we now recognize as equivalent
weights.

For, in truth, Dalton had no means of determining atomic
weights. The value 7 (or 8) for oxygen is derived from the
analysis of water: 8 parts by weight of oxygen combine with
1 part by weight of hydrogen to form 9 parts by weight of water.
Who shall say from this that the atomic weight of oxygen is 8?
That depends on the number of atoms of each element which
combine together to form a unit of water, a fact clearly recognized
by Dalton.

Thus, we have the ratio O : H = 8 : 1 or 16 : 2 or 24 : 3, &c.,
and if 1 atom of oxygen combines with 1 atom of hydrogen,
then the atomic weight of oxygen is 8; if 1 atom of oxygen
combines with 2 of hydrogen, the atomic weight of oxygen is 16;
if 1 combines with 3, it is 24; if 2 combine with 1, it is 4; and
so on. There was, however, no evidence on which to base a
decision between these alternatives. Just at this point Dalton
made a regrettable mistake. Instead of recognizing the limita-
tions of his experimental knowledge, he made the assumption

THE ATOMIC AND MOLECULAR THEORIES 11

that since only one compound of hydrogen and oxygen was
known, it necessarily had the simplest possible composition, and
so was formed from 1 atom of each of its constituent elements.
Consequently, the atomic weight of oxygen was thought to be
7 (or 8); and for a similar reason the atomic weight of nitrogen
(azote) was supposed to be 5, and that of carbon also 5.

It is worth while to notice, however, that Dalton applied the
term atom to the ultimate particles of substances known to be
compounds as well as to those of elements; it is noteworthy
also that his numerical values furnish examples of the__law of
multiple proportions; for instance, the composition of the two
oxides "of "T^rbdnr""""""

Dal ton's system of atomic symbols was ingenious: (^ stood
for oxygen, Q for hydrogen, ^ for carbon, &c.; whilst for

compounds such formulae as /TV- * which stands for sulphuric

CFO

acid (S0 3 ), had to be constructed. In these formula, however,
picturesqueness did not compensate for practical inconvenience;
and the suggestion of Berzelius (1811), that initial letters should
replace Dalton's hieroglyphics, found general acceptance. 1

4. The Molecular Theory

Dalton made no further advance along the road that he had
traversed. His assumption that the simplest formulae for a com-
pound is the right one was a subterfuge which marked the end
of the road. Advance must therefore be sought in another
direction; and it is found in the study of gases; for gases are
the simplest form of matter, and, it* atoms exist, the properties of
gases will best elucidate their existence.

In 1805 Gay-Lussac and Humboldt studied the volume pro-
portions in which oxygen and hydrogen combine to form water;
and announced that " 100 volumes of oxygen required for com-
plete saturation 199-89 volumes of hydrogen, for which 200 may
be put without error'*. This is a single example of a law, Gay-
Lussac's law of volumes, which are thus expressed:

The volumes in which gases combine are simply related to each
other, and to the volume of the compound gas which is formed,, j

1 For a note on Chemical Nomenclature and Symbols see Appendix.

12 CHEMICAL THEORY

For example:

2 volumes of hydrogen combine with 1 volume of oxygen to form

2 volumes of steam.

1 volume of hydrogen combines with 1 volume of chlorine to form
2 volumes of hydrogen chloride.

3 volumes of hydrogen combine with 1 volume of nitrogen to form

2 volumes of ammonia.

A necessary corollary of this law is the statement that: the
densities, i.e. the masses of unit volumes, of the elementary gases
are simply related to their combining weights.

Thus, since 1 volume of hydrogen combines with 1 volume
of chlorine, and also 1 grm. of hydrogen combines with about
35-5 grm. of chlorine, the density of chlorine compared with
that of hydrogen as unity is about 35-5.

It further follows, if gases combine volume by volume, accord-
ing to the law of Gay-Lussac, and also atom by atom, according
to the theory of Dalton, that there is a simple connection between
the volume and the atom; and, indeed, that equal volumes of
hydrogen and chlorine, for example, contain equal numbers of
atoms. This conclusion, which was quite valid so far as it went,
was reached by Gay-Lussac, but was denied by Dalton, on account
of a difficulty which arose when the volume of the product was
considered.

Now when two separate and different elementary atoms com-
bine to form a compound atom, or whatever it may be called,
it is one entity they form, not two. It is impossible, for instance,
that 1 atom of hydrogen combining with 1 atom of chlorine can
produce two compound atoms of hydrogen chloride. And yet 1
volume of hydrogen combining with 1 volume of chlorine forms
2 volumes of hydrogen chloride. That was a dilemma; and it
was met by Dalton by a spirited denial of the law of Gay-Lussac.
"The truth is", said Dalton, "that gases do not combine in simple
proportions by volume; when they appear to do so, it is due to
an error in our experiments" !

Now, Dalton was wrong; and yet what other solution can
be found, unless indeed the "atoms" are torn in pieces in the
process of chemical synthesis, and the pieces are afterwards joined
together again in a different way?

That is precisely the solution of the difficulty suggested by

THE ATOMIC AND MOLECULAR THEORIES

Avogadro, in his celebrated hypothesis. In this hypothesis, which
will now be expounded, two orders of particles were distinguished,
which we now call atoms and molecules. Atoms are indivisible
in ordinary chemical changes; molecules are aggregates of atoms
with a few exceptions which maintain their integrity in
ordinary physical changes, but suffer disruption in the course
of chemical change, so that their constituent atoms may be re-
arranged to form fresh molecules.

Now, when hydrogen chloride is formed from its elements
the volume of the product is twice the volume of the hydrogen
or of the chlorine; therefore it is sufficient to assume that the
molecules of hydrogen and chlorine consist of pairs of atoms
which break into single atoms, and recombine, thus:

so that 1 volume of hydrogen plus 1 volume of chlorine gives
2 volumes of hydrogen chloride, instead of 1 volume, according
to the scheme:

It might be objected, however, that if the molecules of hydrogen
chloride are intrinsically twice the size of the atoms of hydrogen
and chlorine, out of which they are formed, the volume of the
compound gas might be expected in any case to be twice that of
either of the simple gases. Such an objection, however, is invalid,
since the actual size of the molecules of a gas is very small com-
pared with the molecular interspaces, and consequently the question
of a molecule of hydrogen chloride being intrinsically larger than
an atom of hydrogen or of chlorine does not arise.

The formation of 2 molecules of steam from 2 molecules of
hydrogen and 1 molecule of oxygen is thus represented:

CHEMICAL THEORY

The above processes of combination may be set forth in terms
of volumes, by using Dalton's symbols, thus:

1 vol. hydrogen. 1 vol. chlorine.

2 vols. hydrogen chloride.

2 vols. hydrogen

1 vol. oxygen.

2 vols. steam.

Or by means of chemical equations:

H 2 + C1 2 = 2 HC1.
2 H 2 + O 2 = 2 H 2 O.

Thus the molecular formula H 2 for water makes its appear-
ance. The proof of this formula is contained in the preceding
argument, which may be thus epitomized:

Hydrogen and chlorine gases consist of diatomic molecules,
since the volume of hydrogen chloride they produce is twice the
volume of either single gas.

Similarly, oxygen gas consists of diatomic molecules, since the
volume of the steam is twice the volume of the oxygen it contains.
The only formula for steam which agrees with the diatomicity
of hydrogen and oxygen, as well as with the volumetric com-
position of steam, is H 2 O. That the density of steam (H = 1)
is 9 furnishes no additional evidence, since it is deducible from
the densities of hydrogen and oxygen, and the volume of the
steam. That the atomic weight of oxygen is 16 follows from
the fact that its density is 16, and that, like hydrogen, it is
diatomic. The argument would, however, be invalidated if it
were shown that these gases are not diatomic, that in the mole-
cules H x and O x , x is greater than 2. Underlying the whole of
this argument is Avogadro's hypothesis, which is stated thus:

Equal volumes of all gases and vapours, under the same conditions
of temperature and pressure, contain equal numbers of molecules.

But why hypothesis? This statement is not a law, any more
than Dalton's atomic theory is a law. When first put forward
it was properly regarded as a hypothesis, which, indeed, suffered
much at the hands of its friends. Now, however, it is firmly
established, and is of fundamental importance. It ought, therefore,

THE ATOMIC AND MOLECULAR THEORIES

to be dignified with the name of theory. Henceforward we shall
speak of Avogadro'a theory.

It will be seen that this theory is in accord with Gay-Lussac's
law of volumes, and satisfactorily explains the phenomena of
the combination of gases. Thus, 1 volume of hydrogen combines
with 1 volume of chlorine to form 2 volumes of hydrogen chloride,
because 1 molecule of hydrogen reacts with 1 molecule of chlorine
to form 2 molecules of hydrogen chloride. The language of
volumes may be exchanged for the language of molecules; that
is the significance of Avogadro's theory.

That equal volumes of different gases contain equal numbers
of atoms is true only when the molecules of these gases contain
equal numbers of atoms. It is a statement of limited truth, and
of no permanent importance. The same may be said of the state-
ment that the densities of elementary gases are in the same ratio
as their atomic weights. The important fact is that the densities
of all gases are in the same ratio as their molecular weights',
and further, that since the molecular weight of hydrogen is 2,
and its density, which is taken as the standard, is 1, therefore
the molecular weights of all gases are twice their densities. Thus,
the molecular weight of a gas or vapour is revealed by its density,
as the following approximate figures show:

Elementary Gas or Vapour.

Density.

Molecular
Weight

Atomic
Weight.

Molecular
Formula.

Hydrogen

n 2

Oxygen
Nitrogen
Chlorine

16
14
35-5

28
71

16
14
35-5

N 2

Ozone

Phosphorus

124

Mercury

100

200

Sulphur

128

256

S 3

It may be remarked, incidentally, that the magnitude of the
atomic weight of an element cannot be deduced from its gas or
vapour density unless the number of atoms contained within
the molecule of the element, i.e. its atomicity, is known indepen-
dently. As a rule, however, the atomic weight of the element
is known independently, and then the atomicity is deduced from
the density.

16 CHEMICAL THEOKY

The breadth of Avogadro's generalization was not realized in
the time of its originator; and, owing to the persistence of the
volume-atom theory of Gay-Lussac, and its unwarrantable exten-
sion by Berzelius, 1 there was much confusion on the subject until
Cannizzaro, in 1858, reinstated Avogadro's theory on a permanent
basis.

It should be added that Avogadro's theory applies strictly
only to an ideal gas. When a gas deviates from Boyle's law it
deviates to the same extent from Avogadro's theory.

A useful fact to remember in connection with gas densities is
that a litre of hydrogen at C and 760 mm. pressure, i.e. normal
temperature and pressure (N.T.P.), weighs almost exactly 0-09 grin.,
or that 1 grm. measures 11-125 litres. Thus a gram-molecule (i.e.
the molecular weight in grams) of hydrogen at N.T.P. measures
22-25 litres; and from Avogadro's theory it follows that the
volume of a gram-molecule of any gas or vapour, reduced to
normal temperature and pressure, is 22-25 litres. To determine
the weight in grams of 22-25 litres of any gas or vapour, reduced
to and 760 mm., is therefore to discover its molecular weight
referred to H = 1. It is usual now, however, for a reason which
will appear later (p. 20), to accept the atomic weight and density
values H = 1-008 and O = 16-00, so that a gram molecule of any
gas or vapour at N.T.P. measures 224 litres.

SUMMARY

AN ELEMENT is a substance which hitherto has not been resolved
chemically into two or more dissimilar kinds of matter.

LAWS OF CHEMICAL COMBINATION. 1. Law of definite or fixed
proportions. The same chemical compound always contains the
same elements united together in the same proportions; or, the
proportions between the constituent elements of a chemical com-
pound are always the same.

2. Law of Multiple Proportions. When one element combines
with another in more than one proportion, these proportions bear
a simple ratio to one another.

3. Law of Reciprocal Proportions. The proportions of two

*The practice of referring all gaseous molecules to 2 volumes, which was a pernicious
outcome of the theorizing of Berzelius, appears now, fortunately, to be dying out. Why,
indeed, should every molecule be regarded as a microcosm of 2 volumes, as if it could
necessarily be dichotomized?

THE ATOMIC AND MOLECULAR THEORIES 17

elements which separately combine with a fixed proportion of a
third element are also the proportions of these elements which
combine with each other, or else in accordance with the law of
multiple proportions they bear a simple ratio to these proportions.
THE ATOMIC THEORY. 1. All matter consists of discrete
particles called atoms, which remain unbroken throughout chemical
change.

2. Atoms of the same element are ordinarily supposed to be
similar in all respects.

3. Chemical compounds are formed by the union of the atoms
of different elements in simple numerical proportions.

4. The proportions in which elements combine to form com-
pounds are determined by the atomic weights of the elements.

GAY-LUSSAC'S LAW OF VOLUMES. The volumes in which gases
combine are simply related to each other, and to the volume
of the compound gas which is formed.

Corollary. The densities of the elementary gases are simply
related to their combining weights. 1

AVOGADRO'S THEORY. Equal volumes of all gases and vapours
under the same conditions of temperature and pressure contain
equal numbers of molecules.

Corollary.- Since the molecule of hydrogen contains 2 atoms,
the molecular weight of any gas or vapour is twice its density
compared with that of hydrogen.

A litre of hydrogen at N.T.P. weighs 0-09 grm., and 1 gram-
molecule of hydrogen (2-016 grm.) measures 22-4 litres. It follows
from Avogadro's theory that this is also the volume at N.T.P.
of 1 gram-molecule of any gas or vapour.

AN ATOM of an element is the smallest particle of matter
which takes part in a chemical change; it is the unit of chemical
exchange.

A MOLECULE is the smallest particle of matter which exists
independently; it is the physical unit. The molecule of an element
contains similar, that of a compound dissimilar atoms.

The number of atoms contained within the molecule of an
element is called the atomicity of the element.

'The term "combining weight" has sometimes signified equivalent weight, and some-
times atomic weight. Since the term is ambiguous, a use is found for it during the
development of the molecular theory when non-committal language is employed. After-
wards the term should be dropped.

(1)60) 3

CHAPTER II
EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS

i. Equivalent and Atomic Weights

It was shown in the last chapter that Dalton's "atomic weights"
were really equivalent weights, and that the equivalent weight of
an element, when not identical with its atomic weight, is a sub-
multiple of the latter. Thus, whilst the equivalent weight of
oxygen referred to that of hydrogen as unity is approximately 8,
the atomic weight of this element, referred to the same standard, is
approximately 16. In general

Atomic weight = n x equivalent weight,

where n is a small whole number, which indicates the valency of
the element. Valency, or atomic value, is a new idea, necessary to
connect together the ideas of atomic weight and equivalent weight.
It will be more fully developed later.

It will now be useful to define equivalent and atomic weights.

EQUIVALENT WEIGHT. The equivalent weight of an element is
that weight of it which combines with, or displaces from combina-
tion, an agreed weight of a standard element.

ATOMIC WEIGHT. The atomic weight of an element is the ratio
between the weight of its atom and that of the atom of a standard
element.

When these definitions are considered, it appears that the
equivalent weight of an element is an experimental value, inde-
pendent of theory, whilst the atomic weight is connected with
the atomic theory.

It further appears that since equivalent and atomic weights
are ratios, they are not really weights at all, nor masses, but pure
numbers. That the atomic weight of an element is not the weight
of one of its atoms appears plainly enough when it is considered
that the standard of atomic weights has varied from time to time.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 19

Farther, since equivalent weights are values to be determined
experimentally, their determination may well form the starting-
point in the estimation of atomic weights. As a matter of fact the
accuracy with which the atomic weight of an element is known
depends as a rule on the accuracy with which the quantitative
observation of some chemical transformation has been carried out,
so as to determine its equivalent weight.

In some cases, however, atomic weights have been estimated
accurately by the determination of gas density.

For determining equivalents, comparison between reacting
quantities may be made by combination as well as by displace-
ment, because an element combines with, as well as displaces, what
is equivalent to itself. Thus, if there are two elements, A and B,
the chemical equivalent of B referred to A as standard is found by
estimating the amount of B which combines with a known weight
of A, as well as by causing B to displace A, or A to displace B from
combination with another element or group of elements.

When the equivalent weight of an element is known, it is
necessary to determine the value of n in the above equation before
the atomic weight can be fixed. What multiple of the equivalent
weight the atomic weight may be, has to be decided by reference to
one or more of several distinct principles, which lie chiefly in the
domain of physical chemistry, and will shortly be discussed in
detail.

Standard for Equivalent and Atomic Weights.

The question of a standard needs first to be considered; and,
since hydrogen has the least atomic weight of all the elements, and
as small an atomic value (valency) as any element, it is natural to
choose hydrogen as the standard both of atomic and equivalent
weights, and so to make its equivalent and its atomic weight both
equal to 1.

Now, although hydrogen combines with non-metals, and a few
metals, and is displaced from its combination in acids by some
metals, its chemical activity is too limited to permit its use as
a general standard of comparison. Oxygen, however, with very
few exceptions, combines with all the elements, metals and non-
metals alike; on this account it was called by Berzelius the "pole
of chemistry". As a matter of practical experience, therefore,
equivalent and atomic weights are more often estimated with

20 CHEMICAL THEORY

reference to oxygen than to hydrogen; the hydrogen equivalent
may then be calculated from the oxygen equivalent by multiplying
the latter by the equivalent weight of oxygen, and thence the
corresponding atomic weight may be found.

Now, although Dalton (1808) chose hydrogen = 1 as the atomic
weight standard, oxygen was soon adopted in preference, so that
Wollaston (1814) used oxygen = 10, Thomson (1825) oxygen
= 1, Berzelius (1830) oxygen = 100, and Stas (1860-5) oxygon
= 16.

During a recent period the two standards H = 1 and O = 16
were in use, but the latter is now the standard adopted by the
International Union of Pure and Applied Chemistry. Although
unity as the standard is sacrificed by this procedure, the = 16
has at least two advantages over the H = 1 standard.

It was pointed out by Stas that the standard atomic weight
should, as far as possible, be directly connected with the atomic
weight to be determined, and this is the case when oxygen rather
than hydrogen furnishes the standard. Otherwise the ratio H : O
is involved in the calculation when the data are derived from the
composition of an oxide; and whilst this ratio has been determined
with great accuracy to be 1 : 15-88, any future modification of the
ratio would involve the recalculation of all atomic weights de-
pendent upon it. If, however, the ratio is written 1-008 : 16, the
atomic weight of oxygen being fixed at 16, any future alteration
will involve only the atomic weight of hydrogen. The advantage
of this is plain.

Another advantage of the modern system is the fact that
when = 16 several other important atomic weights approximate
very closely to whole numbers; e.g. C = 12-00, N = 14-01,
Na = 22*997. The reason for this approximation will appear
later when modern views of the atom are considered.

It is unlikely that any further modification of the standard
will now be made, but an unfortunate confusion even now remains
in the minds of those who have employed several standards. For
example, the atomic weight of chlorine has been variously given as
35-37, 35-18, 35-46; and these differences are due not to different
estimations of the atomic weight of this element but to the adop-
tion of three different standards for oxygen, viz. = 15-96
(Dumas), 15-88, and 16-00.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 21

2. Determination of Equivalent Weights

The following are the more important methods ordinarily
employed in the laboratory to determine the equivalent weights of
elements.

i. The measurement of the volume of hydrogen displaced from
dilute sulphuric or hydrochloric acid by a weighed amount of
a metal.

ii. The conversion of a weighed quantity of a metal into its
oxide which is weighed, or the reduction of a weighed quantity of
oxide to metal.

iii. The displacement of a metal from a solution of one of its
salts by a weighed quantity of a more chemically powerful metal.

iv. The separation of elements at the electrodes during the
passage of an electric current through a series of electrolytes.
This method yields the electro-chemical equivalent of an element;
but this value is numerically identical with the chemical equivalent.

i. The chemical equivalent of magnesium, zinc, or aluminium
may be easily determined by dissolving a weighed quantity of the
metal in the dilute acid contained in a piece of apparatus designed
for collecting the evolved hydrogen. The gas is measured over
water at atmospheric temperature and pressure; it will consequently
be moist, and the pressure of water vapour at the observed tempera-
ture must be subtracted from the atmospheric pressure, before the
volume of the gas is corrected to normal temperature and pressure.

The weight of metal divided by the weight of the evolved
hydrogen gives the hydrogen equivalent of the metal. This must
be multiplied by 1-008 if the equivalent on the modern atomic
weight basis is desired; though in view of the likely experimental
error such a correction is superfluous.

The experiment may easily be carried out on the lecture-table or
by students. The following result has been obtained by a student:

Weight of magnesium taken = 0-033 grm.

Volume of moist hydrogen measured at\ _. r n

1C C. and 756 mm. / "~ '

Pressure of water vapour at 12 = 10* 5 mm.

Volu-ne of dry hydrogen at N.T.P. - ^^\~^^

= 30-6 c. c.
Weight of hydrogen = 30-6 X 0-00009 = 0*002754

Equivalent of magnesium = A l???, = 12-0.

- ^f^u^m

22 CHEMICAL THEORY

ii. Magnesium may be converted quantitatively into oxide by
the ignition of the metal in the air under suitable conditions, or
by dissolving it in dilute nitric acid, evaporating the solution, and
igniting the nitrate until brown fumes cease to be evolved. These
methods are not without sources of error, but it may be shown that
0-30 grm. of magnesium yields almost exactly 050 grm. of oxide, so
that the equivalent weight of magnesium, that of oxygen being 8, is

The method of conversion into oxide through the nitrate is
applicable to such metals as zinc and copper, which dissolve in
nitric acid and yield stable oxides by the decomposition of their
nitrates. The equivalent of tin may be determined by the con-
version of the metal into hydrated dioxide by means of nitric acid,
since the ignition of the product yields the pure dioxide.

It would be possible to determine the equivalent of carbon by
burning a weighed quantity of the element in a stream of dry air
or oxygen, and collecting and weighing the carbon dioxide formed;
but the great difficulty of obtaining pure carbon free from hydrogen
under ordinary conditions stands in the way of this determination.

For the determination of an equivalent by the reduction of an
oxide to metal, copper furnishes the usual example, since the
reduction is easily carried out by passing a stream of hydrogen
over oxide of copper contained in a boat in a heated tube. Thus
1-00 grm. of black oxide of copper leaves a residue of 0*799 grm. of
copper; whence the equivalent of copper in this oxide is

x 8 " 31 ' 8 '

There is another oxide of copper, however, the red oxide, whose
equivalent weight is 31-8 x 2 = 63-6. This fact is connected with
the exhibition of a dual valency by copper, which again furnishes
an example of the law of multiple proportions. This phenomenon
will be further dealt with under the subject of valency.

iii. A well-known example of the displacement of a metal
from the solution of one of its salts by another metal is the action
of zinc upon a solution of copper sulphate, when the zinc is sup-
posed to displace from combination its equivalent of copper which
may be collected and weighed. This takes place almost quantita-
tively when a cold concentrated solution of copper sulphate is

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 23

employed; but the method is generally unreliable, because other
reactions occur between the displacing metal and the solution
simultaneously with the main reaction, and these vitiate the results.
Therefore the method is not to be recommended.

iv. When a suitable electric current is passed through acidified
water contained in a "voltameter", hydrogen and oxygen are

evolved in the proportion of two
volumes of hydrogen to one
volume of oxygen. Provided the
densities of hydrogen and oxygen
are known, and the conditions
of temperature under which the
gases were measured have been
observed, the hydrogen equiva-
lent of oxygen might be calcu-
lated from the volume relations

Oxygen
008 g

Hydrogen
01003

Fig. 1

of the gases. The estimation would not, however, be very accurate,
owing to several sources of experimental error.

If, however, the same current passes in succession through
several salt solutions for example, copper sulphate, silver nitrate,
gold chloride solutions it will liberate at the cathodes or negative
electrodes amounts of the metals chemically equivalent to the
hydrogen which is liberated in the voltameter. Thus, whilst
0-01008 grm. of hydrogen gas is being evolved, and 0-08 grin, of
oxygen, 0-318 grm. of copper, 1-079 grm. of silver, and 0-657 grm.

24 CHEMICAL THEORY

of gold will be deposited in the successive electrolytic cells. The
necessary arrangement is shown in fig. 1. Thus the equivalent
weights of these metals are determined.

*$. Determination of Atomic Weights

It has been suggested in the previous pages that two distinct
considerations have to be taken into account in the problem of
atomic weight determination. These are:

i. An exact estimation of the chemical equivalent of the
element must be made, generally by carrying out some suitable
chemical transformation, occasionally by other means.

ii. A decision must be arrived at as to the order of magnitude
of the atomic weight, so as to discover the small whole number by
which the equivalent weight must be multiplied to give the atomic
weight.

The order in which the two parts of the problem are here
placed is that which would naturally occur to the mind. Never-
theless it is not the order of historic sequence in relation to modern
atomic weights. The approximate magnitudes of the atomic weights
of all the elements have long since been settled and are not discussed
in modern research upon atomic weights; but the determination of
the exact values of all these atomic weights is a laborious task
which is not yet completed.

The methods for determining chemical equivalents which have
been described above are suitable for demonstration purposes, but
not all of them are equally useful in the actual determination of
atomic weights. Illustrations of the methods that have been
employed in accurate atomic-weight determinations will be given
in the sequel.

The principles which have led to decisions upon the order of
magnitude of the atomic weights of the elements will now be
dealt with.

It has already been seen that Dalton was in need of some
guiding principle to enable him to fix the magnitude of his atomic
weights; and that such a principle came to light in the discovery
by Gay-Lussac of the law of gaseous volumes, and the proper
interpretation of this law by Avogadro. Thus, by means of
Avogadro's theory it was shown that the atomic weight of oxygen
is very probably 16 and not 8; but clearly this theory is limited in

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 25

its application, since it can only be of use in the case of gaseous
or gasifiable substances. Here may be mentioned the method of
chemical displacement, which is of some value in deciding the
magnitude of atomic weights.

In 1819 two other and quite distinct principles became avail-
able in the law of specific heats of Dulong and Petit, and the law
of isomorphism discovered by Mitscherlich. These laws are espe-
cially valuable in furnishing guidance as to the magnitude of
atomic weights, because they are applicable to solid elements and
their solid compounds.

The former of these two laws is the more important, and has
the wider application. Finally, the periodic law, established by
Mendel^eff in 1869, has been of distinct value in several ways in
fixing the approximate magnitude of atomic weights.

So the five guiding principles that aid in settling the order of
magnitude of atomic weights are:

i. Avogadro's theory,
ii. Chemical displacement,
iii. Dulong and Petit's law of specific heats,
iv. Mitscherlich's law of isomorphism,
v. Mendele'eff's periodic law. y/

i. The Method of Avogadro's Theory.

It will be remembered that according to Avogadro's theory the
molecular weights, not the atomic weights, of gases and vapours are
proportional to their densities. It follows, therefore, that the rela-
tive magnitudes of molecular weights, and not of atomic weights,
are directly deducible from Avogadro's theory. So the question
arises how far a knowledge of the relative weights of molecules
can be of use in fixing the relative weights of any of their con-
stituent atoms. Such knowledge may be employed in two ways.

Consider the following volatile hydrocarbons:

Methane.

Ethylene.

Propane.

Benzene.

Naphthalene.

Approximate Density)
(0 = 16) /

Approximate Molecular \
Height f

128

Molecular Proportion of \
Carbon /

120

26 CHEMICAL THEORY

Approximate estimations of gas or vapour density yield ap-
proximate molecular weights; whilst quantitative analysis shows
the proportion of carbon within the molecular proportion of each
compound. Now, it is evident that all these hydrocarbons, except
the first, contain more than 1 atom of carbon in their molecules.
The molecule of methane might indeed contain more than 1 atom,
though the fact than no submultiple of 12 appears in the pro-
portions of carbon in the other molecules is evidence, so far as it
goes, that the figure 12 represents an indivisible unit, or in other
words that 12 is approximately the atomic weight of carbon. And
since by the examination of the very large number of hydrocarbons
that exist, every molecular proportion has been found to contain
12, or a multiple of 12 parts of carbon, the probability that 12 is
the atomic weight of carbon reaches a practical certainty.

/The principle thus illustrated may be put in the following words:
/ The least proportion of an element found within the molecular propor-
tion of any of its volatile compounds is likely to be the atomic weight of
the element; and if the number of compounds which have been examined
is large, the value indicated is very probably the atomic weight.

The question may be asked, however, whether atomic weights
can be determined exactly by the method of Avogadro's theory, i.e.
by the determination of gas density, and the answer is in the
affirmative, provided an ideal gas density can be determined and
the molecular composition of the gas is known.

The density of a gas is determined by weighing a large glass
globe of about 10 litres capacity, first evacuated, and then filled
with the gas, at known temperature and pressure, corrections
being applied for the air displaced by the globe, and for the
slight shrinkage which the glass undergoes when the globe is
evacuated. Thus it has been found, as the mean result of the
experiments of Eayleigh, Morley, and Leduc, that 1 litre of
oxygen at and 760 mm. at the latitude of Paris weighs 1-42895
grm., whilst 1 litre of hydrogen, under similar conditions, accord-
ing to the experiments of Morley and Leduc, weighs 0-08985 grm.

To conclude, however, that the atomic weights of oxygen and

hydrogen are in the ratio ' J*~, although we know that the

(70*700

molecules of both gases are diatomic, would be erroneous, because it
would be to assume that the gases are ideal gases which behave
in perfect accord with the gas laws (q.v.), and so with Avogadro's

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 27

theory. Such, however, is not the case, and the deviation of these
gases from the ideal must be discovered, and allowed for. This
may best be done, in the present case, by determining the effect
of the deviation upon the volume relations in which oxygen and
hydrogen combine to form water.

Now it has been estimated that 2 00268 litres of hydrogen
combine with 1 litre of oxygen, at and 760 mm. at the latitude
of Paris. This complex ratio is due, not to any discrepancy
between the simple proportions in which the molecules of these
two gases interact, but to the > fact that equal volumes do not
contain quite equal numbers of molecules because oxygen is a
little more compressible than hydrogen. But since the densities
relate to equal volumes it may be concluded that

2-00268 x 0-08985 grm.

of hydrogen combine with 1-42895 grm. of oxygen, and therefore
that the hydrogen equivalent of oxygen is

1-42895 ^

2-00268 x 0-08985 '

and its atomic weight 15-88 when H = 1; so that H = 1-0076
when O = 16. A similar method may be applied to determine
the atomic weight of a constituent element of a compound gas.
Thus, by the calculation of the ideal densities of carbon monoxide,
carbon dioxide, methane, and acetylene, by applying a correction
for compressibility to the estimated densities, several observers
have accurately determined the molecular weights of these gases,
and thence the atomic weight of carbon. ^

METHODS OF DETERMINING VAPOUR DENSITY. .*

The determination of gas density always consists in weighing a
certain volume of the gas; but for determining the vapour density
of a volatile liquid or solid an alternative procedure may be
adopted: the volume of the vapour produced by a weighed quantity
of the liquid or solid may be measured under known conditions.

There are three well -recognized methods of vapour density
determination: the methods of Dumas, Hofmann, and Victor
Meyer. In the first of these three methods the weight of a
known volume of the vapour is ascertained; in the two latter the
volume of a weighed quantity of the substance is measured. The
method of Victor Meyer is the easiest and most often employed.

28 CHEMICAL THEORY

(a) Dumas's Method of Vapour-density Determination

A glass globe of the shape shown in fig. 2, and capable of
holding from 50 to 100 c. c. or more, is weighed, and then filled
with the vapour of the substance in the following manner.

A few cubic centimetres of the liquid are introduced into the
globe, which is then immersed in a bath of another liquid whose
temperature is kept constant, and from 20 to 30 above the
boiling-point of the liquid in the globe. As the latter liquid
boils it displaces the air from the globe, and vapour issues from
the neck as long as any liquid remains within the globe. When
the stream of vapour ceases, the globe is filled with
the vapour at atmospheric pressure, and at the tem-
perature of the bath in which it is immersed. The
neck is then sealed by means of a blowpipe; and
the temperature of the bath, and the pressure of
the atmosphere at the time of sealing are recorded.
After being cleansed, the sealed globe is weighed,
and the temperature and pressure of the air in the
vicinity of the balance are also observed.
Since the true weight of the sealed globe with its contents
is equal to its apparent weight plus the weight of the air which
it displaces whilst it is being weighed, the weight of this air must
be calculated and added to the apparent weight. For this calcu-
lation, as well as to ascertain the volume of the vapour at the
time of sealing, the cubical capacity of the globe must be de-
termined. This is done by breaking off the end of the neck of
the globe under water, which should then enter and fill the globe.
The quantity of water in the globe is determined by another
weighing, the weight of the air displaced being in this case
negligible; then the weight of the water in grams shows the
volume of the globe in cubic centimetres with sufficient accuracy.
From these data the weight of the known volume of the vapour
contained by the globe at the temperature and pressure at which
it was sealed is calculated. The volume is then reduced to N.T.P.,
and the weight of hydrogen or air corresponding to it is calculated.
The ratio of the weight of the vapour to that of the hydrogen
is the vapour density of the substance.

EXAMPLE. Calculate the density of ether vapour from the
following data:

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 29

Weight of open globe in air = 22 -649 grm.

Temperature of bath at time of sealing = 60 C.
Atmospheric pressure at times of sealing and weighing = 760 mm.

Apparent weight of sealed globe + vapour in air = 22*662 grm.

Temperature of air at time of weighing =15

Capacity of globe, indicated by weight of water it\ _ 75 c c

can contain J

Weight of 1 c. c. of air at C. and 760 mm. = 0-001293 grin.

Weight of 1 c. c. of hydrogen at C. and 760 mm. = 0-0000899 grm.

Calculation

Weight of air displaced when sealed globe is weighed

_ 0-001293 X 76 X 273

288

0-092 grm.

Weight of vapour in globe = 22 662 + 092 22 549 = 205 grm.
Weight of an equal volume of hydrogen at 60 C. and 760 mm.

_ 0-0000899 X 75 X273

So density of ether ((C 2 H 6 ) 2 O) vapour

333
0-205 = 37.!.

0-00553

(b) Hofmanns Method of Vapour-density Determination
A weighed quantity (about
0-05 grm.) of the liquid con-
tained in a small, drawn-out
bulb or stoppered bottle is in-
troduced into the Toricellian
vacuum of a graduated barom-
eter tube surrounded by the
vapour of a liquid boiling at
a suitable temperature, which
may be below the boiling-point
of the liquid whose vapour den-
sity is being determined. As
the liquid is vaporized it de-
presses the mercury in the
barometer tube; and when the
volume has become constant it
is read off, and the temperature
of the vapour jacket is observed.
The pressure of the vapour is
equal to atmospheric pressure
less the height of the mercury
in the tube above its level in
the vessel in which the tube stands. Strictly speaking, the height of

To Condenser

Fig. 3

CHEMICAL THEORY

the mercury column should be corrected for expansion by heat; but
this need not be considered. From these data the vapour density
of the liquid can be calculated; as the following example shows:

Weight of stannic chloride (B.P. 114 ) 1 taken
Volume of vapour
Temperature of vapour jacket
Barometric pressure
Height of mercury column
Whence pressure of vapour

Volume of vapour reduced to N.T.P.
Weight of 3-75 c.c. of hydrogen at KT.P.

Vapour density of stannic chloride (SnCl 4 ) =

0*0445 grm.
16-2 c.c,
99

752 mm.
512 mm.
240 mm.

16*2 X 273 X 240 _ ?
372 X 760 3-75 c.c.

3-75 X 0-00009 grm.
0-0003375 grm.

0-0445
0-0003375

= 131-8.

: E

-i

7fo

In this method a weighed quantity of the
substance is made to evaporate into a space sur-
rounded with the vapour of a boiling liquid whose
boiling-point is at least 25 higher than that of
the substance. The volume of the vapour is not
directly measured, but the air displaced by it is
collected and measured at atmospheric tempera-
ture and pressure; while all the displacing vapour
remains in the locality of its production. The
weight of an equal volume of hydrogen is then
calculated, and the weight of substance taken
divided by this weight of hydrogen gives the
vapour density of the substance, since the vapour
of the substance, if it could be obtained at atmo-
spheric temperature and pressure without conden-
sation, would occupy the same volume as the air.
The tube A (fig. 4) is closed at the lower end, and
is furnished with a bent delivery tube B which
dips under water in the dish C. The upper end
of A is closed by a rubber stopper. The lower
part of the tube is heated by the vapour of a
' D liquid, e.g. water, boiling in the outer jacket D,
and, owing to expansion, air escapes by the side

1 Since the compound is vaporized into a vacuum the temperature
of its vapour may be lower than the B.P. of the compound.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 31

tube through the water. When no more bubbles of air are seen,
the graduated tube E is placed over the end of B, and a little tube
or loosely-stoppered bottle, containing a weighed quantity of the
substance under investigation, is dropped to the bottom of A, being
received on a pad of asbestos or glass wool, which prevents frac-
ture. For the introduction of the little vessel containing the sub-
stance the rubber stopper is removed, and quickly replaced, or the
vessel may be held by a mechanical contrivance at the top of the
tube A, where the air is nearly cold, 1 and then allowed to fall at
the right moment, without opening the tube.

EXAMPLE. 0444 grm. of chloroform displaced 28-6 c. c. of
moist air measured at 14 and 756 mm. pressure. Pressure of
water vapour at 14 = 12 mm.

Vol. of air at N.T.P. = 28-6 X 273 x (756-12) = 26>6 c c

287 X 7oO

Weightof an equal vol. of hydrogen = 266 X 0-0000899 grm.

= 0-00239 grm.

Vapour density of chloroform "I 0*144 __ fin o
CHC1 3 J 0-00239 ZZ2,

The method of Victor Meyer is more easily carried out than
either of the other methods. It employs very little of the substance
and is sufficiently accurate for most purposes. Consequently, it
is the method usually employed.

*It might be supposed that since the air is colder in the upper part of the tube,
which extends beyond the vapour jacket, than in the lower part, which is within it, too
much air will be displaced, and a high result obtained. This, however, is not the case,
because of the contraction of the air that rises in the body of the tube to take the place
of the air driven out. The following proof of this statement has been given by Dr. E. B. R.
Prideaux.

First, suppose temperature constant in the V. Meyer tube, so that the heated vapour
immediately displaces its own volume of heated air, which is then cooled. Let T = abs.
temperature of vapour and air when first expelled, and T abs. temperature of cooled air
leaving the end of delivery tube under water ; let V = vol. of vapour formed and therefore

T V
of air expelled, and V vol. of air collected. Then V = ~^-.

Second, let there be two temperatures T and Tj within the tube, with corresponding
volumes of equal masses of air V and Vj. f

T V
Then let V be expelled into the T! region, and thereby become V b so that Vj = -JL..

Vi, not V, will now displace its own volume of air, which will be cooled so as to become,

TV T T V TV

say, V 2 at the end of the delivery tube. Then V 2 = jl = ^ = Ml = V ; and

similarly with any number of temperature zones.

Thus a temperature gradient within the V. Meyer tube docs not affect the volume of
air displaced from the end of the delivery tube.

32 CHEMICAL THEORY

ii, The Method of Chemical Displacement.

Somewhat related to the above principle is another by which
the molecular formula of a compound may be determined, and so
the atomic weight of a constituent element.

Consider methane. The hydrogen in this compound can be
displaced by chlorine in four distinct stages, the following sub-
stitution products being formed: methyl chloride, methylene
chloride, chloroform, and carbon tetrachloride. The carbon, how-
ever, cannot be displaced fractionally. From these facts the
inference is drawn that the molecule of methane contains 4
atoms of hydrogen and only 1 atom of carbon; but if methane
is CH 4 , Dalton's problem of the number of atoms in the molecule
is solved, and the atomic weight of carbon is 12.

A similar argument may be applied to water. The composition
of sodium hydroxide proves that half the hydrogen of the water
molecule has been displaced by sodium. By no means, however,
can any fraction of the oxygen of the water molecule be displaced.
Thence it is concluded that water is H 2 O, and that the atomic
weight of oxygen is 16.

The principle of this method of fixing the magnitude of atomic
weights may be stated thus:

When -th of the proportion of a constituent element in a chemical
n

compound can be displaced by another element, a molecule of the
compound contains at least n atoms of that element.

iii. The Method of Dulong and Petit's Law.

The specific heat of a substance is the ratio of the amoant of
heat required to raise unit weight of it through one degree of
temperature to the amount of heat required to raise unit weight
of a standard substance through the same temperature interval.
The standard substance is water.

In 1819 Dulong and Petit published the specific heats of
thirteen elements, and showed that the product of specific heat
into atomic weight is approximately a constant quantity, the
average of which, on our modern atomic weight basis, is 6-4.
In the following table, containing the elements studied by Dulong
and Petit, modern values are given throughout.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 33

Element.

Specific Heat.

Atomic Weight.

Atomic Weight
x Specific Heat
=3 Atomic Heat.

Bismuth

0-0305

208-0

6-34

Lead ...

0-0315

207 2

6-53

Gold ...

0-03035

197-2

5-99

Platinum

0-03147

195-2

6-14

Tin ..

0-0559

119-0

6-65

Silver ..

0-0559

107-88

6-03

Zinc

0-0939

65-37

6-14

Telluriun

0-0475

127-5

6-06

Copper..
Nickel ..

0-09232
0-10842

63-57
58-68

5-81
6-16

Iron

0-10983

55-84

6-13

Cobalt ..

0-10303

58-97

6-08

Sulphur

0-1712

32-07

5-49

The law of Dulong and Petit may therefore be stated thus:

The specific heats of the solid elements are inversely proportional
to their atomic weights.

The product of specific heat and atomic weight, which is ap-
proximately a constant, is called the atomic heat because it is
the heat capacity of a quantity of an element proportional to
its atomic weight. Thus, for example, 55-84 parts by weight
of iron require the same amount of heat to raise them through
one degree of temperature as, say, 208-0 parts of bismuth. But
these quantities of the elements contain equal numbers of atoms.
So, in the words of Dulong and Petit, " the atoms of all substances
have exactly the same capacity for heat".

In order to reach this result, however, Dulong and Petit made
some drastic changes in the accepted atomic weight values, which
aroused the opposition of Berzelius, their author. Thus, taking
the atomic weight of sulphur as a true magnitude, they halved
the atomic weights of the metals in relation thereto. This pro-
cedure was, however, justified, even in the opinion of Berzelius,
after Mitscherlich, his pupil, had arrived at similar conclusions
by an application of the law of isomorphism.

Now, since

Specific heat x atomic weight = 6*4 (approx.)
atomic weight 6

specifieat

here is a valuable method for fixing the magnitude of the atomic

(D60) 4

34 CHEMICAL THEORY

weight of an element. All that it is necessary to do is to deter-
mine the specific heat of the element, and divide 6*4 by this value.

It must be clearly understood, however, that the value thus
obtained is only* approximate, for the atomic heat value, 6-4, is
only approximate, since it is a mean value, even if the specific heat
is accurately known. The method serves to indicate what multiple
of an accurately determined equivalent weight is the atomic weight.
To divide 6-4 by the given specific heat of an element, and report
the quotient as its atomic weight, is a gross error.

The following illustration will make plain the use of Dulong
and Petit's law: Marignac 1 found that 100 grm. of lead yielded
134-201 grm. of the chloride. The specific heat of the metal is
0-0315; find its atomic weight; Cl = 35-46.

The equivalent weight of lead is found from the proportion:

Wt. of chlorine : wt. of lead : : equivalent Cl : equivalent Pb,
so 34-201 : 100 :: 35-46 : 103-68.

The approximate atomic weight of lead, as indicated by its
specific heat, is:

Therefore the atomic weight of lead is twice its equivalent weight;
so Pb = 103-68 X 2 = 207-36.

Dulong and Petit's law applies strictly only to solid elements,
generally metals, whose atomic weights exceed 30. The specific
heats of other solid elements vary with temperature, but become
approximately constant at high temperatures, when they give an
atomic heat value of about 5-5.

iv. The Method of the Law of Isomorphism.

Isomorphism is similarity of crystalline form. It was supposed
by the earlier mineralogists that identity of crystalline form
generally indicated identity of chemical composition; but it was
shown by Mitscherlich in 1819 that compounds of analogous as
well as identical composition crystallize in similar forms belonging
to the same crystal systems. Thus sodium di-hydrogen phosphate
and sodium di-hydrogen arsenate, which are now represented by
the formulae NaH 2 PO 4 -H 2 O and NaH 2 As0 4 -H 2 O, were found
to be isomorphous. Careful measurements of the crystal angles
of isomorphous salts show that these angles are not quite equal,

1 Marignac, (Euvres Complies, 1846, I, 186.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 35

but the following criteria of isomorphism have been established:
(i) great similarity of crystalline form, (ii) analogous composition,
(iii) power to form mixed crystals by simultaneous crystallization,
(iv) power of crystal overgrowth, so that a crystal of one com-
pound may form the matrix on which the growth of the crystal
may be continued by the deposition of another substance.

With regard to the second criterion, it must be noted that
isomorphism is sometimes observed in pairs of compounds which
are not chemically analogous, but have the same numbers of atoms
within their molecules. Thus calc-spar (CaCO s ) is isomorphous
with Chili saltpetre (NaNO 3 ), and aragonite (CaCO 8 ) with nitre
(KNO S ).

Mitscherlich stated the law of isomorphism as follows:

"The same number of atoms combined in the same manner
produce the same crystalline form; the crystalline form is independent
of the chemical nature of the atoms, and is determined solely by
their number and mode of combination."

Nevertheless isomorphism such as that illustrated by the case
of sodium phosphate and arsenate is the rule; that is to say, not
only do the molecules of isomorphous compounds contain the same
number of atoms similarly combined, but these atoms themselves
are analogous, as, for instance, are phosphorus and arsenic. Indeed,
isomorphism is taken to be a sign of chemical analogy.

Therefore, for practical purposes, the law of isomorphism may
be stated more briefly:

The molecules of isomorphous substances contain equal numbers
of atoms, which when not of identical are of analogous elements.

The consequence of this law, when applied to the case already
mentioned, is that the atomic weights of phosphorus and arsenic
can be directly compared, and if one atomic weight is known the
other is derivable from the results of chemical analysis.

A simple numerical example is furnished by the following
results of the analysis of the isomorphous salts potassium sulphate
and potassium selenate, carried out by Mitscherlich:

K 2 SO 4 K 2 Se0 4

100 parts contain 100 parts contain 127-01 parts contain

K 44-83 K 35-29 44-83

O 36-78 O 28-96 36-78

S 18-39 Se 35-75 45-40

100-00 100-00 127-01

36 CHEMICAL THEORY

In the third column is shown the proportion of selenium in an
amount of the selenate which contains the same amounts of potas-
sium and oxygen as are shown in the percentage analysis of the
sulphate. Whence it follows that 45-40 parts of selenium take
the place of 18-39 parts of sulphur.

Now, the law of isomorphism declares that the ratio between
these quantities is the ratio between the atomic weights of the two
elements. Therefore, if the atomic weight of sulphur is 32-0 that

of selenium is S 2 ' ^ 5 ' 40 = 79^.

The phenomena of isomorphism are somewhat confused by those
of dimorphism and polymorphism. Thus calcium carbonate, as
shown above, is dimorphous in calc-spar and aragonite; ammonium
nitrate, NH 4 N0 3 , is tetramorphous, crystallizing in four distinct
forms; arsenious and antimonious oxides, As 4 O 6 and Sb 4 O 6 are
isodimorphous, that is to say, they are both similarly dimorphous.
Nevertheless, the phenomena of isomorphism have been of value,
not only in confirming atomic -weight magnitudes derived from
other considerations, but in correcting erroneous magnitudes.

For example, previous to the recognition of isomorphism,
Berzelius regarded various metallic monoxides MO as dioxides
MO 2 ; similarly, Fe 2 O 3 was written FeO 3 , Cr 2 O 3 was pr0 3 , and Cr0 3
was CrO 6 . But when this chemist recognized the isomorphism of
chromates with sulphates he altered CrO 6 to CrO 3 to agree with
S0 3 , the oxide known to be present in sulphates. Consequently,
the former CrO 3 became Cr 2 3 ; and since chromic and ferric alums
were isomorphous, what was formerly Fe0 3 became Fe 2 8 , and so
FeO 2 became FeO. But compounds of copper, nickel, cobalt, man-
ganese, zinc, and magnesium are isomorphous with corresponding
iron compounds, and so if FeO 2 should be FeO the corresponding
dioxides of all these metals should really be monoxides.

This sweeping change would involve the halving of a number
of accepted atomic weights; nevertheless, the change was made
by Berzelius in accordance with the principles of isomorphism; and
it was at once ratified by the law of specific heat, which required
the same atomic weight magnitudes for the elements concerned.

The alums which conform to the general formula

M 2 SO 4 -X 2 (S0 4 ) 3 -24H a O
are amongst the best -known isomorphous compounds; and the

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 37

atomic weight of the element X can be determined by an analysis
of its alum. For this purpose it is best to ignite the ammonium
alum, which leaves a residue of the oxide X 2 O 8 .

Thus, if a grm. of the alum leaves 6 grm. of oxide, the value of
X is calculated from the expression:

a : b = (NH 4 ) 2 S0 4 -X 2 (S04) 3 .24H 2 : X 2 O 3

= [132 + 2 X + 288 + 432] : [2 X + 48].

By this means Lecoq de Boisbaudran, who discovered gallium,
found that 3-1044 grm. of its ammonium alum left on ignition a
residue of 0*5885 grm. of the sesquioxide; whence Ga = 70-1.

v. The Method of the Periodic Law.

An adequate account of the periodic law is necessary to an
appreciation of its value as a guide to the magnitudes of the
atomic weights of the elements; but this must be deferred to a
later chapter.

It will be sufficient to state here that a natural connection
exists between the properties of an element and its atomic weight,
and therefore that the order of magnitude of the atomic weight of
an element may be judged from a study of the properties of the
element and its compounds. Examples of this use of the periodic
law will be given later.

The application of the foregoing methods of atomic - weight
determination is well illustrated by the case of carbon.

4. The Atomic Weight of Carbon

The determination of the atomic weight of carbon has consisted
of two parts:

i. The determination of the order of magnitude,
ii. The determination of the exact value.

i. Determination of the Order of Magnitude of the Atomic Weight.

Dalton and his contemporaries attributed the value 6 to the
atomic weight of carbon, but this was really only an equivalent
weight. The following is the evidence that the atomic weight is
about 12:

(a) Avogadro's Theory. Never fewer than 12 parts by weight of
carbon are present in a molecular proportion of any of the gaseous
or volatile compounds of this element.

(6) Chemical Displacement Use might be made of the argu-

38 CHEMICAL THEORY

ment that, for example, hydrogen is displaceable from methane in
four equal fractions, but carbon not fractionally; whence it follows
that the formula for methane is CH 4 and the atomic weight of
carbon 12.

(c) The Law of Specific Heat. Although Dulong and Petit's law
does not apply strictly to an element whose atomic weight is less
than 30, and the specific heats of diamond and graphite differ
widely from each other at ordinary temperatures, at 600 the
specific heats of these two allotropic forms of carbon, which vary
with temperature, become almost constant and equal, and give an
atomic heat of 5-5, if C = 12, a value which is comparable with
the atomic heats " of analogous elements.

(d) The Law of Isomorphism. The iodides of carbon and
silicon are isomorphous; therefore they are similarly composed, and
the atomic weights of carbon and silicon are in the ratio 12 : 28.

(e) The Periodic Law. With an atomic weight of 12, carbon is
appropriately placed in the periodic table between boron (10-82)
and nitrogen (14-01); and is thus the first or "typical" element of
the fourth group. If carbon forfeited its place owing to an altera-
tion in the magnitude of its atomic weight, there is no vacant place
in the periodic table which this element could fill, nor is any element
known which could occupy the place of carbon.

ii. Determination of the Exact Value of the Atomic Weight.

There are two ways in which the atomic weight of carbon has
been determined exactly:

(a) By estimating the densities of its gaseous compounds.

(6) By the combustion of carbon or the analysis of its compounds.

(a) It has already been pointed out that gas or vapour density
is simply related to molecular weight only when Avogadro's theory
is rigidly true. This, however, is never the case; but an "ideal"
density can sometimes be calculated from carefully ascertained
data. This has been done 1 for the three gases: carbon monoxide
(CO), carbon dioxide (CO 2 ), and acetylene (C 2 H 2 ).

00 C0 2 ja

Experimental density (O 2 = 1) 0-87495 1-38324 0-8194

"Ideal" density ......... 0-87516 1-37516 0-81331

Molecular weight ...... 28-005 44-005 26-026

Atomic weight of carbon ... 12-005 12-005 12-005

Ann. Cfam. Phys. y 1898 [vii.], 15 t 5; 1910 [viii.], 19, 441.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 39

It appears from these figures that the method leaves nothing to
be desired from the point of view of accuracy.

(6) The atomic weight of carbon has been determined by
several chemists by burning diamond or carefully purified graphite,
weighing the carbon dioxide produced, and then calculating the
result from the proportion:
Weight of CO 2 : weight of C : : 32 + atomic weight C : atomic weight C.

The following are the results, as originally given, and as cor-
rected by Scott:

Uncorrected. Corrected by Scott. 8

Dumas and Stas 1 11-9975 11-9938

Erdmann and Marchand 2 12-0093 12-0054

Koscoe 3 12-0029 11-9973

Friedel 4 12-0112 12-0056

Vander Plaats 5 12-0031 12-0017

The ignition of organic silver salts, such as the acetate and
tartrate, which leave a residue of pure silver, serves as a means
of estimating the atomic weight of carbon; or the silver may be
estimated electrolytically, as was done by Hardin, 7 with the follow-
ing results, obtained with silver acetate and benzoate respectively:

(1) C 2 H 3 O 2 Ag : Ag = 100 : 64-637
whence atomic weight of carbon = 12-000.

(2) CVH 6 2 Ag : Ag = 100 : 47-125
whence atomic weight of carbon = 12-001.

[Ag = 107-880, H = 1-00762,0 = 16-00.]

The above results are selected from amongst others as typical;
they serve to show the degree of accuracy which has been attained
in the determination of the atomic weight of carbon. This value
lies between 12-000 and 12-005, and may be taken to be 12-003.

5. Determination of Molecular Weights (in Solution)

The establishment of molecular weights by the determination
of gas and vapour densities has been fully considered in the pre-
ceding pages. By the study, however, of the influence of dissolved

i Dumas, Pogg. Annalen, 1838, 44, HO.

8 Erdmann and Marohand, J. prakt. Chem., 1841, 23, 159.

3 Roscoe, Compt. rend., 1882, 94, 1180.

4 Friedel, Bull. Soc. C/iim., 1884 [ii.], 4*, 100.

Van der Plaats, Compt. rend., 1885, 100, 52.

Scott, Trans. Ctiem. Soc., 1897, 71, 550.

'Hardin, J., Amer. Chem. Soc., 1896, 18, 990.

40 CHEMICAL THEORY

substances on the solidifying- and boiling-points of liquids, the
molecular weights of substances in solution in these liquids may be
determined; and it will be appropriate to consider here these newer
methods of molecular- weight determination.

It is well known that salt water freezes at a lower temperature
than fresh water, and that sea ice when melted yields fresh water.
Thus, when a dilute solution of salt in water is cooled, crystals of
pure ice begin to separate from the solution at a temperature a
little below 0. Blagden, in 1788, showed that the depression of the
freezing-point of water by a dissolved salt is directly proportional
to the amount of salt present. The boiling-point of water, on the
other hand, is raised by salt in solution, and the elevation of boiling-
point is directly proportional to the amount of salt dissolved. In
1883-4 F. M. Raoult discovered that not only are the depression
of freezing-point and rise of boiling-point of a solvent proportional
to the number of molecules of a particular substance in solution,
but that equimolecular proportions of different substances have the
same influence on the freezing- and boiling-points.

RaouWs law, which applies equally to freezing- and to boiling-
points of solvents, may be stated thus:

The depression of freezing-point and elevation of boiling-point of
a solvent by a quantity of dissolved substance are directly proportional
to the number of molecules of the substance in solution, and con-
sequently inversely proportional to its molecular weight.

Or, otherwise:

Equimolecular solutions, with the same solvent, have the same
freezing- and boiling-points.

Evidently these facts provide a means of comparing molecular
weights, or of determining them if a substance of known molecular
weight is chosen as a standard of comparison. It should be added
that the extent to which a freezing- or boiling-point is affected
depends also upon the solvent; consequently the first procedure
is to determine the freezing- or boiling-constant (K) for a particular
solvent by the use of a substance of known molecular weight. This
constant is the number of degrees the freezing-point is lowered or
boiling-point raised by 1 grm.-molecule of the substance dissolved
in 100 grm. of the solvent.

For instance, 2 grm. of cane sugar dissolved in 100 grm. of
water cause a depression of the freezing-point$A = 0-11. Since

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 41

the molecular weight of cane sugar (C IZ H^O U ) is 342, the freezing-
constant, K, for water, sometimes called the molecular depression, is

0il x 342 = 19

The same quantity of sugar dissolved in the same amount of
water raises the boiling-point of the water 0-030. Therefore
the boiling-constant or molecular elevation, K, for water is:

0-030 X 342 _. 5 . 2

2 " '

When the freezing- or boiling -constant K for a solvent is
known, an unknown molecular weight is calculated from observed
data as follows:

Let K = depression or rise caused by 1 grm.-mol. of a substance

in 100 grm. of solvent (known constant).
M S = weight of substance taken.
L = weight of solvent taken.
A = observed depression or rise.
M = required molecular weight.

Then, since the observed depression of freezing-point or rise of
boiling-point is directly proportional to the amount of substance
taken, and inversely proportional to the amount of solvent,

A K X S X 100 ,, 100 KS
A = TfrrrT or M = T-= .
M X L A L

Practical Methods.

The prime necessity for the experimental determination of
molecular weights of substances in solution is a thermometer which
will indicate accurately hundredths of a degree. If this ther-
mometer is to be used both for freezing- and boiling-points, it
would appear necessary for it to have a long range in addition.
Real temperatures, however, have not to be read; only temperature
differences. Consequently a thermometer has been devised by
Beckmann with a range of about six degrees, the scale being
divided into hundredths of a degree, and furnished with a reser-
voir of mercury from which mercury can be added if low tem-
peratures are to be recorded, and into which mercury can be
driven when the instrument is to be used for higher temperatures.
By the use of this device the same thermometer can be employed
for temperatures near the freezing- as well as the boiling-point
of water or otherf sol vent.

CHEMICAL THEORY

The Cryoscopic Method, v

The determination of molecular weights by the cryoscopic

method, that is, by observing the depression of freezing-point, is

carried out in the apparatus of Beckmann shown in the figure.

The tube (A), furnished with a side limb for the introduction of
the substance, is fitted with a cork through
which the thermometer (T) and platinum
stirrer (S) pass. The lower part of this tube
is surrounded by a wider tube (B) which
provides an air jacket between the tube (A)
and the freezing-mixture contained in the
outer vessel (C). This freezing - mixture,
whose temperature should be about 5 below
the freezing-point of the solvent employed, is
also furnished with a stirrer (S 1 ). A weighed
quantity of water, or other solvent, is placed
in the tube (A) and then frozen. Owing to
under-cooling the temperature indicated by
the thermometer falls below the freezing-
point, and then quickly rises again, and be-
comes stationary at that point as soon as ice
separates. When the freezing-point of the
solvent has been indicated on the arbitrary
scale of the thermometer, a weighed quantity
of the substance is introduced and the freez-
ing-point of the solution determined. The
amount of substance added should produce a
depression of about 0-5. The determination
may be repeated after the addition of a
further quantity of substance. The reading
should, however, be taken when a minimum

quantity of the pure solid solvent has separated, so that the

concentration of the solution may not be appreciably increased.

The following are important freezing-constants (K):

Water 18-6; acetic acid 39; benzene 50; phenol 73.

EXAMPLE. Successive quantities of 0-317, 0-394, and 0-5152
grm. of a substance were dissolved in 18-054 grm, of benzene,
the depressions of freezing-point being 0-278, 0-348, and 0-452
respectively; what is the molecular weight of the substance? The

Fig. 6

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 43

molecular lowering of the freezing-point of benzene (K) is 50.
(Institute of Chemistry, July, 1902.)

100 KS

"sir-

100 X 50 X 0-317

i. M =

ii. M =

M1U
S=B

0-278 X 18-054
100 X 50 X 0-394

0-348 X 18-054
100 X 50X0-5152

0-452 X 18-054

= 315-8.

= 313-6.

= 315-7.

The Ebulliscopic Method Beckmann's Apparatus.

The tube (A) (fig. 6) em-
ployed in the Beckmann appa-
ratus for determining elevation
of boiling-point resembles that
in which freezing-point deter-
minations are carried out; but,
in addition to the side tube
for the introduction of the sub-
stance, it is provided with an-
other tube (B) fitted with a
reflux condenser for the con-
densation of the vapour arising
from the boiling liquid. In
order to prevent super-heating
of the liquid, and consequent
irregular boiling, a short piece
of stout platinum wire (C) is
fused into the bottom of the
tube, which also contains some
small beads which surround the
lower part of the thermometer
bulb, and serve to break and
distribute the bubbles of vapour
as they rise. In addition to
this, the boiling-tube is sur-
rounded with a wider vessel (D)
packed with some non-conduct-
ing material to prevent loss of
heat by radiation, or sometimes Fig. e

44 CHEMICAL THEORY

with a glass envelope containing the vapour of the boiling solvent
The whole apparatus stands upon a sheet of asbestos (E), below
which the burner for heating is placed.

In carrying out an experiment a weighed quantity of the
solvent is heated until it boils briskly, and its temperature has
become constant. If the condenser is acting efficiently the solvent
should not lose weight; but about 0-3 grm. should be subtracted
from its weight to allow for the quantity required to wet the
internal walls of the tube and condenser. After the boiling-point
of the solvent has been recorded, the weighed quantity of the
substance is introduced, and a reading again taken when the
temperature has become constant. As in the case of freezing-point
determinations, successive quantities of substance may be added to
the same quantity of solvent, and corresponding readings taken.
If much time elapses between the observations of the boiling-points
of solvent and solution, it is necessary to read the barometer, and
make a correction for change of atmospheric pressure during the
interval.

The Modified Landsberger Apparatus.

A method of determining elevation of boiling-point, introduced
by Sakurai, 1 modified by Landsberger, 2 Walker and Lumsden, 3 and
others, and more recently by Turner and Pollard, 4 consists in raising
the solvent to its boiling-point by passing into it the vapour of the
same liquid boiling in another vessel. The vapour condenses, and
its latent heat eventually causes the solvent to boil, although the
boiling-point after the addition of the substance is above that of
the pure solvent. As the amount of the solvent continuously in-
creases by condensation of vapour, it is estimated by weighing
or measuring after condensation has been arrested instead of
before heating is begun.

By this method all possibility of superheating is avoided, and
accurate results are rapidly obtained.

The construction of the apparatus is shown in fig. 7. The vessel
(A), about 16 cm. high and 3 cm. in diameter, is fitted with a two-
holed cork through which pass the thermometer (T) and the delivery
tube (B) by which vapour is conveyed to the bottom of the vessel

i Trans. Cham. Soc., 1892, 61, 994. 2 tier., 1898, 31, 4G1.

8 Trans. Ckem. Soc., 1898, 73, 502.

* Trans. Ckem. Soc., 1910, 97, 1184, Proc. Chem. Soc., 1913, 29, 349.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 45

from the liquid boiling in the flask (F). A small hole (C) in the
upper part of the vessel
allows uncondensed va-
pour to pass into the
outer vessel (D), where
it forms a vapour jacket
and is then condensed,
either here or by sub-
sequent cooling after
escaping by the side
tube (E).

The boiling-constants
(K) of important liquids
are: Water 5*2; ethyl
alcohol 11-7; chloroform
39; benzene 27.

EXAMPLE. Turner
found that 1-150 grm.
of diphenylamine
(C 6 H 6 ) 2 NH, dissolved in
42-82 grm. of chloro-
form caused the boiling-
point of the solvent to rise 0-618. K = 39. Find the molecular
weight of diphenylamine.

M = 1Q 9 KS = 100X39 X 1-150 _ 169t5
AL 0-618 X 42-82

Theory for (C 6 H 5 ) 2 NH = 169-1.

Fig. 7

6. Molecular Complexity

The methods and results of determining the molecular weights
of gases and vapours and of substances in solution have been
reviewed in the preceding pages; and it appears that the molecules
of substances in solution are often of the same order of magnitude
as those of the same substances in the state of vapour. For example,
ferric chloride in a state of vapour at about 750 consists of mole-
cules represented by the formula FeCl 3 , and the elevation of the
boiling-point of ether or alcohol by dissolved ferric chloride points
to the same molecular formula. The reason for this identity of
molecular state is to be found in the fact that the vaporous state

46 CHEMICAL THEORY

and the state of solution are analogous to each other, and that the
process of vaporization of a solid or liquid, with the consequent
distribution of its molecules through space, resembles the process of
solution of the same substance, and the distribution of its molecules
throughout the solvent.

Yet the molecular state of a dissolved substance depends some-
times upon the liquid in which it is dissolved. Hydrogen chloride,
for example, forms molecules when dissolved in benzene and nitro-
benzene which may be as much as five times as great as the gaseous
molecule; its molecules are then said to be associated. With regard
to liquids themselves, there is good reason to believe that their
molecules are often associated. Consider water, for example.
Water is the first of the series of four hydrides: H 2 O, H 2 S, H 2 Se,
H 2 Te; three of these are gases; why, therefore, is water a liquid?
Since the atomic weight of oxygen is the least of the atomic weights
of the four elements combined with hydrogen in this series, and
volatility diminishes from H 2 S to H 2 Te, water would be expected
to be the most instead of the least volatile of the four hydrides.
The reason water is a liquid at atmospheric temperature must be
that it forms complex molecules (H 2 0) n . Steam even appears to
contain a very small proportion of molecules, which are regarded as
double, since its density is a little greater near the point of con-
densation than corresponds with the simple formula H 2 O. Liquid
water undoubtedly consists of associated molecules, e.g. (H 2 O) 2 and
(H 2 O) 3 , whilst ice is believed to be (H 2 0) 3 only. It is noteworthy
that hydrogen fluoride, which follows water in the periodic classifi-
cation, also contains associated molecules, and has an anomalous
boiling-point.

Regarding benzene, C 6 H , there is evidence that near its point of
condensation the saturated vapour begins to contain double mole-
cules, (C 6 H 6 ) 2 , and that liquid benzene consists entirely of these
molecules, until near its freezing-point, when (0 6 H 6 ) 4 molecules
begin to appear, and increase in number until solid separates,
which consists wholly of (C 6 H 6 ) 4 molecules.

It is remarkable that the association of various molecules is
promoted by the entire absence of water. Thus benzene, which
usually boils at 80, was found by H. B. Baker 1 to boil at 106 after
being dried over phosphoric oxide for eight years. This rise of
boiling-point must, no doubt, be attributed to molecular association.

1 Ckem. Soc. Trans., 1922, 121, 570.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 47

Concerning the molecular state of solids little or nothing has
been known until recently; but the examination of the X-ray
spectra produced by solids is throwing much light on their mole-
cular structure. The nature of the problem presented by solids
may be realized by considering a simple example. Carbon and
silicon are related elements; yet carbon when it burns forms a gas
and silicon a solid. Thus the molecules of C0 2 are not associated,
but those of Si0 2 are, if indeed Si0 2 molecules can be said to exist
at all. What is the reason for this difference? The difference
may be expressed, [though not explained, by saying that any field
of influence outside the C0 2 molecule, by which other molecules of
the same kind might be attracted, and associated, is very limited,
or else the attractive force exerted therein is very small; whereas
in the case of the Si0 2 molecule the external field of attractive
force is considerable in extent or strength or both.

The difference between the two cases is referable to the difference
between the carbon and silicon atoms, and must depend ultimately
upon the different structures of these atoms. Meanwhile it may be
stated that the " crystal unit ", i.e. the smallest unit that takes part
in crystal growth of silica, is (SiO 2 )3.

With regard to the formulae to be applied to liquid and solid
compounds there is much difficulty if real molecules are to be
represented. Indeed it can hardly be asserted that molecules,
NaCl, of common salt exist either in the solid state or in solution.
It is not necessary, however, to have molecular formulas in order to
represent chemical reactions, since the simplest formulas represent-
ing the inherent properties of compounds are sufficient for all
ordinary purposes. The formula CaC0 3 , for example, serves for
chalk; it would be interesting to know what is the molecular or
crystal unit of this compound, but such knowledge is not necessary
for the representation of its common reactions.

In some cases, however, the smallest empirical formula would
be untrue; e.g. hyponitrous acid is H 2 N 2 2 and not HNO; and
benzene is C 6 H 6 and not CH nor C 2 H 2 .

7. The Molecular Compositions of Compound Gases

The evidence on which the molecular formula H 2 O for water is
based has already been considered. There are a number of com-
pound gases whose molecular formulae may be established by the

48 CHEMICAL THEORY

application of the principles set forth in this chapter; and these
will now be dealt with.

It has already been shown that the molecules of hydrogen,
chlorine, and oxygen are diatomic. This follows, it will be re-
membered, from the fact that the hydrogen chloride formed from
equal volumes of hydrogen and chlorine occupies twice the volume
of each separate gas, and that steam occupies twice the volume of
its constituent oxygen at the same temperature and pressure. By
an extension of the principle here employed the number of atoms
of a gaseous element within the molecule of a compound gas may
always be determined.

Thus, since it can be shown that 2 volumes of ammonia gas
yield when decomposed 3 volumes of hydrogen and 1 volume
of nitrogen, it follows, provided the nitrogen molecule is diatomic,
that ammonia must be represented by the formula NH 3 ; for the
molecular change on the decomposition of ammonia is:

2 mols. ammonia yield 3 H 2 + N 2 ,
consequently 2 NH 3 = 3 H 2 + N 2 .

The argument may be put in another way. Since the volume
of the ammonia is to that of the hydrogen as 2 : 3, the atomic con-
centration of hydrogen in ammonia is to that in free hydrogen as
3:2; and since the volume of ammonia is to that of nitrogen as
2:1, the atomic concentration of nitrogen in ammonia is to that
in free nitrogen as 1:2; whence the formula NH 3 follows.

In the case of a gas containing a solid element, such as sul-
phurous anhydride, the additional estimation of the density of the
gas suffices to show how many atoms of the solid element it
contains, provided the atomic weight of this element is known.
Thus, (a) the gas produced by burning sulphur in oxygen measures
the same volume as the oxygen; therefore the molecule of this gas
contains 2 atoms of oxygen; (6) the density of the gas is 32, and
its molecular weight consequently 64, whilst the weight of oxygen
within its molecule is 32, and the atomic weight of sulphur is 32;
therefore it follows that the molecule of the gas contains 1 atom
of sulphur, and that its molecular formula is SO 2 .

The following statements epitomize the evidence for the mole-
cular formulae of a number of the best-known gases:

Hydrogen Chloride.

That 1 volume hydrogen + 1 volume chlorine give 2 volumes

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 49

hydrogen chloride is fundamental to the molecular theory. The
following facts suffice to prove this relation:

(a) Electrolysis of an aqueous solution of hydrogen chloride
under suitable conditions yields equal volumes of hydrogen and
chlorine.

(6) Sodium amalgam removes the chlorine from hydrogen
chloride gas, and the remaining hydrogen occupies half the volume
of the hydrogen chloride.

Water and Steam.

(a) Electrolysis of acidified water yields hydrogen and oxygen
in the proportion of 2 volumes of the former to 1 of the latter.

(6) When a volume of electrolytic gas, i.e. a mixture of 2
volumes of hydrogen with 1 volume of oxygen is exploded in
a eudiometer kept at a temperature above the boiling-point of
water, the volume of the resulting steam is two-thirds the volume
of the mixed gases. Therefore

2 vol hydrogen + 1 vol. oxygen yield 2 vol. steam.

Carbonic Anhydride.

When carbon is burnt in oxygen gas the volume of the gas
remains unaltered. Therefore a molecule of the gaseous product
contains 2 atoms of oxygen (O 2 = 32).

The density of carbonic anhydride is 22; therefore its molecular
weight is 44. Within this molecular proportion are 32 parts
(0 2 ) of oxygen, and therefore 12 of carbon. But 12 is the atomic
weight of carbon. Therefore carbonic anhydride is CO 2 , and is
rightly called carbon dioxide.

Sulphurous Anhydride.

When sulphur is burnt in oxygen the volume of the gaseous
product is the same as that of the oxygen. The density of sul-
phurous anhydride is 32, and its molecular weight 64. The atomic
weight of sulphur is 32; therefore, by the same argument as applies
to carbon dioxide, sulphurous anhydride is sulphur dioxide, SO 2 .

Hydrogen Sulphide,

When hydrogen sulphide gas, confined over mercury, is decom-
posed by electric sparks, or when its sulphur is removed by means
of tin heated in the gas and so converted into sulphide, the volume

(DCO) 5

50 CHEMICAL THEORY

of the remaining hydrogen is equal to the volume of the original
hydrogen sulphide, whose formula is consequently H 2 S n . That
n = 1 is proved by the fact that the gas density is 17 and
molecular weight 34; for of this 32 parts must be sulphur, and 32
is the atomic weight of sulphur. Thus the formula for hydrogen
sulphide is proved to be H 2 S.

Nitrous Oxide.

Potassium, sodium, copper, and other metals remove the oxygen
from nitrous oxide when heated in the gas, leaving nitrogen.
There is some risk of nitrite being produced if the two former
metals are heated too strongly in the gas, but strongly heated
copper removes only the oxygen, and leaves all the nitrogen in a
pure state. By this means it may be shown that nitrous oxide
contains its own volume of nitrogen, and therefore that its mole-
cule contains 2 atoms of this element. The density of nitrous
oxide is 22, and its molecular weight is 44, and this weight contains
28 parts (N 2 ) of nitrogen, and therefore 16 parts of oxygen. Since
16 is the atomic weight of oxygen the molecule of nitrous oxide
contains 1 atom of this element, and therefore the molecular
formula for the gas is N 2 O.

The same conclusion is reached by mixing nitrous oxide with
its own volume of hydrogen and exploding the mixture. After
condensation of the steam pure nitrogen remains equal in volume
to the nitrous oxide. Thus it is shown, not only that nitrous
oxide contains its own volume of nitrogen, but that the oxygen it
contains would occupy half that volume, since it combines with
a volume of hydrogen equal to that of the nitrous oxide. These
facts are sufficient to establish the formula N 2 O for nitrous oxide.

Nitric Oxide.

If potassium is heated in nitric oxide the vigorous combustion
which takes place results in the formation of nitrite and nitrate;
but a spiral of iron wire heated electrically removes all the oxygen
from the gas without combining with the nitrogen, and the residual
nitrogen then occupies half the volume of the original nitric oxide.
This proves that a molecule of nitric oxide contains 1 atom of
nitrogen (N = 14). The density of nitric oxide is 15, and, since
its molecular weight is 30, the molecule contains 1 atom of
oxygen (O = 16), and the molecular formula is NO.

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 51
Ammonia.

When ammonia solution is dropped into chlorine gas, hydrogen
chloride is formed, and nitrogen set free. The experiment may be
carried out in a long graduated tube, sealed at one end and pro-
vided at the other end with a cork furnished with a tap funnel.
Ammonia solution is passed through the funnel into the chlorine,
and the reaction is accompanied by a greenish flame and fumes
of ammonium chloride. After the ammonia has been added in
excess, dilute sulphuric acid is introduced to combine with the
excess of ammonia, after which water is allowed to enter until
the gas in the tube is at atmospheric pressure, when the flow of
water ceases. Then it is found that the gas, which is nitrogen,
fills one-third of the tube. Since hydrogen and chlorine combine
in equal volumes to form hydrogen chloride, the hydrogen of the
ammonia from which the nitrogen has been liberated would have
occupied three times the volume of this nitrogen. This shows
that ammonia, when decomposed, yields 1 volume of nitrogen
to 3 of hydrogen; but since the volume of ammonia gas which
is thus decomposed is unknown, all that this experiment reveals
is that the molecule of ammonia is (NH 3 ) n .

The relation between the volume of ammonia and the volumes
of its decomposition products may be determined by confining a
measured volume of the gas over mercury and passing electric
sparks through it until expansion ceases. The gas will then
have been decomposed into a mixture of hydrogen and nitrogen
which will occupy twice the volume of the ammonia. That this
mixture consists of 3 volumes of hydrogen and 1 volume of
nitrogen is known from the previous experiment, or it may be
shown by adding excess of oxygen and exploding the mixture.

Thus for example:

Volume of ammonia = 10 c. c.

Volume of nitrogen + hydrogen after sparking = 20

Volume after addition of oxygen =75.0

Volume after explosion =52-5

Thus 22*5 c. c. of gas have disappeared, of which 15-0 c. c. must
have been hydrogen. So it follows that 10-0 c. c. of ammonia were
decomposed by electric sparks into 15 c. c. of hydrogen and 5 c. c.
of nitrogen; and, as shown before, this proves the molecular
formula NH 3 for ammonia; for 2NH 3 = 3H 2 + N 2 .

52 CHEMICAL THEORY

Phosphine. *"

The case of phosphine differs from that of ammonia because,
when the gas is decomposed by electric sparks, the liberated phos-
phorus remains as a solid whose volume is negligible. Conse-
quently, the proportion of phosphorus in the molecule must be
discovered by density determination as in the case of sulphur
dioxide, hydrogen sulphide, &c.

Two volumes of phosphine, decomposed by electric sparks,
yield 3 volumes of hydrogen. Therefore a molecule of the gas
contains 3 atoms of hydrogen.

The density of phosphine is 17, and its molecular weight 34.
Consequently, the proportion of phosphorus within the molecular
proportion of phosphine is 31. But 31 is the atomic weight of
phosphorus. Therefore the molecule of phosphine contains 1
atom of phosphorus, and so its formula is PH 8 .

Carbon Monoxide.

Carbon monoxide can be converted into carbon dioxide by
exploding it with oxygen, when it is found that 2 volumes of the
gas combine with 1 volume of oxygen to form 2 volumes of carbon
dioxide; or, since the molecular formulae C0 2 and O 2 are known, in
the equation,

2 C x O y + O 2 = 2 CO 2 ,

x and y both = 1, so that the molecular formula CO is proved.
This conclusion is confirmed by the density of the gas, which is
14, whence the molecular weight is 28; and C = 12, O = 16, so
that CO = 28.

Methane, Ethylene, and other Hydrocarbons.

If a certain volume of a hydrocarbon is exploded with a
known volume of oxygen used in excess, the resulting moist gas,
measured at atmospheric temperature and pressure, consists of
carbon dioxide mixed with unused oxygen. The volume of
carbon dioxide formed is estimated by absorbing this gas in
sodium hydroxide solution, and the total volume of oxygen used,
part of which has produced carbon dioxide, and part water, is
shown by the difference between the original and the remaining
volume of oxygen. These data are sufficient to establish the
molecular formula of the hydrocarbon.

EQUIVALENT, ATOMIC, AND MOLECULAE WEIGHTS 53

For, consider the gaseowa hydrocarbon C x H y . The result of
its explosion with oxygen is represented by the equation

C x H y + (x + |)0 2 = *C0 2 + |H 2 0.

The volume of steam formed and condensed is not measured;
but when the volume of carbon dioxide, referred to that of the
hydrocarbon as unity, which is x, has been ascertained, the value
of y is found by subtracting this from the total volume of oxygen
used, referred to the same standard, and multiplying the remainder
by 4.

When x and y are found, the formula of the hydrocarbon
is settled. Vapour density will confirm the formula, but is not
necessary to establish it.

Methane.

When a mixture of 10 c. c. of methane with 30 c. c. of oxygen
is exploded, the resulting gas, measured at the same temperature
and pressure, is a mixture of 10 c. c. of carbon dioxide and 10 c. c.
of oxygen.

Thus 1 volume methane requires for combustion 2 volumes
oxygen, and yields 1 volume carbon dioxide.

So in the equation

C x H y + (x + ^)O 2 = #CO 2 + |H 2 O, x = 1 and f = 1;

\ / 25 4

therefore the formula for methane is CH 4 ; or otherwise, because
the volume of the carbon dioxide produced is equal to the volume
of the methane, a molecule of the latter contains 1 atom of carbon;
and, because the volume of the oxygen required to burn the
hydrogen of methane is equal to the volume of the methane,
the atomic concentration of hydrogen in the methane molecule
is twice what it is in the free hydrogen molecule; i.e. there are
4 atoms of hydrogen in methane. Thus the molecular formula
for methane is CH 4 .

Ethylene.

When a mixture of 10 c. c. of ethylene with 40 c. c. of oxygen
is exploded, the resulting gas, measured at the same temperature
and pressure, is a mixture of 20 c. c. of carbon dioxide and 10 c. c.
of oxygen.

54 CHEMICAL THEORY

Thus 1 volume ethylene requires for combustion 3 volumes
oxygen and yields 2 volumes carbon dioxide.
So in the equation

consequently the formula for ethylene is C 2 H 4 .

Or, to employ the alternative argument, since the volume of
the carbon dioxide produced is twice the volume of the ethylene,
a molecule of this hydrocarbon contains 2 atoms of carbon; and
since the volume of oxygen required to burn the hydrogen of
ethylene is equal to the volume of the ethylene, this hydrocarbon
contains 4 atoms of hydrogen. Thus, again, the molecular formula
for ethylene is C 2 H 4 .

In a similar way the molecular formula of any gaseous hydro-
carbon may be established.

SUMMARY

EQUIVALENT WEIGHT. The equivalent weight of an element
is that weight of it which combines with, or displaces from com-
bination, an agreed weight of a standard element. The standard
is: O = 8-00.

ATOMIC WEIGHT. The atomic weight of an element is the
ratio between the weight of its atom and that of the atom of
a standard element. The standard is: O = 16-00.

DETERMINATION OF ATOMIC WEIGHT:

(a) Exact estimation of chemical equivalent.
(6) Decision as to order of magnitude.
Guiding principles: i. Avogadro's theory.

ii. Chemical displacement.
iii. Law of specific heats.
iv. Law of isomorphism.
v. Periodic law.

PRINCIPLE OF CHEMICAL DISPLACEMENT. When l/7i th of the
proportion of a constituent element in a chemical compound can
be displaced by another element, a molecule of the compound
contains at least n atoms of that element.

LAW OF SPECIFIC HEATS: DULONG AND PETIT'S LAW. The
specific heats of the solid elements are inversely proportional to

EQUIVALENT, ATOMIC, AND MOLECULAR WEIGHTS 55

their atomic weights; or the atoms of the solid elements have
the same capacity for heat.

Specific heat X atomic weight = 6 4 (appro*.)
or atomic weight =

specific heat

LAW OF ISOMORPHISM: MITSCHERLICH'S LAW. The molecules
of isomorphous substances contain equal numbers of atoms, which
when not identical are analogous.

DETERMINATION OF MOLECULAR WEIGHTS: RAOULT'S LAW.
The depression of freezing-point, and elevation of boiling-point of a
solvent caused by any dissolved substance are directly proportional
to the number of molecules of the substance in solution, and
consequently, inversely proportional to its molecular weight; or
equimolecular solutions, with the same solvent, have the same
freezing- and boiling-points.

CHAPTER III

OLDER VIEWS OF VALENCY AND CHEMICAL
CONSTITUTION

In the preceding pages the experimental foundations of the
atomic and molecular theories, as these were laid by the chemists
and physicists of the nineteenth century, have been considered;
and it has been found possible to exhibit and expound these
without reference to modern conceptions of the atom; this is
because the phenomena concerned have been superficial and have
not dealt with the inter-relations of the atoms themselves in
chemical compounds.

When, however, the subjects of valency and chemical constitu-
tion are approached the case is otherwise. It is impossible at the
present time to consider these subjects adequately without bringing
into view the atom as it appears in the light of to-day's knowledge.
Moreover, the tide of this new knowledge is so powerful that
much that was considered sound and stable has been broken by its
flood; so that the first task of the chemical philosopher is to
strengthen what remains of the harbour of his thought, whilst the
flotsam disappears.

Consequently this chapter on "Older Views of Valency and
Chemical Constitution" is historical; touching lightly the great
subjects with which it deals, it brings chemical knowledge up to
the boundary of the new domain, and leaves for a further chapter
the task of exploration. It will be enough for the present purpose
if what was temporary and must disappear can be distinguished
from what is permanently useful. If this purpose is achieved, if
impedimenta are dropped, and only useful tools and weapons are
retained, there is hope that in the new field valuable possessions
may be acquired.

It has already been seen, with regard to an element, that
Atomic weight = n X equivalent weight;

and that n is the valency or atomic value of the element.

VALENCY AND CHEMICAL CONSTITUTION 57

The doctrine of valency, in the form in which it was held during
the latter part of the nineteenth century, was a matter of slow
development. It arose during the growth of organic chemistry,
because of the need of a theory of structure in systematizing that
branch of the science. Chemists developed the habit of referring
various organic compounds to a few simple inorganic types, and this
idea proved fertile, for it stimulated research, and led to important
discoveries. These simple types were:

H\ H\ H
HJ Cl/ H

It cannot be said that the conception of valency was definitely
contained in these types, yet it was not far off.

The chief exponent of the theory of types was Gerhardt; but
it was E. Frankland who first introduced the idea of saturation
capacity or valency. Frankland showed that whilst 1 atom of tin
was capable of combining with two atoms of oxygen to form the
dioxide SnO 2 , a molecule of the compound tin diethyl, Sn(C 2 H 5 ) 2 ,
or SnEt 2 , could combine with only 1 oxygen atom forming the
compound SnEt. 2 O. Thus it appeared that the tin atom had a
certain saturation capacity, that it could combine with not more
than 2 atoms of oxygen or their equivalent; and that its power of
combining with oxygen was diminished by the extent to which it
was already combined with other atoms or groups of atoms. >

The principle was further illustrated by Frankland by reference
to such compounds as j^H NI

PH 3 r PCI*

in which the atoms of nitrogen and phosphorus combined with
3 atoms of hydrogen or halogen; and by Kekul, who showed that
the carbon atom could combine with four other atoms, as in the
compounds

The phenomenon here illustrated is now called valency, about
which the following statement may be made:

The valency of an element indicates the number of otter atoms
with which one of its atoms can directly combine.

An atom may be uni-, bi-, tri-, quadri-, quinqui-, sexi-, septi-,
or even octi-valent; 1 corresponding terms, formerly used, are monad,

1 These numerical prefixes are proposed as standard ones by J. D. Main Smith, Chemistry
and Industry, 1927, 188.

58 CHEMICAL THEORY

dyad, triad, tetrad, &c. In the compounds cited above the nitrogen
and phosphorus atoms are trivalent, and the carbon atom is quadri-
valent; whilst the hydrogen, chlorine, and iodine atoms are univalent.
Hydrogen is never more than univalent, and therefore its atom
is chosen as the standard of valency; chlorine is univalent with
regard to hydrogen and metals, and, indeed, probably to all ele-
ments except oxygen; it may therefore replace hydrogen as a
standard.

The following hydrides exhibit the valency of a number of
elements:

Valency 1234

FH OH 2 NH 3 CH 4

C1H SH 2 PH 3 SiH 4

BrH AsH 3

IH SbH 3

and the following halides 1 illustrate valency more extensively:

078

SF fl - OsF 8
TeF r , -

Oxygen is shown above to be bivalent. It is seldom other than
this; and if oxygen is bivalent, the large number of oxides that
exist may be classified to show valency, on the assumption that the
valency of an element is equal to twice the number of oxygen
atoms with which one of its atoms combines.

Valency 1

NaCl

MgCl,

BC1 3

CC1 4

PF &

KC1

ZnCl 2

PC1 3

SiCl 4

AsF fi

AgCl

HgCl.

A1C1 3

SnCl 4

SbF 6

Valency 1

Na 2

MgO

B 2 3

CO 2

S0 3

CIA

OsO 4

K 2 O

CaO

A1 2 S

SiO 2

P 2 6

Cr0 3

(I 2 7 )

Eu0 4

Ag 2

ZnO

Fe 2 3

Pb0 2

Bi 2 6

UO 3

Mn 2 O r

These oxides are in a different category from that of the
foregoing hydrides and halides. In those the molecular formulae
have in all cases been established by vapour density or other
measurements, and the valency of the element concerned is directly
indicated by the number of attached hydrogen or halogen atoms.
The formulae for these oxides, however, are seldom truly molecular.
Phosphoric oxide, for example, is (P 2 O 6 ) 2 even as vapour, and the
crystal unit of silica is (SiO 2 ) 3 ; but little, as a rule, can be said
about the molecules of solids. Moreover, according to the above
statement, valency as a property cannot strictly be judged from
oxides at all, for these compounds do not exhibit, attached to a nuclear
1 Halide = fluoride, chloride, bromide, or iodide.

VALENCY AND CHEMICAL CONSTITUTION 59

atom, a number of peripheral atoms corresponding with its valency.
Nevertheless, there is good reason to regard the valencies indicated
by oxides such as those in the table to be correct.

The establishment of the idea of valency was soon followed by
a device by which the facts of atomic union were represented
graphically.

Bonds were introduced by Couper to show the joining together
of the atoms in the following way:

Cl-H, H-O H, N, H-C H.

H H

Thus graphic or constitutional formulae were constructed, with
bonds to show units of valency, or units of affinity, which they
might be called, if they are thought of as standing for the forces
by which the atoms are united.

Oxides containing bivalent oxygen have been represented by
graphic formulae, such as the following:

Mg=O, B O B, O=C=O, \N O-N/ , O=S=O,

Vox o> NO j

or
O=B-O B=0,

O O O

O=C1-0-C1=O, O=0s=0.

o & A

Formulae such as these are chiefly of historic interest, for they
have to be reconsidered carefully in the light of modern knowledge
and theory concerning the atom. It will appear later that Couper's
bonds ought not to be used indiscriminately or similarly for all
these compounds; e.g. while they are appropriate in the case of
carbon dioxide they are hardly proper in the case of magnesium
oxide, since the mode of chemical union in this case seems to be
different from that in the case of the gaseous oxide.

When bonds were first employed it was thought that the atoms
in all compounds were united together in a similar way; and all
that has been done in constructing these formulae has been to
arrange the atomic symbols in relation to one another so as to

60 CHEMICAL THEORY

represent known or supposed facts of chemical constitution, and
then to join these symbols by bonds to represent the supposed
acting valencies of each atom. How artificial such formulae are is
seen by comparing the two formulae given for B 2 3 . Each satisfies
the requirement that boron be trivalent and oxygen bivalent,
yet both cannot truly represent the constitution of this oxide.

Is, then, the writing of graphic formulae merely an interesting
geometrical exercise based on the facts of valency alone? Con-
sider, for example, a substance with the formula C 3 H 6 O. Since
carbon is quadrivalent, oxygen bivalent, and hydrogen univalent,
two graphic formulae are possible for this substance:

H O H H H H

H-C-C-C-H and H-C C-C=O.

H (i) H H H (ii)

Does it matter which formula is adopted? The answer is that
two quite different substances are known, both of which are
C 3 H 6 O; and that one of them, acetone, certainly has the constitution
(i), whilst the other, propaldehyde, as certainly possesses the con-
stitution represented by (ii).

Thus graphic formulae are constitutional formulae, and only so
far as they represent the ascertained constitutions of compounds are
they valid ; therefore the construction of graphic formulae for com-
pounds which have not been definitely proved to have certain
constitutions is to be deprecated.

A fundamental question connected with this subject was raised
by comparing magnesium oxide with carbon dioxide. This question
has lately assumed great prominence, and may be illustrated more
clearly by considering the two chlorides NaCl and CC1 4 . Until
recently these compounds have been formulated thus:

Na Cl; Cl C Cl;

there are, however, great differences between the two chlorides,
both in physical and chemical properties, which suggest different
modes of union of their constituent atoms. Sodium chloride is a
solid whose separate atoms 1 of sodium and chlorine are arranged

1 Or, more accurately, ions.

VALENCY AND .CHEMICAL CONSTITUTION

Fig. 8

like cubes 1 packed close together (fig. 8),
until cubic crystals of various sizes are
produced; whilst carbon tetrachloride con-
sists of self-contained CC1 4 molecules which
easily separate from one another, and exist
apart in the state of .vapour. Further,
when sodium chloride dissolves in water
its chlorine becomes reactive in a manner
in which the chlorine of carbon tetrachloride
never becomes reactive; i.e. it shows the reaction of chloride with
silver nitrate which carbon tetrachloride fails to show. Can it
reasonably be maintained that the same kind of bond unites chlorine
with sodium as unites this element with carbon? It cannot; and
consequently there appear to be two kinds of valency; the kind of
valency which exists in carbon tetrachloride can be properly repre-
sented by "bonds", whilst the kind which unites sodium with
chlorine cannot. Therefore the use of Couper's bonds in expressing
the constitution of many inorganic compounds is being discontinued.
The modern view regarding the mode of union of the elements in
these compounds will appear later.

Variability of Valency.

Early in the development of the theory of valency the question
arose whether valency is a fixed and inherent property of an atom,
like its mass, or whether it can vary under varying circumstances.
Kekul^, who showed the quadrivalency of carbon, believed valency
to be unalterable; and the study of carbon compounds alone ap-
peared to justify Kekul^'s opinion. The following compounds were
cited by Kekul6 to illustrate the constant quadrivalency of carbon:
H H Cl Cl

H C-H, H-C-C1, C1-C-C1, H C Cl, O=C<f ,
II I I X C1

H H Cl Cl

O=C=O, S=C=S, H-C-N.

Frankland, on the other hand, observed that nitrogen formed not
only NH 3 , in which the element is evidently trivalent, but also
NH 4 C1 , in which it was apparently quinquivalent. Thus was
expressed the idea of a maximum potential valency, and an actual
valency, exercised in specific compounds, which might be less than
this. And it was observed that the actual valency frequently fell

1 It must not, however, be supposed that the atoms are cubical in shape.

62 CHEMICAL THEORY

short of the potential valency by two units, as, for example, in the
pairs of compounds

NH 3 , NH 4 C1; P 2 Og, P 2 O 6 ; SO 2 , SO 3 ; SnClj, SnCl 4 ;

so it was supposed that when valency decreased from the maximum
it was always by two units, and that consequently the valency
of an element remained either odd or even. Then it was thought
that the two valencies which remained disengaged in the lower
compounds satisfied each other, so that no valencies remained free.

There are, however, some notable exceptions to this supposed
rule, and it cannot be regarded as a natural law. Examples of
these exceptions are shown in the sets of compounds

NO, N 2 3 , NO., N 2 6 ; C10 2 , C1 2 O 7 ; IO 2 , I 2 O 6 ; FeCl 2 , FeCl 3 ;
InCl, InCl 2 , InCl 3 ; WC1 6 , WC1 6 .

The lower compounds are unsaturated, and combine with oxygen,
chlorine, &c., to form higher compounds.

When rise of temperature causes dissociation it thereby causes
the acting valency of the nuclear atom or atoms of a compound
to diminish. Tlyis when ammonium chloride, NH 4 C1, in which the
nitrogen atom is regarded as quinquivalent, dissociates into ammonia
and hydrogen chloride, the nitrogen atom becomes in consequence
trivalent. Similarly tungsten hexachloride, WC1 6 , dissociates at high
temperature into the pentachloride WC1 6 and chlorine. Occasion-
ally dissociation involves the halving of molecules, as the following
examples show:

Fe 2 01 4

A1 2 C1 6 A1C1 3 + A1C1 3 .

Presumably this dissociation involves a reduction in operative
valency, since such valency is necessary to hold together the two
parts of the double molecule, but becomes inoperative on dissociation.
It may be added that, unless there is reason to the contrary,
association into double molecules, such as those represented above,
is supposed to be effected by means of 1 unit of valency. Mole-
cular association in liquids and solids must also be accounted for
by the exercise of additional valencies. Thus liquid water contains
double molecules, or molecules of even higher complexity; and the
existence of these complex molecules is accounted for on the older
theory of valency by assuming oxygen to be quadrivalent, thus:

VALENCY AND CHEMICAL CONSTITUTION

since oxygen is known to be quadrivalent in some other compounds.
The existence of double salts and salts with water of crystal-
lization cannot be explained by the narrower conceptions of valency.
Consider, for example, potassium alum, K 2 S0 4 A1 2 (S0 4 ) 8 24H 2 0.
The constitutional formulae for potassium and aluminium sulphates
have been constructed thus:

K CK /O-S0 2 -<

| = 8 >S02 and ^o-sg <

but it is difficult to see how these formulae are to be united together,
and 24 molecules of water to be incorporated in the scheme as well.

At one time it was customary to describe such compounds as
"molecular" rather than "atomic", but such a distinction is no
longer regarded as valid, and several theories have been proposed
to account for the constitution of these compounds on the ground
that auxiliary or latent valencies come into play in their formation.
These theories cannot, however, be considered here.

The variation of valency with the kind of compound formed
has been illustrated in the lists of hydrides, halides, and oxides
already given. Thus it appears that whilst the valency of an
element towards oxygen and the halogens 1 may rise as high as
8, valency towards hydrogen is never greater than 4; no single
atom is known to combine with more than 4 hydrogen atoms.
Now hydrogen and oxygen are reciprocally related, and it is a
noteworthy fact that as the valency for hydrogen diminishes in
a series of elements with increasing atomic weight, the valency
for oxygen correspondingly increases, and the sum of the oxygen
and hydrogen valencies remains equal to 8. This is shown in the
following compounds, although fluorine and bromine fail to form
oxides, and iodine is not known to form the oxide I 2 7 ; moreover,
tin has recently been shown to form an unstable hydride, though
this has not been proved to be SnH 4 .

CH 4 C0 2
NIT 3 NA
OH 2
FH -

SiH 4 SiO 2
PH 3 PA
SHo SO 3
C1H CIA

QeH 4 GeO 2
AsH 3 As 2 O 6
SeH 2 SeO 3
BrH -

SnH 4 1 SnO 2
SbH 3 Sb 2 5
TeH 2 TeO 3
1H dA)

These phenomena have a deep significance, which will appear later.
With regard to valency for the halogens, it must be noted that

i The halogen elements are fluorine, chlorine, bromine, iodine.

64 CHEMICAL THEORY

as a rule halides are not so stable as the corresponding oxides. For
example, NC1 3 is so unstable as to be highly explosive, whilst N 2 O 3
does not split off oxygen; PC1 5 dissociates into PC1 3 and C1 2 , whilst
P 2 6 is stable; S0 2 may be united with oxygen to form S0 3 , whilst
SC1 4 , formed below C., easily loses chlorine.

Fluorides, however, are much more stable than the other halides:
PF 6 and SF 6 are stable gases, and the existence of OsFg, 1 in addition
to OsF 6 and OsF 4 , shows a valency of 8 towards a halogen.

The Double Bond in Carbon Compounds.

Consider the two hydrocarbons ethane, C 2 H 6 , and ethylene, C 2 H 4 .
The former is saturated, the latter is unsaturated; that is to say, it
is capable of combining with 2 more hydrogen atoms or their
equivalent. This state of unsaturation of ethylene is represented
by a double bond, the graphic formulae for the two compounds
being

Ethane. Ethylene.

H H H H

-C C-H and H-C=C-H.

The question may be asked whether the double bond is simply
employed to keep up the appearance of the quadrivalency of carbon,
or whether it has any real meaning; whether, indeed, carbon is not
really trivalent in ethylene, so that the formula might as well be

H H

H-C C H.

This question may be answered in the negative for several reasons.

First, no such compound as CH 3 CH 2 is known, in which
one atom of carbon is quadrivalent, whilst the other is trivalent; so
that both atoms must be either saturated or unsaturated. Here, at
least, the idea that the two unsaturated atoms in ethylene satisfy
one another appears justified; and the double bond expresses their
mutual dependence.

Further, the double bond between carbon atoms, the ethylene
linkage, as it is called, expresses something more than unsaturation;
for the nature of this union differs from that represented by the
single linkage. It is weaker than the single linkage, for when

^uff and Tschirch, Ber. t 1913, j6 t 929.

VALENCY AND CHEMICAL CONSTITUTION

a compound contains a chain of carbon atoms in which there is
a double linkage, this is the point at which the chain breaks when
the compound comes under disruptive influence. The fact that
the double is weaker than the single linkage shows that no
mechanical significance must be attached to bonds.

There is still a third characteristic of the double bond, which,
however, can only be made clear by the study of the stereo-
chemistry of carbon compounds.

"Chemistry in Space."

How far, it may be asked, is the graphic formula

H
H C-H

supposed to represent the real configuration of the molecule of this
simple hydrocarbon, methane? The answer may at once be given
that it is probably an imperfect representation of the truth, because
it is a flat formula, a formula in two dimensions, whereas matter
exists in three dimensions; the formula has length and breadth,
but the molecule of methane has thickness as well as length and
breadth. Moreover, the adequacy of the formula may be tested
in a very 'simple way. The formula suggests that there might be
two methylene chlorides, CH 2 C1 2 ,

Cl Cl

H C H and H C Cl,

in which the two chlorine atoms are respectively opposite and

adjacent to each other. Two such chlorides

do not, however, exist; therefore a method

of formulation must be found which does

not suggest their existence. Only when

the valencies of the carbon atom are

equally distributed in tridimensional space

is this requirement met; that is to say,

when they are directed from the centre to

the angular points of a regular tetrahedron,

thus:

(D60)

Pig 9

66 CHEMICAL THEORY

Since this figure is symmetrical, the positions of the 2 hydrogen
and 2 chlorine atoms in methylene chloride shown upon it may
be interchanged in any way without causing a difference in the
relative positions of these 4 atoms. This conception of the
disposition in space of the valencies of the carbon atom, which
is due chiefly to van 't Hoff, has been very fruitful in organic
chemistry. The aspect of the science thus suggested has been
called "chemistry in space", or stereochemistry. Space-formulas
should, of course, be applied to all chemical compounds, and some
progress has been made with elements other than carbon; but these
formulae are mainly of use in elucidating the structures of carbon
compounds, where the question of constitution is of such vital
importance.

It may be added that double and triple bonds are represented
stereochemically by the joining of two tetrahedra along their
edges and adjacent surfaces respectively. For example, ethylene,
CH 2 ~CH 2 , and acetylene, CH~CH, are thus represented:

I'lg. 10

The Criterion of Valency.

The facts recorded in the foregoing pages suggest that valency
might furnish a means of chemical classification of the elements,
were it not that the exercise of this property varies somewhat
irregularly. On the other hand, an independent classification of
the elements might be expected to furnish information regarding
valency. Such information is supplied by the Periodic Classifica-
tion, which will shortly be studied. It will be sufficient to state
here that in this classification the elements fall into nine groups
Groups to VIII; and that the maximum valency of each element
appears to be identical with the number of the group which
contains it. Thus, the no- valency elements of the argon family
are in Group O, the univalent metals of the alkalis in Group I, the
bivalent metals of the alkaline earths in Group II, and so on.

VALENCY AND CHEMICAL CONSTITUTION 67

Very seldom does the acting valency of an element exceed that
indicated by the group to which it belongs; nevertheless in
Group IB copper forms CuCl 2 and gold AuCl 8 ; more often, how-
ever, it falls below it. For example, the halogens belong to the
seventh group, and should therefore have a maximum valency of
seven. This is realized by chlorine in C1 2 7 , and by iodine in
H 6 IO 6 ; but not by fluorine or bromine. Iron, nickel and cobalt,
as well as osmium, ruthenium, &c., belong to the eighth group; but
whilst the two latter metals realize octivalency in OsO 4 and RuO 4
the three former metals appear never to be octivalent.

Nevertheless the Periodic Law is the true criterion of the
valency of an element. This will appear later when atomic struc-
ture is considered.

The Nature of Valency.

A study of the operation of valency, however detailed, or the
graphic representation of the union of elements in chemical com-
pounds by the use of bonds or solid geometrical figures, leaves the
nature of valency itself quite unexplained. It may be said that
the force which binds the atoms together is chemical affinity; but
this explains nothing, and, moreover, the term " chemical affinity "
has received a meaning in physical chemistry which is not closely
associated with the idea of units of valency acting in specific
directions through space.

More than a century ago H. Davy 1 expressed the opinion that
" electrical effects are exhibited by the same bodies, when acting on
masses, which produce chemical phenomena when acting by their
particles". Berzelius extended this idea in his electro-chemical
theory, whence is derived the method of classifying the elements
as electro-positive and electro-negative. Faraday, later, showed
that during electrolysis a definite quantity of matter is always
associated with a definite quantity of electricity, a fact which
suggests that electricity as well as matter is atomic. This sug-
gestion starts a trail which might be followed into all the intri-
cacies of modern knowledge and theory concerning the structure of
the atom. The purpose of this chapter, however, has now been
fulfilled; but when the earlier development of the periodic law
has been considered in the next chapter the way will have been

i H. Davy, Phil. Trans., 1807, 1.

68 CHEMICAL THEORY

fully prepared for an excursion into this new domain; and the
promise may be made that in the course of this adventure the
" nature of valency " will become illuminated in such a remarkable
way that an entirely new and impressive conception will be gained
concerning it.

SUMMARY

VALENCY. The valency of an element indicates the number of
other atoms with which one of its atoms can directly combine.

CHAPTER IV
CLASSIFICATION OF THE ELEMENTS

The Periodic Law according to Mendeteeff

When the elements are regarded collectively, and in view of
their ascertained atomic weights and properties, two considera-
tions present themselves: (i) How may the elements be classified?
(ii) What is their origin? These considerations are related, for the
classification of material species is likely to lead to questions
regarding the origin of such species.

Probably the first systematic classification of the elements was
derived from the electro-chemical theory of Berzelius, to which
reference has already been made. This theory grew out of the
facts of electrolysis. Thus, for example, when an electric current
passes through an aqueous solution of sodium chloride, the sodium
appears at the cathode or negative electrode, and the chlorine at
the anode or positive electrode. Consequently sodium was regarded
as electro-positive, being attracted to the electrode of opposite sign,
while chlorine was, for a similar reason, electro-negative. Or, more
generally, metals were considered to be electro-positive and non-
metals electro- negative. Further, it was recognized that some
metals are more electro-positive than others, power of metallic
replacement being regarded as a criterion of electro-positiveness.
For example, since zinc displaces copper from copper sulphate in
solution, zinc is more electro-positive than copper; and, conversely,
since chlorine displaces iodine from potassium iodide in solution,
chlorine is more electro-negative than iodine.

So this method of classification served not only for the distinc-
tion of metals from non-metals, but also for the recognition of
metallic and non-metallic intensity.

When the atomic weights of a sufficient number of the elements
had been established with some degree of accuracy, it was perhaps
inevitable that numerical relationships should be sought for between

70 CHEMICAL THEORY

them, and that attempts should be made to discover a connection
between the properties of an element and its atomic weight.

The first attempt to establish numerical relations between the
atomic weights was made in 1815-6 by an Edinburgh physician
named Prout, who tried to prove that all the elements are con-
densations of hydrogen as the primordial substance, by affirming
that all the atomic weights are whole numbers when that of
hydrogen is unity. This affirmation was unjustified at the time,
for Berzelius subsequently showed that a number of atomic weights,
determined with accuracy by the use of material ordinarily avail-
able, were far removed from whole numbers. Nevertheless the
fact remained that when the atomic weight of oxygen is made
equal to 16-00 "the atomic weights tend to approximate to whole
numbers far more closely than can reasonably be accounted for by
any accidental coincidence"; 1 and therefore it appeared, even a
quarter of a century ago, that the complete rejection of Prout's
hypothesis was unwarranted. Recent work, of which an account
will be given in the sequel, has gone far to re-establish Prout's
hypothesis; which, however, considering the time of its promul-
gation, must be regarded as a philosophic guess rather than a
conclusion of inductive science.

Another attempt was made by Dobereiner, in 1817 and 1829,
who showed that in various triads of related elements the central
member of each group possesses properties and an atomic weight
which are approximately the mean of the properties and atomic
weights of the extreme members of the triad. These triads are:
lithium, sodium, potassium; calcium, strontium, barium; phosphorus,
arsenic, antimony; sulphur, selenium, tellurium; chlorine, bromine,
iodine.

It will be sufficient to give numerical details for the first and
last of these triads.

Element.
Lithium

Atomic Weight.
6-94

Differences.

Mean of Extreme
Atomic Weights.

Sodium

23-00

16-06
i f if\

23-02

Potassium

39-10

It)' 1U

Chlorine

35-46

Bromine

79-92

44-46

81-19

Iodine

126-93

47-01

******

It will be observed that the atomic weight of sodium is almost

1 R. J. Strutt, Phil. Mag. [vi], 1, 311 (1901).

CLASSIFICATION OF THE ELEMENTS 71

exactly the mean of the atomic weights of lithium and potassium,
but that the atomic weight of bromine is considerably less than the
mean of the atomic weights of chlorine and iodine. The relations
suggested by these triads are therefore approximate only. It has
been objected, moreover, that triads should not be made up to the
exclusion of other related elements; that, for example, there are
four halogens, and that it is arbitrary to exclude fluorine by form-
ing a triad with the other three. But when it is recognized that
fluorine differs from the other halogens, not only in atomic weight
relationship, but also widely in the properties of its compounds,
this objection loses force. So that without doubt the relationships
shown by Dobereiner's triads are remarkable; nevertheless their
value is historic only, for they are now merged in the generalization
known as the periodic law.

Another kind of triad was, however, observed by Dobereiner,
in which the three related elements had nearly identical atomic
weights. These triads are:

Iron 55-84
Cobalt 58-94
Nickel 58-69

Ruthenium 101-7
Rhodium 102-9
Palladium 106-7

Osmium 190-8
Iridium 193-1
Platinum 195-2

They also find a place in the periodic classification. Dobereiner's
observations were limited to the elements cited above. These
observations could not give rise to a generalization, since they were
concerned with only a minority of the elements; the majority did
not form triads; and therefore it is difficult to see what significance
could have been attached at the time to the existence of these
triads.

Strecker, in 1859, initiated the idea of seeking relations between
the elements placed in atomic weight sequence; whilst de Chan-
courtois, in 1862, placed the elements in sequence in a spiral round
a cylinder divided into sixteen equal sectors to represent atomic
weight magnitudes. Thus analogous elements of low atomic weights
fell into places in vertical columns because properties recur in such
elements after atomic weight differences of 16.

In 1863-6 J. A. R Newlands arranged the elements in ascend-
ing order of their atomic weights, commencing with hydrogen,
thus:

H Li Be B C N O

F Na Mg Al Si P S

Cl K Ca Cr Ti Mn Fe, <kc.

72 CHEMICAL THEOEY

In this way he discovered that the eighth element is " a kind of
repetition of the first ", the ninth a repetition of the second, and so
on; Na, for example, is a repetition of Li, Si of C, Cl of F. This
discovery he called the

Law of Octaves. "Members of the same group of elements
stand to each other in the same relation as the extremities of
one or more octaves in music."

This simple "law" did not apply to the elements of higher atomic
weight; even in the above table manganese is wrongly classified
with phosphorus; and it was suggested by a contemporary of
Newlands that it would be as useful to arrange the elements in
alphabetical order as in the order of their atomic weights! Never-
theless, the law of octaves is valid as an introduction to the
periodic law. In the year 1869 Mendeteeff arranged all the
elements in the order of their atomic weights, and discovered
a periodicity in their properties. The fact of this periodicity he
enunciated in the following statements:

1. The elements arranged according to the magnitudes of their
atomic weights show a periodic change of properties.

2. Chemically analogous elements have atomic weights either in
agreement (Pt, Ir, Os), or increasing by equal amounts (K, Rb, Cs).
(Of. Dobereiner's triads.)

3. The arrangement according to atomic weights corresponds
with the valencies of the elements, and to a certain extent the
difference in chemical behaviour, for example: Li, Be, B, C, N, O, F.

4. The elements most widely distributed in nature have small
atomic weights, and all such elements are distinguished by their
characteristic behaviour. They are thus typical elements, and the
lightest element, hydrogen, is therefore rightly chosen as the typical
unit of mass.

5. The magnitude of the atomic weight determines the properties
of the element, whence in the study of compounds regard is to be
paid not only to the number and properties of the elements and
their mutual action, but to the atomic weights of the elements.
Hence the compounds of S and Te, Cl and I, show, beside many
analogies, yet striking differences.

6. It allows the discovery of many new elements to be foreseen;
for example, analogues of Si and Al, with atomic weights between
65 and 75.

7. Some atomic weights will presumably experience a correction;

CLASSIFICATION OF THE ELEMENTS 73

for example, Te cannot have the atomic weight 128, but 123 to 126.

8. From the table new analogies between elements become
apparent. . . .

Some of these statements are open to criticism or require modi-
fication. Thus regarding statement 4, lithium and beryllium are
not so widely distributed as the heavier metals sodium and potas-
sium, and magnesium and calcium respectively, and conversely
among the heavier metals tin and lead are more widely distributed
than is the lighter germanium. Also the difficulty in statement 7
has been overcome recently, but not in the way suggested by
Mendeteeff.

Nevertheless these generalizations marked a great advance on the
position of earlier chemists, and Mendel^eff, whilst acknowledging
the work of his predecessors, rightly claimed that he was the
first " to foretell the properties of undiscovered elements, or to alter
the accepted atomic weights ".

The Periodic Law, according to Mendel^eff, may therefore be
stated thus:

The physical and chemical properties of the elements and their
compounds are periodic functions of the atomic weights; or

If the elements are arranged in the order of increasing atomic
weight, their properties vary definitely from member to member of
the series, but return to a more or less similar value at fixed points
in the series.

The periodic system, according to Mendel^eff, may now be
developed.

Hydrogen, the element of lowest atomic weight, became the
sole member of Series 1 in Mendeteeff's system. Series (2) and (3)
were the same as in the octaves of Newlands, thus:

(2) Li Be B C N O F

(3) Na Mg Al Si P S Cl.

The next two series were:

(4) K Ca Sc Ti V Cr Mn F c N -

(5) Cu Zn Ga Ge As Se Br,

being linked together by the triad Fe, Co, Ni; for to place these
three elements in the consecutive positions occupied by Cu, Zn, Ga,
thus displacing all that follow them, would have been not only to
obliterate periodicity from the scheme, but also to ignore the

LONG PERIODS

PQ I

t * i

< (55 1

_l CO
< Q

<u c l
O <n 1

Q. OC
> Ld
I QL

<tf d i

o 1

1 1

U, U

e -d |
N U I

o en

3 bp i
U <3 '

Z 0.

2 S 1

M
P.

U CO

(3 S 1

U
I-H

D
O

,...._.-.

i <

<U 3
PH (^ 1

<u bo
OQ g

1 1 1

ctf
J 2

*- 5 1

u 2 1

^ &

X 2

< <
5H I

a 111

^ U. O,''''

H N U

/
1 (2

& > J

I 1

u c/5 (3

1 *

M ^ Q

1 I

^ 4

I 1

CO
Q.

D
O
DC
O

oc
cp

CO
CO

CLASSIFICATION OF THE ELEMENTS

peculiar relations these three elements bear to one other as members
of a triad.

Series (6) and (7) connected by another triad were:

(6) Rb Sr Y Zr Cb Mo

(7) Ag Cd In Sn Sb Te I.

Ru Rh Pd.

Now it will be seen that Series (4) and (6) begin with the alkali
metals K and Rb, whilst Series (5) and (7) begin with the metals
Cu and Ag, which, whilst allied to each other, differ widely from
the alkali metals. Similar differences exist between subsequent
members of odd and even series. Elements in vertical columns
constitute groups; of which, according to Mendeteeff, there
were eight: seven groups corresponding with Newlands's
octaves, and an eighth group in which Dobereiner's triads of
nearly equal atomic weight were placed. When the inert gases
were discovered, these were placed in a group by themselves:
Group 0, which preceded the other groups. Except with regard
to the elements of Series (2) and (3), Groups I to VII were sub-
divided into A and B Sub-groups, to show the above-mentioned
differences between consecutive members of the same group. Thus
the complete periodic system took the following form.

Groups

III

VII

VIII

Sub-

A B

groups

Series 1

Fe Co Ni

Hu Ilh Pd

Rare

___

Earth

Metals

_ _

Os Ir Pt

Tli

Oxides

X 2

XO 2

X0 3

XO 4

The above arrangement was improved upon, and a clearer
view obtained, by recognizing the existence of short and long
periods. Thus Series (2) and (3) constituted short periods; Series

76 CHEMICAL THEORY

(4) and (5), with the linking elements of the eighth group, formed
one long period. Other long periods followed, and the whole scheme
shown in the chart on p. 74 resulted.

The great advantage of this mode of presenting the Periodic
System was that the A and B Sub-groups were separated, so that
elements which have little resemblance to one another were not
classified together. For instance, it may well be objected that
Cu, Ag, and Au, being very unlike the alkali metals, should not
be placed with them in Group I. This objection is sufficiently
answered when it is shown that these metals occupy positions
near the centres of the long periods, whilst the alkali metals
are quite differently situated at the beginning of these periods.
Similar remarks apply to the relation between manganese and
the halogens.

The arrangement of elements in any group now takes this
form, illustrated by Group I:

(A) Na (B)

K Cu

Kb Ag

The table on p. 74 represents the final and most useful form of
the Periodic System according to Mendeleeff, but before proceeding
further it is desirable to point out some of its shortcomings, and
thus to give a hint of the modification the system has necessarily
undergone on account of recent knowledge.

Mendeleeff did not classify the metals of the rare earths. For
one reason the number of these was unknown, and for another their
properties did not progress from member to member as did the
properties of elements in the recognized periods. Therefore the
position of the rare-earth metals in the scheme could not be given
in detail; but it was indicated that they intervened between Ce in
Group IV and Ta in Group V.

Whilst the rare-earth metals could not be spaced, there remained,
nevertheless, a large number of blank spaces following these elements;
and in course of time it became increasingly improbable that these
spaces ought to be reserved for elements hitherto undiscovered. It
was scarcely credible that if 18 elements indicated by blank spaces
existed not one of these should have been discovered. So it was

CLASSIFICATION OF THE ELEMENTS 77

proposed to fill these blank spaces with Ta and the elements that
follow it, moved up from the series below, so making one very long
period including the rare-earth elements.

Thus A. Werner 1 suggested a long period of 33 elements from
caesium to the higher analogue of xenon, now known as radon; and
except that his estimate of the number of existing rare-earth
metals was one too many, time has proved that Werner was
right.

It must be observed, however, that on account of the rare-eartii
metals, progression of properties from member to member is not
shown throughout this long period; but that these metals may be
regarded as functioning as a single element in the same sense as
the individual members of the triads in the eighth group function
together as a single element.

Atomic Weight Differences in the Periodic System.

Attention may now be drawn to atomic-weight differences
between analogous elements in consecutive short and long periods.
These differences are shown for a number of the elements in the
following tables:

THE Two SHORT PERIODS

He Li Be B C N O jb'

Ne Na Mg Al Si P S Cl

Differences 16-2 16-06 15-30 16-15 16-06 17-02 16-06 16-46

THE FIRST Two LONG PERIODS

Ar K Ca Sc Ti V Cr Mn Fe

Kr Rb Sr Y Zr Cb Mo Ru
Differences 42-99 46-34 47-56 43-8 42-9 42-14 43-99 45-86

Co Ni Cu Zn Ga Ge As Se Br
Rh Pd Ag Cd In Sn Sb Te I
Differences 43-97 48-01 44-31 47-03 45-08 46-1 46-81 48-3 47-02

It will be observed that the differences in the short periods
are approximately 16, and in the long periods about 45; in the
short periods 8 elements intervene before a recurrence of properties,
and in the long periods 18 elements. The differences are by no
means constant, for no mathematical relations exist between the
atomic weights; but anomalies are seen in the differences between

i Ber. y 1905, S3, 914.

78 CHEMICAL THEOEY

krypton and argon, palladium and nickel, tellurium and selenium,
in accordance with the anomalies in the atomic weights of argon
nickel and tellurium, to which attention will be drawn. It must
be confessed, however, that there are other anomalies which are
not pronounced enough to affect the order of the atomic weights
of the elements.

Stress, however, must not be laid upon atomic- weight differences,
because atomic- weight values themselves are now known to be of
only secondary importance in matters of theory and classification of
the elements. Indeed all the anomalies in atomic-weight relations,
whether they affect the relative positions of the elements in series
or not, are now removed, because the conception of atomic number
(q.v.) has displaced that of atomic weight as of primary impor-
tance.

The elements of the short, or so-called typical periods, may be
allied to those either of the A or the B Sub-groups. In the case of
Group I, Li and Na are plainly related to the other alkali metals
K, Rb, Cs, in the A Sub-group, rather than to Cu, Ag, and Au
in the B Sub-group, but in Group VII, F and Cl are related to
Br and I in the B Sub-group, rather than to Mn in the A Sub-
group. This latter relationship obtains in all groups from II to
VII.

The periodic law, according to Mendeleeff, states that the physical
and chemical properties of the elements and their compounds are
periodic functions of their atomic weights. This statement must
now be illustrated.

Periodicity of Physical Properties.

Perhaps the most obvious property of a solid element is its
density. It was shown by Lothar Meyer, in 1870, that the
densities of the elements vary periodically. Instead, however,
of using the densities of the elements directly, L. Meyer calculated
from them the atomic volumes, and plotted these values on a curve
as ordinates, together with the atomic weights as abscissae. The
atomic volume of an element is related to its density in the
following manner:

The reciprocal of the density is the specific volume,

Specific volume = = j = volume of unit mass :

density

CLASSIFICATION OF THE ELEMENTS

1 F

S 3 IN fllOA O I W OJL V

fiO CHEMICAL THEOKY

the atomic volume is this value multiplied by the atomic weight
thus:

Atomic volume = fttomic weight
density

For example, the atomic weight of copper is 63-6, and its
density 8-9; consequently

Atomic volume Cu = ?1? = 7-15.
8*9

This figure stands for the relative volume of a mass of copper
proportional to the atomic weight of the element; it does not
express the relative size of the copper atoms themselves; it could
only do this if the atoms were packed without interspaces, or if
the interspaces were constantly related in volume to the atomic
material of the elements. What it does express is the relative
volume of the atom plus its share of atomic interspace. The
atomic volume curve shows a remarkable periodicity; for it is
like a series of waves consisting of crests and hollows; moreover,
the crests of successive waves increase in height with increasing
atomic weight. The most important fact connected with the curve,
however, is that related elements occupy analogous positions upon
it. For example, the alkali metals, potassium, rubidium, and
caesium, are at the apices of successive curves, the halogens,
chlorine, bromine, and iodine, are on ascending, and the alkaline
earth metals, calcium, strontium and barium, on descending parts
of the curves.

The following other physical properties of the elements and
their compounds are periodic.

Melting-point, malleability, coefficient of expansion, atomic
refraction, conductivity for heat and electricity, colours of salts
in solution. Consequently there are certain regions, which are
similar on successive curves, where these properties are manifested,
or reach their maxima. The student may test this statement with
reference to the melting-points of the elements.

The periodic occurrence of colour in compounds is very striking.
Thus all the metals whose salts give coloured solutions are included
in the following series:

Ti V Cr Mn Fe Co Ni Cu

Mo Ku Kh Pd

W - Os Ir Ft Au
U

CLASSIFICATION OF THE ELEMENTS 81

These series consist of metals in atomic weight sequence, and
they occupy the lowest portions of successive parts of the atomic
volume curve. There are other coloured compounds, however,
which do not give coloured solutions; e.g. various sulphides and
iodides. In the case of these compounds there is generally a
deepening of colour with rise of atomic weight in a group, as in
the sulphides of zinc, cadmium, and mercury. It is noteworthy
that the colours of these compounds belong only to the solids;
for when scarlet mercuric iodide is dissolved in alcohol, and
yellow lead iodide in water, colourless solutions are obtained.
This is to be expected, since colour is not associated with mercuric,
lead, or iodide ions; thus the nitrates of mercury and lead are
colourless, and so are the iodides of the alkali and alkaline earth
metals.

Periodicity of Chemical Properties.

The fundamental chemical division of the elements is into
metals and non-metals; and, according to the classification of
Berzelius, metals are electro - positive, and non-metals electro-
negative.

As the elements are traversed in the order of ascending atomic
weights the variation of metallic and electro -chemical properties
is periodic. Thus in the two short periods from lithium to
fluorine, and from sodium to chlorine, there is continuous and
regular transition from great metallic and electro - positive to
sxtreme non-metallic and electro - negative characters. In the
long periods which follow, for example the period from potassium
bo bromine, there are two phases; the first phase is from potassium
through manganese to the eighth-group metals iron, cobalt, and
lickel; the second phase is from copper to bromine. The transition
Eroin potassium to bromine is similar in degree to that from sodium
:o chlorine, but the period contains more than twice as many
elements; and the stages of this transition present an interesting
Dhenomenon.

The elements of the first phase (K to Fe, Co, Ni) are all
netals, but there is a continuous diminution of electro-positiveness
liroughout them; the elements of the second phase begin with
,he comparatively inert and electro-negative metal, copper, and
Jiere is actually a rise in metallic strength to zinc, followed by
i regular fall to the non-metallic and electro-negative bromine.

(DGO) 7

82 CHEMICAL THEORY

Similar relations exist in the subsequent long periods; but the
inertness of the central elements, i.e. those of the eighth group
and of Group IB, becomes more pronounced with elements of
higher atomic weight.

When the transition of properties within the separate groups,
i.e. the elements in vertical columns, is considered, an increase of
metallic nature or decrease of non-metallic nature is found to be
the rule. Thus, for example, the alkali metals increase in electro-
positiveness with rise of atomic weight; and the halogens similarly
show a diminution of electro-negativeness with rise of atomic
weight. Within the region of chemical inertness and metallic
electro-negativeness, i.e. the eighth group, Group IB, and to a less
extent Group II B, an opposite state of things, however, exists;
there 'is a diminution of electro-positiveness and chemical reactivity
with rise of atomic weight. Thus the inert metals, platinum,
gold, and mercury, occur consecutively as the last members of
Groups VIII, IB, and II B.

From all this it follows that the most powerful metals are to
be found at the extreme left of the periodic diagram; csesium,
the most electro-positive metal, being in the lower left-hand corner;
whilst the non-metals occupy the upper right-hand portion of
the diagram; fluorine, the most powerful non-metal, being in the
upper right-hand corner. The dotted line in the diagram on p. 74
delimits the region of non-metals.

Periodicity of Valency.

The following statement is generally true.

The maximum valency of an element corresponds with the
number of the periodic group to which it belongs. The statement
is illustrated by the formulae of the typical oxides appended to the
table on p. "75. In Chapter III valency was illustrated by lists of
hydrides, halides, and oxides, and in most of the formulas for
the halides and oxides, but not the hydrides, the numerical
value of the valency indicates the group to which the element
belongs. A valency of seven is not always realized in the seventh
group; less often is a valency of eight seen in the eighth group.
On the other hand, copper and gold in Group IB show bi- and
tri-valency respectively in CuCl 2 and AuCl s , but the elements
of this group are in any case somewhat anomalous in their
relationships. A more striking exception is shown in the case of

CLASSIFICATION OF THE ELEMENTS 83

boron, which forms the hydride B 2 H 6 and other hydrides, in which
the element can hardly be less than quadrivalent.

It was seen in the chapter on valency that the sum of the
oxygen valencies and hydrogen valencies in volatile hydrides of
an element is equal to eight; and this is true irrespective of the
periodic group to which the element belongs. The elements of
Groups I, II, and III, however, excepting boron, form no volatile
hydrides, and exhibit only the lower valencies in the oxides. Thus
in the first and second short periods the oxides and hydrides show
valencies as follows:

Ill

VII

Li 2 O

BeO

C0 2

Oxides.

B 2 H fl , <fec

. CH 4

NH 3

OH 2

Hydrides.

Na 2 O

MgO

A1A

SiO 2

S0 3

CIA

Oxides.

SiH 4

PH 3

SH 2

C1H

Hydrides.

Metals which form non-volatile hydrides exhibit the same
valencies in these compounds as in the oxides; for example,
K 2 0, KH; CaO, CaH 2 .

The periodic law constitutes a valuable criterion of valency,
because the periodic group to which an element belongs indicates
almost invariably the valency of the element in the highest oxide
which it can form. The existence of super-oxides, such as NaA
and Ba0 2 , constitutes no real exception to this rule, because these
compounds are constituted thus:

Na O O Na, Ba<^ I, or ^>0:O, Ba=O:O;

and so the valencies of the metals are the same as in the corre-
sponding basic oxides. The elements of Groups VI, VII, and VIII
often fail to realize their maximum valency; and indeed some,
e.g. iron, never exhibit the group valency. Since the chemical
character of a compound depends largely upon the active valency
of its nuclear element, the elements of these higher groups show
great variety in the properties of their compounds, because they
exhibit highly variable valency. The highest oxide is to be
regarded as the typical oxide, provided it exhibits the valency
of the group to which the element belongs; it is then found
that lower oxides and their derivatives show relationships to
oxides and their derivatives of similar type, but belonging to
elements in other groups.

84 CHEMICAL THEOBY

For example:
Derivatives of

Mn 2 O 7 in Group VII are isomorphous with those of C1 2 O 7 in Group VII.
Mn0 3 VII S0 3 , VI.

Mn 2 3 VII Fe 2 O 3

and A1 2 O 3

MnO VII FeO

and ZnO

VIII,
III.
VIII,
II.

Other examples might be given, all of which show that poly-
valent elements, forming several classes of compounds, exhibit
several relationships corresponding to these classes, and therefore
that the type is the determining factor in chemical relationship.
Consequently manganese, which can be septivalent, is not dis-
qualified from appearing in Group VII by reason of relationships
to metals in Groups VIII, VI, III, and II.

Uses of the Periodic Law.

Prediction of Unknown Elements. In the periodic scheme,
as first formulated by Mendel^eff, there were some significant
omissions. The positions now occupied by scandium, gallium,
and germanium were left blank, since no elements were known
qualified to fill them. If every available space had been filled
with the known elements, placed in the order of their atomic
weights, there would have been no periodic system, or but a
distorted one, because every element which now follows a
space that should have been left unoccupied would thereby have
been moved one space forward, and the arrangement of analogous
elements in groups would have been interfered with. On the
other hand, deliberately to leave certain spaces blank, so as to
preserve the desired periodicity, was to suggest that elements
remained to be discovered to fill these spaces, and so to provoke
a severe test of the truth of the periodic law.

The latter alternative was chosen by Mendel^eff, and in par-
ticular the existence of three elements was foretold, which were
named provisionally eka-boron, 1 eka-aluminium, and eka-silicon.
The first of these lay between calcium and titanium in the periodic
table; the other two were placed consecutively to fill two blank
spaces between zinc and arsenic. Moreover, by reference to the
properties of neighbouring elements in series and in group, it

1 Eka is Sanskrit for one.

CLASSIFICATION OF THE ELEMENTS

was possible to foretell with considerable accuracy the properties of
these undiscovered elements.

This prophetic use of the periodic system by its discoverer has
been rightly compared with the employment of mathematical cal-
culation by Adams and Le Verrier to foretell the existence of the
planet Neptune from observed irregularities in the movements of
Uranus, and it has had an equally satisfactory vindication. For
the elements scandium, gallium, and germanium, subsequently dis-
covered, have been found to possess properties closely agreeing
with those foretold by Mendel<5eff. This is illustrated in the follow-
ing comparison of eka-aluminium with gallium.

EKA-ALUMINTUM.

GALLIUM

Atomic weight, cir. 68.

Metal of density 59 and low
melting-point; not volatile; un-
affected by air; should decompose
steam at a red heat and dissolve
slowly in acids and alkalis.

Oxide should have formula E1 2 O 3 ,
density 55, and dissolve in acids to
form salts of the type E1X 3 . The
hydroxide should dissolve in acids
and alkalis.

There should be a tendency
towards the formation of basic salts.
The sulphate should form alums.
The sulpnide should be precipitated
by H ? S or (NBL) 2 S. The anhydrous
chloride should be more volatile
than zinc chloride.

The element will probably be dis-
covered by spectrum analysis.

Atomic weight, 69 9.

Metal of density 5*94; melting at
3015; not volatile; unchanged in
air; action on steam not known;
dissolves slowly in acids and alkalis.

Oxide, Ga 2 O 3 ; density not known;
dissolves in acids, forming salts
GaX 3 . The hydroxide dissolves in
acids and alkalis.

Salts readily hydrolyze and form
basic salts. Alums are known. The
sulphide can be precipitated by
H 2 S or (NH 4 ) 2 S, but only under
special circumstances. The anhy-
drous chloride is more volatile than
zinc chloride.

Was discovered by spectrum
analysis.

There are other blank spaces in the periodic system which were
presumed to correspond with hitherto undiscovered elements.
Modern research, however, has shown that probably only two
spaces remain unfilled, since most of the spaces shown in the table
on p. 74 are obliterated when the table is rearranged according to
recent knowledge.

The two elements required to fill these spaces are: another
alkali metal to precede radium, and a halogen element to precede
radon. A fruitless search has been made for the analogue of
caesium, but the discoveries of the two missing analogues of
manganese and of a rare earth metal have recently been announced.

CHEMICAL THEORY

Correction of Atomic Weight Values.

Since the periodic law requires sequence of atomic weight
values and sequence of properties to be in accord, grossly erroneous
atomic weight value placed in sequence must disturb the sequence
of properties; or, conversely, if sequence of properties is maintained
it will necessitate a departure from atomic-weight sequence. In
either case the erroneous value is revealed when the element in
question is considered in the light of the periodic law. Indeed
the erroneous atomic -weight value must cause a position to be
claimed for the element, which, according to its properties, should
be occupied by another element, and must consequently leave
vacant a place suited to the element and in accord with its true
atomic weight. Therefore the periodic law is of value, not only
for detecting false atomic -weight values, but also for suggesting
true ones. For example, the atomic weight of caesium was at
first erroneously thought to be 123*4. This value would place
caesium after antimony, and, of course, cause the displacement
of tellurium, iodine, and other elements one place to the right.
Such a condition cannot be thought of; therefore the value
123-4 is condemned. On the other hand, since caesium is an alkali
metal it should follow rubidium in group, and consequently have an
atomic weight of about 131-8, so that Cs Rb = Rb K = 4535.
The atomic weight of caesium is now known to be 132-81, and so
this metal occupies its proper place in the scheme. In the cases
of beryllium, indium, and uranium the periodic law has furnished
the means of deciding what multiple of the equivalent is the atomic
weight. The equivalent weight of beryllium (glucinum) is 4-55,
and the atomic weight of this element was at first thought to
be 4-55 x 3 = 13-65. This value would place beryllium in an
impossible position between carbon and nitrogen, whereas 4-55 X 2
= 9-1 would give it a place in harmony with the periodic law.
Subsequent considerations have confirmed the value Be(Gl) = 9-1.

Indium with the equivalent weight 38-27 was thought to have
an oxide InO, and atomic weight 76-54, which would place this
element between arsenic and selenium, whore it cannot stand. An
atomic weight of 38-27 x 3 = 114-8, with the corresponding
oxide In 2 O 3 , would satisfy the periodic law; and this value has
subsequently been accepted on the grounds of specific heat.

The atomic weight of uranium was originally thought to be
about 60, or else 120: but neither of these values enables the

CLASSIFICATION OF THE ELEMENTS 87

element to be placed suitably in the periodic scheme. The value
240 was required by Mendeteeff, so that the element might become
the last member of the sixth group, following tungsten. This high
value, or more accurately 238-2, has been supported by the vapour-
density method applied to the halides, and by the fact that
uranium is radio-active, since radio-activity is characteristic of the
heaviest atoms.

The criticism of the atomic weights of the elements by means
of the periodic law may be carried further. The accepted atomic
weight of argon is greater than that of potassium, that of cobalt is
greater than that of nickel, and that of tellurium is greater than
that of iodine; yet the individual members of these three pairs of
elements are placed in the reverse order of their atomic weights
in the periodic scheme, because their properties do not permit of
any other arrangement. Repeated attempts were made in the case
of tellurium to reduce the value of its atomic weight below that of
iodine, but without avail, and it was supposed that the relationships
of these three pairs of elements constituted exceptions to the periodic
law. Such a conclusion was unsatisfactory; but the difficulty has
been removed by the recognition of atomic number as the criterion
which decides the position of an element in the periodic scheme;
and the atomic numbers of the elements in question place them in
a sequence which accords with their properties.

The Suggestiveness of the Periodic Law.

In spite of the apparent imperfections and anomalies it contains,
the periodic law is true in principle. Indeed it cannot be doubted
that the truth, beauty, and value of this law as an index to the
material world have been enhanced by the discoveries of recent
years. This fact is a challenge to the scientific imagination; it
must provoke questionings and research. For example, in a group
of allied elements, such as the alkali metals, Li, Na, K, Rb, Cs,
there are series of compounds such as oxides, hydroxides,
chlorides, sulphates, carbonates, and so forth, which may be ex-
pected to be related to one another somewhat as the metals
themselves are related. The examination of the physical and
chemical properties of these compounds may therefore be under-
taken with a view to discovering the gradations which exist
between them. Interesting relations will thus be established
and this fact will become apparent: that there is a break in the

88 CHEMICAL THEORY

gradation of properties between Na and K; in other words, that
K, Rb, and Cs and their compounds are closely related, while Na
and its compounds, as well as Li and its compounds, stand apart
from them. The periodic classification affords an explanation of
this phenomenon; it is that Na is situated in the second short
period, whilst K occupies a different kind of position near the
beginning of the first long period, and Rb and Cs follow K in
quite analogous positions in subsequent long periods. Having
observed this, the student may then remember that although
caustic soda and caustic potash are thought of as very similar
substances, sodium salts are after all not very similar to potassium
salts, for they do not crystallize with the same amounts of water
of crystallization as the latter, and frequently they are not iso-
morphous with them, while their solubilities in water are so
different from those of potassium compounds that solutions of
sodium salts are used to precipitate potassium, and vice versa.

At the other extremity of the periodic table the halogens pre-
sent another interesting subject for study. The fact that the
affinity for hydrogen diminishes from F to I in the hydrides HF,
HC1, HBr, HI is well known, and is quite in accord with what occurs
in other groups; e.g. in the hydrides OH 2 , SH 2 , SeH 2 , TeH 2 in
Group VI, or NH 3 , PH 3 AsH s , SbH 3 , (BiH 3 ) in Group V; but fluorine
is widely different from the other halogens. Why is this?

This is a sort of question that must be answered ultimately by
reference to the constitution of the atom; but consideration elicits
this remarkable fact: that all the elements of the first short period
are unique, being widely separated in properties from those in the
same groups which follow them. It suffices to draw attention to
carbon, nitrogen, and oxygen, which cannot be properly classified
with the elements succeeding them. Again, hydrogen fluoride differs
remarkably in condensibility from the other halogen hydrides; is
there any analogy to this phenomenon in neighbouring groups?
Assuredly there is; if water were no more condensible than hydro-
gen sulphide, the world would be a very different place to live in!

In the region of the periodic chart where volatile hydrides
occur the following compounds are found:

CH. NH 3 OH 2 FH

SiBL PH 3 SH 2 C1H

GeH 4 AsH 3 SeH 2 BrH

rSnH 4 ) SbH 3 TeH 2 IH

CLASSIFICATION OF THE ELEMENTS 89

The periodic law suggests a comparison between them in series
and in group; and thus the following gradations of properties are
discovered.

The hydrides diminish in stability with rise of atomic weight in
every group. Thus, for example, in the fifth group ammonia is very
stable, and is decomposed only slowly by the passage of electric
sparks; phosphine, PH 3 , is less stable than ammonia, and is rapidly
decomposed by the same agency; AsH 3 is broken up into its
elements when passed through a tube heated to 230, SbH 3 is
similarly decomposed at 150, and BiH s is too unstable to be
isolated.

In series, i.e. in the hydrides standing in horizontal lines, there
is an increase of stability with rise of atomic weight, corresponding
with the increase of non-metallic characters, and also the diminution
of hydrogen valency, so that there is less hydrogen to be retained.
Thus hydrogen fluoride is the most stable volatile hydride, and ger-
manium and bismuth hydrides the least stable. It may be observed
that Ge, As, Sb, and Bi are metalloids, that is, almost metals. No
true metal forms a volatile hydride. The power to form alkyl
compounds, i.e. compounds with radicles, such as methyl, CH 3 , and
ethyl, 'C 2 H 6 , is more extensive than that to form hydrides; so that
some metals in the B sub-groups preceding in series the above non-
metals form these so-called organo-metallic compounds. Perhaps
the best known of these substances is zinc ethyl, Zn(C 2 H 6 ) 2 ; but,
in addition to zinc, cadmium, mercury, tin, lead, and bismuth form
them, and thus come into line with the above non-metals, all of
which form alkyl compounds as well as volatile hydrides.

Another interesting but rather difficult question is that of the
relative acidic or base-producing power of these volatile hydrides.
Consider the four hydrides:

CH 4 , NH 3 , OH 2 , FH.

Methane is inert; ammonia is base-producing, for its solution in
water is alkaline owing to the reaction: NH 3 + H" + OH' =^
NH 4 * + OH'; water is neutral, and hydrogen fluoride is acid. Why
is not methane, CH 4 , more base-producing than NH 3 ? the grada-
tion of properties seems to require it to be. The answer is that in
CH 4 carbon is already saturated with hydrogen, so that this sub-
stance cannot form an additive compound with water or an acid as

90 CHEMICAL THEORY

ammonia does; for the peculiar base-producing power of ammonia
is an additive property, viz.: NH 3 -f- H* = NH/. 1
Consider again the hydrides:

NH 3 OH 2
PH 3 SH 2 .

There is a loss of base -producing power from NH S to PH 8 , and
an apparently analogous increase in acidity from OH 2 to SH 2 ; but
it is difficult to generalize here, for ammonia is unique in base-
producing power, just as nitrogen is unique as an element; and
water, again, like oxygen, is unique in its properties. Moreover,
it must not be concluded that increase in acidity of hydrides with
rise of atomic weight in a group is general, for C1H, BrH, and IH
are acids of about equal strength.

The comparison of properties of the oxides of elements in the
various groups of the periodic system is a simpler and more satis-
factory exercise. For there is in general a loss of acidic and a
corresponding gain of basic properties with rise of atomic weight
in a group. This is shown, for example, in the oxides

N 2 3 P 2 3 As 2 3 Sb 2 O 3 Bi 2 3 ,
and N 2 O 6 P 2 O 6 As 2 O 6 Sb 2 O 6 Bi 2 O 6 .

In the trioxides there is a gradual transition from wholly
acidic, through amphoteric 2 to purely basic properties, and in the
pentoxides from powerfully to very feebly acidic properties.

Again, the trioxides of Group VI A,

Cr0 3 , Mo0 3 , W0 3 , U0 3 ,

form an interesting series; for, in accordance with the above
generalization, basic properties actually appear, together with
acidic properties, in the oxide U0 3 , which is basic with regard to
one oxygen atom only, forming basic salts, such as U0 2 (N0 3 ) 2 , the
uranyl salts. V

Objections to the Periodic Law.

A consideration of the criticisms to which the periodic system
has been submitted is valuable. If the criticisms are baseless, as
some of them are, the process of their refutation will be illumi-
nating; if they are valid, their consideration may exhibit the

1 The meaning of this will appear when the subject of ionization is considered.
9 Both basic and acjdic, ap<6repo? = both.

CLASSIFICATION OF TfiE ELEMENTS 91

relations of the elements from a new point of view, and so
increase our knowledge concerning them.

The most sweeping accusation which has been brought against
the periodic system is that it places together dissimilar elements,
whilst separating similar ones. It brings together the alkali metals
and copper, silver, and gold in Group I, it is said a most unnatural
alliance. This objection has already been met by a denial of the
statement that these dissimilar metals are brought together. It is
further objected that the periodic classification separates copper
from mercury and barium from lead. But it may be maintained
that such separation is proper; for the similarities between the
metals in these several pairs are superficial rather than funda-
mental, for copper and mercury are widely different in physical
properties and in oxidizability; and, in spite of the fact that both
metals form two series of salts, and that their lower chlorides
are insoluble in water, there is little further resemblance between
their corresponding salts. The differences between barium and
lead are even more fundamental, so that to regard the elements
as similar on account of the insolubilities of their sulphates, and
the isomorphism of some other salts, is a grave error of judgment.

The discovery of argon, and the determination of its atomic
weight, furnished material for adverse criticism of the periodic law.
For not only was it supposed that no room could be found in the
scheme for an element with such extraordinary properties as argon
possessed, but the atomic weight of argon was found to be greater
than that of potassium; and it was manifestly impossible to place
this element between potassium and calcium. Then other inert
elements were discovered helium, neon, krypton, xenon, the
companions of argon; and these have atomic weights less than
those of the neighbouring alkali metals. Thus the atomic weight
of argon is recognized as anomalous, like that of tellurium, and the
inert gases therefore form a new group, which is like a buffer
between the extremely different halogen elements and alkali metals;
just as the metals of the eighth group intervene between manganese
in Group VIlA and copper, silver, and gold in Group IB. So it
is recognized that the elements of the argon family are properly
placed as Group O, the periodic law is vindicated, and, in recog-
tion of their analogy with the noble metals, the elements concerned
are sometimes called the noble gases.

The periodic law, however, needs no vindication. Modern

92 CHEMICAL THEORY

research, it is true, has modified it by causing the conception of
atomic number to displace that of atomic weight; but this has
served only to strengthen the law by removing its anomalies, so
that it has now become the supreme generalization concerning the
origin and constitution of matter as revealed by the inter-relations
of the elements.

SUMMARY

PERIODIC LAW ACCORDING TO MENDEL^EFF. The physical and
chemical properties of the elements and their compounds are
periodic functions of the atomic weights; or

If the elements are arranged in the order of increasing atomic
weight, their properties vary definitely from member to member
of the series, but return to a more or less similar value at fixed
points in the series.

USES OF THE PERIODIC LAW. Prediction of unknown elements.
Correction of atomic weight values. Stimulation of thought and
research regarding the elements.

CHAPTER V
THE MODERN VIEW OF THE ATOM

When the atom was introduced into science by Dalton it
appeared in a theory brought forward to account for the laws of
chemical combination; and since it was supposed to be an ultimate
particle of matter the question of its structure did not arise, for
structure involves parts. Newton had supposed that matter con-
sisted of "solid, hard, impenetrable particles"; and although the
atoms of different elements differed in weight, and presumably
therefore in size, the reason for this was no more an active question
than is the reason for the difference in size of the marbles in a
bag to the boy who plays with them.

The question what the atoms were made of was, nevertheless,
soon raised by Prout; and if it had been conceded that the atoms
were made of hydrogen the great diversity of properties between
the elements must have caused inquiry as to how one primordial
material could give rise to such diversity. This inquiry might
have become more urgent when periodicity of properties amongst
the elements was discovered; but, owing to the suppression of
Prout's idea, no inquiry regarding atomic constitution appears tp
have been made until recently. The generalization of Mendeteeff,
however, that the properties of the elements are periodic functions
of their atomic weights, appears inadequate apart from some idea
regarding atomic constitution. For, consider the two short periods:

Li Be B C N F Ne

Na Mg Al Si P S Cl Ar.

Continuous increase of atomic weight is connected with progressive
change of properties from Li to Ne; but why should this progres-
sive change stop at Ne; why should the addition of about 3 units
of atomic weight to Ne produce an element (Na) which with some
modification reproduces the properties of Li? This fact was long
ago represented by de Chancourtois by means of the " telluric

94 CHEMICAL THEORY

screw ", a spiral curve on which the elements were marked; but
the representation of a fact is a very different thing from its
explanation. There is no explanation, unless the elucidation of
atomic constitution can provide it. So the periodic system
demands a theory of atomic constitution to give it meaning.

The facts of electrolysis investigated by Davy, the electro-
chemical theory of Berzelius, and the laws of electrolysis estab-
lished by Faraday, have some bearing on the constitution of
the atom, though this was not realized by these chemists.
Metals were elements whose atoms could carry positive charges
and travel to the negative electrode or cathode during electro-
lysis; non-metals were elements whose atoms carried negative
charges and in electrolysis travelled to the positive electrode or
anode. Thus elements were distinguished as electropositive or
electronegative, and electricity and chemical affinity were seen to
be closely allied; but these were forms or components of energy
rather than of matter; and that electricity itself could form part
of a material atom was an idea not entertained.

Nevertheless Faraday showed that an ion during electrolysis
was always associated with a fixed quantity of electricity, a
bivalent ion being associated with twice as much electricity as a
univalent ion. This fact is now interpreted as signifying that
electricity, like matter, is atomic, but such a conclusion was not
reached by Faraday. The smallest quantity of electricity associated
with an atom of matter in electrolysis was called by Johnstone
Stoney, in 1874, an electron, and so at that date was recognized as
an atom of electricity.

No real beginning, however, was made towards any knowledge
regarding the constitution of the atoms of matter until these atoms
themselves furnished evidence regarding their internal contents
and structure. The first evidence of this kind was the outcome of
the work of Crookes on high vacua. Crookes found that when an
electric discharge took place through a high vacuum rays travelled
from the cathode in straight lines, that these rays caused the glass
of the containing vessel to fluoresce, but that they were intercepted
by a material object which thus caused a shadow. These rays were
considered by Crookes to consist of matter in an ultra-gaseous
state, and they were subsequently called "cathode rays" or
" cathode particles". Sir J. J. Thomson, in 1897, investigated
these particles, found that they travelled with a velocity about

THE MODERN VIEW OF THE ATOM 95

one-tenth that of light, and proved that their mass was l/1850th
part of the mass of a hydrogen atom. The most significant dis-
covery concerning them, however, was that their nature was in-
dependent of the gas originally present in the vacuum tube, and
of the metal used as cathode. Consequently they were judged to
be not only disintegration products of material atoms, but in-
variable constituents of those atoms. This was the first piece of
evidence regarding the constitution of the atoms of matter.

The next evidence was furnished by the facts of radioactivity,
which began to be discovered after attention had been drawn by
Rontgen to the fluorescence of the glass of the Crookes tube as the
source of those peculiar rays called X-rays.

The radioactivity of uranium, radium, and thorium was found
to be caused by the emission of two kinds of particles known
respectively as a- and /3-particles; and after a time these particles
were recognized to be actually disintegration products of the atoms
of those heavy metals, and were likewise identified. Thus an
a-particle was found to be an atom of helium carrying a double
charge of positive electricity, and a /6-particle to be the same as a
cathode particle which was now also identified with the electron,
the atom of negative electricity.

Yet^a-particles, i.e. positively charged atoms of helium, are not
the smallest known particles of matter, for there are ions of
hydrogen, or hydrogen nuclei, i.e. positively charged atoms of
hydrogen which owe their charge to the loss of an electron. It
has been inferred, without direct experimental evidence it is true,
that the atom of helium has been formed by the condensation of
four hydrogen atoms; i.e. four hydrogen nuclei plus four electrons
have produced one helium atom; and further, that two of the
electrons in the helium atom are detachable from that atom so as
to leave a helium ion or doubly charged helium atom, which is the
a-particle derived from radioactive matter. Incidentally it must
be noted, however, that helium ions are unknown in chemistry,
although they have been recognized in work upon positive rays;
we have no power of removing two electrons by chemical means
from the helium atom so as to produce a helium ion or a-particle.
Moreover, a-particles, ejected from the atoms of radioactive ele-
ments, soon take to themselves electrons, and become helium atoms,
as was shown by Ramsay and Soddy.

A helium atom or ion has never been known to yield hydrogen

96 CHEMICAL THEORY

atoms or ions by disruption, and hydrogen atoms or ions have
never been observed as the products of spontaneous radioactive
change. Nevertheless there is direct evidence that some of the
lighter atoms of matter contain hydrogen nuclei as integral parts
of their structure. This evidence has been furnished by the ex-
periments of Sir Ernest Rutherford, 1 who has shown that hydrogen
nuclei are discharged from the atoms of boron, nitrogen, fluorine,
sodium, aluminium, and phosphorus under bombardment by
a-particles; and it is significant that elements whose atomic
weights are multiples of four, i.e. carbon and oxygen, do not
yield hydrogen nuclei under such treatment. Therefore it is
concluded that the massive parts of those atoms which contain
hydrogen nuclei contain them as such in addition to the requisite
number of helium nuclei. For example: N = 14 = 3He + 2H;
and F = 19 = 4 He + 3H.

Thus a clear idea has been reached concerning the different
parts of which all atoms are composed. They are composed of
hydrogen nuclei, helium nuclei (a-particles), and electrons (/3-par-
ticles); and if each helium nucleus is regarded as reducible to four
hydrogen nuclei, then the atoms of matter consist of hydrogen
nuclei and electrons alone.

Now the electrons are atoms of negative electricity, arid in an
electrically neutral atom these must be balanced, whatever their
number, by an equal number of atoms of positive electricity. These
atoms of positive electricity must be the hydrogen nuclei, which
are the only other constituents of a material atom; they are called
protons^ so that every neutral atom is composed of protons and
electrons in equal numbers.

It is now desirable to discover something concerning the manner
of distribution of the protons and electrons in an atom; and the
first insight into this manner of distribution is gained by consider-
ing and placing in contrast two kinds of change which some atoms
can undergo: chemical change, and radioactive change.

The main characteristic of chemical change in general is that it
can be initiated by man and is reversible, whilst radioactive change
is beyond man's control and, as far as we know, is irreversible. It
is concluded from this and the conclusion is now supported by
strong evidence that chemical change touches only the surface of
the atom, whilst radioactive change affects its internal parts.

i Trans. Chem. Soc., 1922, 121, 400.

THE MODERN VIEW OF THE ATOM 97

That chemical change affects even the surface of an atom is an
idea which would not have been acceptable to Newton or Dalton,
who regarded the atoms of matter as unchangeable. Indeed, a
generation ago this idea would have been thought revolutionary.
Chemical affinity, manifested through valency, was a force exerted
by atoms, but exerted outside themselves; the atoms came un-
scathed through chemical change; they bore no superficial wounds
to show that they had been in action. That chemical change
actually affects and alters the surface of an atom is the idea which
underlies the present-day electronic theory of valency; and in
developing this theory it is well to begin with electrolysis.

It will be remembered that the electric charges upon the ions in
an electrolytic solution are due to definite quantities of electricity
which are the electrons. The question may be asked: whence do
the electrons come which are associated with the atoms of matter in
electrolysis? They are not brought into existence by the current;
they must therefore be derived from the compounds in solution.
Sodium chloride, for example, must contain electrons, which become
available as electric charges when this compound is dissolved in
water. Such a view is consistent with the theory of Arrhenius,
which supposes that when a salt or other electrolyte dissolves in
water it breaks up spontaneously into charged ions, which are
ready to carry or be carried by the current when it comes.

The late Sir William Ramsay represented the electron in a
sodium chloride molecule, and the behaviour of the molecule when
it dissolved in water, in the following manner:

NaECl ^ Na- + ECl'.

Thus an electron, as an atom of the chemical element electricity,
was the binding material between the atoms of Na and 01, but
when the salt was dissolved in water this electron which was previ-
ously shared in common by both atoms became attached solely to
the chlorine atom, with the consequence that the sodium atom by
the loss of negative electricity became positively charged and
functioned as a cation, whilst the chlorine atom, by the gain of
negative electricity consequent on having the electron to itself,
became negatively charged and was the chloride ion. This idea
has now been developed so that the neutral sodium atom is believed
to have on its surface one loosely attached electron which it easily
parts with so as to become a univalent positive ion, whilst the

(D60) 8

98 CHEMICAL THEOKY

neutral chlorine atom is believed to receive easily and accommodate
an electron, thus becoming a univalent negative ion, the chloride
ion. So the chemical union between sodium and chlorine is believed
to consist in the transfer of an electron from each sodium atom to
each chlorine atom.

Such a view, however, necessarily modifies the electrolytic dis-
sociation theory of Arrhenius. Ions are formed, according to the
most recent view, when the compound is produced; for the transfer
of electrons converts neutral atoms into ions. Sodium chloride is
thus always ionized, it would not otherwise be sodium chloride;
but in the solid state the positive and negative ions are held together
by electrostatic attraction, just as two oppositely charged pith balls
are attracted together. When, however, the salt is dissolved in
water, the electrostatic attraction gradually gives way, and the
ions become dissociated. We must not now speak of ionization as
a consequence of solution, for that has occurred already in the
formation of the salt; but electrolytic or ionic dissociation is a suit-
able term to describe the separation, through the medium of the
solvent, of the already existing ions.

The idea of chemical change thus briefly outlined, by which a
salt is produced by the transfer of electrons from metal to non-
metal, clearly suggests that in the act of union the metallic atom
loses part of its substance and the non-metallic atom receives an
addition to its substance; that is to say, the atoms themselves
suffer change in their substance. Thus in contrast with the old
doctrine of the unchangeableness of the atoms of matter we have
the new doctrine that in every chemical change the atoms suffer
change. Nevertheless this change is superficial, and its reversi-
bility depends upon its superficiality.

It is quite different with regard to radioactive change. This
change is believed to affect the innermost recesses of the atom;
that is to say, the nucleus where are situated the protons which,
being hydrogen nuclei, constitute its effective mass. Such a change
is profound; it has not been initiated by man; and when it occurs
it is irrevocable, and so radically alters the properties of the atom
that elemental transmutation is said to take place. Sometimes
only electrons, or ^-particles, are ejected from the atoms of an
element by radioactive change; then the mass of an atom is un-
affected though the properties of the element are altered. If,
however, an a-particle, i.e. a helium nucleus, with an atomic weight

THE MODERN VIEW OF THE ATOM 99

of 4, is cast forth, the atom changes not only its chemical properties
but also its mass, for it becomes an atom having an atomic weight
4 units less.

It is possible, however, to show a little more clearly what is
the effect of the loss of a- and /3-particles by the atoms of an
element through radioactive change.

When a j8-particle is ejected from the nucleus of an element the
predominating positive charge there is increased by one unit, and
the atom as a whole will consequently carry one positive charge
if it was previously neutral; it would therefore become a univalent
positive ion if it did not at once take to itself an electron from
outside. When such an electron is assimilated it does not enter
the nucleus, or the radioactive change would be reversed; it remains
on the exterior of the atom as a valency electron. The negative
valency of the neutral atom, i.e. its power of appropriating electrons,
will consequently be reduced by one. With regard to the periodic
system, then, the loss of an electron from the nucleus of an atom
transfers that atom one place to the right in the table. Examples
of such change, with a corresponding effect on chemical properties,
are known; e.g. radium-B in Group IVB, by losing a /3-particle
from the nucleus of its atom, becomes radium-C in Group VB.
The fact, however, must be emphasized that there is no loss of
mass in this radioactive change, for even the loss of an electron
from the nucleus is compensated for by the gain of an electron at
the exterior. Therefore it appears that two separate elements, as
radium-B and radium-C, judging from the chemical properties,
seem to be, may have the same atomic weights. Such elements
are called isobares. Stress, however, must be laid upon the differ-
ence between tho loss of an electron from the nucleus and from the
exterior or sheath of an atom. The latter is the accompaniment
of any change by which a neutral atom becomes a univalent posi-
tive ion; and it has 110 further .significance.

When an a-particle is ejected from the nucleus of an atom two
kinds of loss are sustained by the atom. The first kind of loss is
that the atom loses two units of positive charge, since the a-particle
is a helium atom carrying two units of positive charge, this particle
being composed ultimately of four protons and two electrons, which
are associated together inseparably, so far as experience goes. Such
a change, when it has been compensated for by the eventual escape
of two electrons from the atomic surface, transfers an element two

100 CHEMICAL THEORY

places to the left in the periodic system, because the positive
valency of the neutral atom, i.e. its power of losing electrons, has
thereby been reduced by two. Various examples of such radio-
active change are known; e.g. the atom of radium in Group HA,
by losing an a-particle, becomes an atom of radium-emanation or
radon in Group O. The second kind of loss sustained by an atom
which ejects an a-particle is a loss of mass. Since an a-particle
is a helium nucleus with atomic weight of 4, 4 units of mass dis-
appear. So whilst the atomic weight of radium is 225*95, that of
radon is 222.

Now the atoms of the heaviest elements are capable of suc-
cessive radioactive changes in which both a- and /3-particles are
ejected. Suppose that an atom loses first an a-particle and then
successively two /3-particles. The loss of an a-particle moves the
element two places to the left in the periodic table, and the loss
of two /3-particles brings it back again two places to the right,
leaving it in the same group it occupied originally. This is the
kind of change uranium, for example, undergoes. Ui in Group VIA
loses an a-particle, with 4 units of mass, and becomes UX 1 in
Group IVA; UX 1 loses a /3-particle, becoming UX 2 in Group VA;
and then UX 2 also loses a /3-particle, becoming Un, which again is
in Group VIA. Thus Ui and Un are both in Group VIA, having
atoms which differ by 4 units of mass, but are chemically indis-
tinguishable. Such elements have been called by Soddy isotopes.

The following conclusion regarding the manner of distribution
of the protons and electrons in an atom has now been reached.

All the protons with some of the electrons arc situated in the
nucleus of the atom; the remainder of the electrons are external
to the nucleus, and some of them are at the extreme superficial
limit of the atom in its sheath, that is to say. These external
electrons are removable by chemical change, but the nucleus remains
intact in all chemical changes, and is affected only in radioactive
change or when submitted to intense bombardment by a-particles
(Rutherford). In a neutral atom the number of protons must be
equal to the number of electrons; therefore the number of electrons
external to the nucleus must* be equal to the excess of protons over
electrons in the nucleus.

It thus appears that there is a number which is both the excess
of protons over electrons in the nucleus of an atom and the number
of electrons external to the nucleus when the atom is uncharged.

THE MODERN VIEW OF THE ATOM 101

This number characterizes the atom as regards its chemical pro-
perties; it is called the atomic number. It does not depend upon
the number of protons in the nucleus; for the loss from the nucleus
of four protons and two electrons, constituting an a-particle, together
with two more electrons, leaves the atomic number the same and
the chemical properties the same, as, for example, was shown in the
case of the uraniums cited above. Nevertheless diminution in the
number of protons involves diminution in atomic weight. Since,
therefore, atomic weight may alter while chemical properties remain
identical, the generalization that the properties of the elements are
periodic functions of their atomic weights no longer appears strictly
true; for it is apparent that atomic number takes precedence over
atomic weight so far as chemical properties are concerned.

Now it has already been seen, in the discussion of the periodic
law according to Mendeteeff, that the question of atomic weight is
not paramount in deciding the position which an element is to
occupy in the system. The cases of argon and potassium, cobalt
and nickel, and tellurium and iodine will be recalled. In each of
these pairs of elements the first-named member has a greater atomic
weight than the second. Yet it was agreed to place the elements
in the order named because their properties demanded this in spite
of their atomic weights.

If every element is to have an atomic number, these numbers
cannot all be decided without reference to the total number of
existing elements from hydrogen to uranium. The idea of number-
ing the elements is not new; it was entertained by Newlands, who,
however, could not carry his law of octaves very far because he did
not recognize gaps in the procession of the elements. Mendeleeft'
recognized the gaps but did not lay stress upon atomic number.
No difficulty now arises in numbering the elements until the rare-
earth metals are reached; for the only vacant place previous to
these metals has recently been filled by the discovery of an analogue
of manganese. A difficulty has existed, however, regarding the
number of rare-earth metals.

That difficulty has now been overcome by the discovery of an
experimental method of determining the atomic number of an
element. This method is the result of the work of Moseley on
X-ray spectra. 1 It was discovered by Moseley that these spectra

1 For an account of this work a special textbook should be consulted, e.g. T/te Structure
of Matter, by Dr. J. A. Cranston.

102 CHEMICAL THEORY

are much simpler and more regular than the luminous radiation
spectra of the elements, since the principal lines of the X-ray
spectra of successive elements follow one another in regular grada-
tion like a flight of steps, so that a missing element would bo
revealed by a gap in the series of spectra. Moreover, the vibration
frequency V of the principal line in the X-ray spectrum of an
element is connected with the atomic number N by the following
formula, where A is a constant:

V = A(N - I) 2 .

Moselcy's results may be summarized thus:

1. Every element is characterized by an integer N, which
determines its X-ray spectrum.

2. This integer N, the atomic number of the element, is
identified with the numerical value of the charge of positive
electricity on the atomic nucleus.

3. The order of the atomic numbers is the same as that of the
atomic weights, except where the latter disagrees with the order
of the chemical properties.

It is interesting to report that the atomic number of uranium,
the last of the elements, is 92; thus there are 92 elements in all,
from hydrogen to uranium. Of these ninety have already been dis-
covered, and two remain to be discovered; these are a halogen to
follow iodine in group, number 85, and an alkali metal to follow
caesium in group, number 87.

Recent discoveries are the elements masurium (Ma) number 43,
and rhenium (Re) number 75, analogues of manganese, and the sole
remaining rare-earth metal illinium (II) number 61. Element 72,
called hafnium (Hf) is now recognized to be an analogue of
zirconium.

Further, the placing of argon and potassium, cobalt and nickel,
tellurium and iodine in their accepted order according to chemical
properties rather than atomic weights has been justified; for the
atomic numbers of these elements are: 18 and 19, 27 and 28, 52 and
53 respectively.

Since the atomic weight of an element must now be subordinated
to its atomic number, it follows that the statement of the periodic
law according to Mendeleeff must be modified.

Thus the statement that the properties of the elements are
periodic functions of their atomic weights becomes:

THE MODERN VIEW OF THE ATOM 103

The properties of the elements are periodic functions of their atomic
numbers,

This is the periodic law in its modern form.

Further, since the order of sequence and total number of the
elements are now definitely known, it is possible to develop the
periodic system itself in a form likely to be permanent. These are
the questions to be decided: (i) how many periods are there; (ii)
how many elements are there in each period?

Already a partial answer has been given to these questions; for
Mendel^eff showed short and long periods containing, if the inert
gases are included, 8 and 18 elements respectively; and, as was
stated on p. 77, Werner, in 1905, proposed a long period of
33 elements to include the rare earths. Rydberg, in 1897, 1 had at-
tempted a classification of the elements according to numbers derived
from their atomic weights, and Rydberg's system, corrected by means
of Moseley's atomic numbers, now furnishes a very simple formula
to express the numbers of the elements in successive periods.

The atomic numbers of the inert gases are these:

He Ne Ar Kr Xe Rd
2 10 18 36 54 86;

therefore the successive periods, which are completed by these
elements, contain the following numbers of elements:

2, 8, 8, 18, 18, 32
= 2 [I 2 , 2 a , 2 2 , 3 a , 3 2 , 4 2 ].

The successive periods or series represented by these numbers
are called the Rydberg series; and although no ultimate explana-
tion of these numerical relationships has yet been given, they are
taken to express the manner of arrangement of the elements in
series in the periodic system. There are thus six series or periods,
with the beginning of a seventh, containing the elements from
87 to 92. So the modern periodic system, according to the
Rydberg series, takes the form on p. 104.

If a clear idea of atomic number has now been gained, this will
furnish a more complete conception of the nature of isotopes. It
will be remembered that isotopes, according to the observations
and definition of Soddy, were elements which, having made
excursions into different groups of the periodic system, on account

*Zeit. phys. Chern., 1897, 14, 66.

<
o

O
O
O

w
PJ

Qoo

ssg

) <M

3-3

ee n

020

104

THE MODEEN VIEW OF THE ATOM 105

of radioactive changes in their atoms, were accommodated in the
same place in the system, either permanently, or else only tem-
porarily because further radioactive changes removed them from
that place. It will now be seen that since isotopy relates to
elements having the same atomic number irrespective of their
atomic weights, it is not necessarily limited to radioactive elements.
Indeed there is no a priori reason why the phenomenon should not
occur widely throughout the whole range of the elements. If, how-
ever, it did so occur, the phenomenon would result from different
atoms of what is chemically the same element with the same
atomic number having different atomic weights because of different
numbers of protons in their nuclei.

The idea that different atoms of the same element may have
slightly differing relative weights is not new. It was put forward
by Crookes in 1888, with reference to yttrium, in the following
words: "The atomic weight which we ascribe to yttrium therefore
merely represents a mean value around which the actual weights
of the individual atoms of the 'element' range within certain limits.
But if my conjecture is tenable, could we separate atom from atom,
we should find them varying within narrow limits on each side of
the mean."

If this possibility is admitted for yttrium, we cannot refuse to
consider it for other elements, as indeed Crookes realized. That
the atomic weight of any of the elements represents not the weight
of every atom of that element, but the mean weight of an un-
numbered host of those atoms, is an idea which evidently has some
connection with the subject of isotopes; but how can such an idea
be put to the test, and how can it be regarded as any other than
an unprofitable speculation? A banker in pre-war days would
weigh a hundred sovereigns instead of counting them, because he
knew the average weight of such a number to be constant Yet
the sovereigns might have been weighed one by one on a delicate
balance, and differences in their weights detected. So the chemist
weighs many atoms of an element together, and finds the average
weight of the same number always the same. To detect differences
in individual weights, however, he would need to weigh the atoms
separately; but that he cannot do.

Yet an instrument called a mass-spectrograph has been devised
by which the atoms of an element are separated in such a manner,
when charged electrically, that they register themselves on a photo-

106 CHEMICAL THEORY

graphic plate in positions which depend only on their individual
masses.

Sir J. J. Thomson began work upon this subject in 1912, carry-
ing out what was called positive-ray-analysis, because the " rays ",
now called " mass rays ", which produced the effects were positively
charged particles or ions. Thus Thomson separated gaseous neon,
with an atomic weight of 20-2, into atoms, most of which were
shown to have a relative weight of 20, and a much smaller number
a relative weight of 22. So it was demonstrated that the element
neon is a mixture of isotopes, its accepted atomic weight being the
mean of the atomic weights of the separate isotopes present in
the requisite numerical proportions, viz. 90 per cent of Ne 20 and
10 per cent of Ne 22 .

After the war Dr. F. W. Aston developed the method of
Thomson and elaborated the instrument, and thus has been able
to show by means of " mass spectra " that a large proportion of the
chemical elements are mixtures of isotopes. Up to the end of 1924,
56 elements had been examined by Aston and others, and of these
25 were found to consist of identical atoms, and 31 of mixtures
of isotopes.

Now the elements, all of whose atoms are identical in weight,
are also elements whose accepted atomic weights approximate very
closely to whole numbers, whilst among the elements which are
mixtures of isotopes are those whose atomic weights are far
removed from whole numbers. Examples of the former are:
C = 12-00, N = 14-01, O = 16, S = 32-06, P = 31-02, Cr = 52-00;
and of the latter: Mg = 24-32, 01 - 35-46, Cu = 63-57, Zn = 65-38,
Se = 79-20, Kr = 82-92, Hg = 200-60.

It is obvious, however, that the possession by an element of an
atomic weight which is approximately a whole number is no proof
that the element does not consist of isotopes; for the mean of a
number of isotopes might happen to be nearly or exactly a whole
number. Such is the case not only with Kr = 82-92 but
also with Br = 79-92, which is a mixture of the isotopes Br 79
and Br 81 .

The question here arises how many isotopes of an element there
may be, and what range of atomic weight, or mass number as it is
now called, is possible. The answer seems to be that 8 is the maxi-
mum difference in mass number, and therefore 9 the maximum
number of isotopes possessed by any element. Thus the following

THE MODERN VIEW OF THE ATOM

107

data for tin and xenon, as well as for potassium and copper, are
given by Aston (Chem. Soc. Ann. Report, 1924):

Element.

Atomic
Number.

Atomic
Weight.

Minimum
Number
of Isotopes.

Mass Numbers of
Isotopes in order
of Intensity.

118-70

7(8)

/120, 118, 116, 124, 119,
\ 117, 122, (121)

130-20

7(9)

J 129, 132, 131, 134, 136,
I 128, 130, (126), (124)

39-10

39, 41

(Ju

63-57

63, 65

The greatest numbers of isotopes are possessed by elements of
even atomic number; indeed elements of odd atomic number, e.g.
potassium and copper, seem to consist of not more than two isotopes
whose mass numbers differ by two units.

It appears that isobares are present amongst the above isotopes;
e.g. Sn 124 and Xe 124 , provided the latter value is substantiated. Yet
although some of their atoms have equal masses, tin and xenon are
entirely distinct elements; for whilst the number of protons in the
nuclei of their atoms may be the same, the numbers of electrons,
and hence of electric charges therein, must differ. Other examples
of isobares are furnished by Ar 40 , Ca 40 and Ge 74 , Se 74 . A striking
fact in connection with the subject of isobares is that tellurium
with mass numbers 120, 130, 126, and atomic weight 127-5, or
more probably 127-8, shares all those numbers with xenon, with
which it is thus trebly isobaric. Iodine, however, with atomic
weight estimated to be 126-92, consists of I 127 only.

The mass numbers of atomic isotopes are always given as whole
numbers, whilst the estimated atomic weights of the elements are
often fractional. Thus is raised a question which it is necessary to
discuss. The mass of an atom is due to its protons, and a proton is
a hydrogen atom minus an electron, which thus has a mass of
1-008 when O = 16-00.

The oxygen atom contains 16 protons, yet its mass is not quite
16 times the mass of a proton. This loss of mass is attributed to
a " packing effect " in the nucleus, where the additive law of mass
is not obeyed. On the basis of O = 16-00, however, the atomic
masses of the isotopes are known to conform to the whole number
rule except for a few small variations.

Thus since the masses of the individual atoms of all the elements

108 CHEMICAL THEORY

are, within a close approximation, whole numbers when O = 1600,
it is clear that the fractional atomic weights with which we have
been long familiar, and which we are still compelled to employ in
accurate analytical work, are averages due to mixtures of isotopic
atoms; and so we understand why the atomic weight of an element
which does not exhibit isotopy is very nearly, if not precisely, a
whole number.

A further question connected with isotopy is this: If the same
element is found in different parts of the world, will it always
have the same estimated atomic weight? The idea that the same
element, obtained from different sources, may have differing atomic
weights is a disturbing one which strikes at the foundation of all
accurate analytical work. For example, a redetermination of the
atomic weight of antimony has resulted in a drastic change from
1202 to 121 .76. Can it be that different Isotopic mixtures of
antimony atoms have yielded those discordant results? This is
believed not to be the case; but it is thought, rather, that the new
figure is a correction of the old. Indeed there is abundant evidence
that the atomic weights of naturally occurring elements which
are not of radioactive origin are always constant whatever the
sources of the elements. Thus cobalt and nickel of meteoric origin
have the same atomic weights as the terrestrial elements, and the
same is true of silicons from cosmic and terrestrial sources. Never-
theless it is highly desirable for those who undertake the redeter-
mination of atomic weights to state the source of their material.

So far as present knowledge goes, the experimental atomic
weights of normal inactive elements are still to be regarded as
constants of nature, since nature has effectively mixed her isotopes,
and never sorts them out again. Nevertheless the proved existence
of isotopes is a challenge to man to separate them. Chemical
methods of separation are unavailing, since isotopes do not differ
in chemical properties; but physical methods, such as fractional
diffusion and distillation, have been attempted with some success
in the case of chlorine, mercury, and perhaps zinc (Ckein. Soc.
Ann. Report, 1922).

The case is very different with an element which is the residue
of radioactive change. The uranium atom, for example, with atomic
weight 238-17, passes through a succession of radioactive changes,
in which it loses eight a-particles as well as /3-particles, the final
product being lead. This lead, therefore, should have an atomic

THE MODERN VIEW OF THE ATOM

109

weight of 238-17 32 = 206-17. Now uranium minerals are
found to contain a small proportion of lead, and the lead extracted
from such minerals has been estimated to have an atomic weight
of 206*46, whilst that of ordinary lead, not associated with radio-
active material, is invariably found to be 207-20.

Again, the thorium atom, with atomic weight 232-15, loses six
a-particles in the series of radioactive changes which end in lead;
therefore the lead derived from thorium would be expected to have
an atomic weight of 232-15 - 24 = 208-15. Lead obtained from
thorite has been found experimentally to have an atomic weight
of 207-77; which, although a little lower than that it would be if
the lead were derived from thorium alone, is considerably higher
than the atomic weight of ordinary lead.

Uranium-lead and tlioriuin-Iead are indistinguishable from
ordinary lead in chemical properties and in all physical properties
except density. Thus three leads have been named which, chemi-
cally speaking, are one lead. These, with their experimental
atomic weights and densities, arc:

Uranium-lead.

Thorium-lead.

Ordinary Lead.

Atomic weight
1 )ensity
Atomic volume

206-08
11-213
18-28

207-77
11-376
18-26

207-20
11-352
18-25

It is seen, moreover, that the densities of these leads vary as their
atomic weights, so that their atomic volumes are constant.

There is more to tell about the internal structure of an atom.
Since an atom consists of a nucleus and surrounding electrons, it is
desirable to gain some idea of the size of the nucleus as compared
with that of the atom as a whole. The experiments of Rutherford
on the scattering of a-particles yield the desired information. These
experiments have already been referred to because they furnish
information regarding the description of the atoms of some of the
lighter elements; but if attention is concentrated on the tracks of
the a-particles themselves rather than on the havoc they work by
their bombardment, some quite different information is obtained.

When the a-particles from a radioactive source traverse a gas
supersaturated with aqueous vapour, their tracks can be made visible
by the condensation of the vapour which occurs along them. Thus
it is discovered that whilst some of the a-particles undergo sharp

110 CHEMICAL THEOBY

deflections, these deflections are many times fewer than they would
be if collision with an atom as a whole caused deflection. So it is
concluded that a large proportion of the a-particles pass through
the atoms as through empty space, and that only when an a-
particle collides with or comes very near to a nucleus is it deflected.
Thus it has been calculated, owing to the rarity of these collisions,
that the diameter of the nucleus of an atom is about one ten-
thousandth part of the diameter of the entire atom.

On account of this relation of the nucleus to the surrounding
electrons, an atom of matter has been compared with the solar
system, the nucleus being analogous to the sun, and the electrons to
the planets.

The different planets of the solar system have different orbits,
and they revolve round the sun. Have the planetary electrons of
an atom different orbits, and do they revolve round the nucleus in
their several orbits? This is a question to which some answer
must now be given, though the details of the answer are not yet
beyond the region of controversy.

The number of electrons external to the nucleus of an atom of
an element is the same as the atomic number of the element; and
thus with elements of high atomic number these electrons are
numerous. In the atom of uranium there are 92 electrons external
to the nucleus, and it is not to be supposed that these are all
situated or revolve on the surface of a single shell or envelope.
Rather must it be supposed that the electrons are distributed in a
number of shells which succeed one another like the layers of an
onion. How many electrons are present in each layer is a question
on which in some cases there is difference of opinion. It must be
remembered, however, that since the electrons on the outer layer
of the atom, which is called the sheath, are the valency electrons,
their number will be related to the valency of the element; and
further, that since the atom of an inert gas has no valency, the
sheath of such an atom will presumably consist of a completed
layer of electrons, to or from which no electron can be added or
removed.

To form a mental picture of the structure of the atoms of
matter it will be best to begin at the beginning, with hydrogen.
The neutral hydrogen atom consists of 1 proton + 1 electron, and
the helium atom of 4 protons + 4 electrons, 2 of these electrons
being bound up in the nucleus with the 4 protons, and the other

THE MODERN VIEW OF THE ATOM

111

two being in a sheath which is complete since the helium atom
manifests no valency, although it can exist momentarily without
these two electrons as an a-particle ejected from a heavy atom
during radioactive change. A consideration of the hydrogen atom
under different circumstances will illuminate the subject of valency.
When a hydrogen atom becomes a cation, i.e. the hydrion, as in the
formation, say, of an acid in aqueous solution, this atom parts
with its solitary electron and becomes reduced to a naked proton. 1
It is possible, however, for the hydrogen atom to assimilate
a second electron and so become an anion. This is shown by the
fact that lithium hydride, LiH, yields on electrolysis hydrogen at
the anode. Thus the hydrogen atom shows a tendency either to
lose its single valency electron, or more rarely to gain another, by
which means it would assume an external configuration character-
istic of the uncharged helium atom.

Now consider the short period Li to Ne.

Atomic number

Number of electrons inl
sheath of neutral atom J

Valencyl normal
(Abegg) /contra

-7

+2
-6

+3
5

-3

-2
+ 6

-1

The atomic number of each atom in this period exceeds the number
of electrons in the sheath of the neutral atom by two, because there
are two electrons in the under layer which correspond with the
two electrons in the complete sheath of the helium atom. Now
when the significance of the numbers in the above table is under-
stood, the nature of valency will stand revealed.

It is the great merit of Abegg 2 to have drawn attention to
the fact that the sum of the hydrogen and oxygen valencies of
a number of elements is equal to 8, e.g.

SiH 4 PH 3 SH 2 C1H
Si0 2 P 2 6 S0 3 CIA,

and to have derived therefrom the theory of normal and contra-
valencies, an example of which is given in the table above. Ac-
cording to Abegg, the normal valencies of an element are the more
usual and characteristic; the contravalencies are more seldom

1 Unless it is hydrated. 2 Z. anorg. Chem. (1904), 39, 330.

112 CHEMICAL THEORY

exercised, at any rate with the more extreme members of a period.
Now if, in addition to this, valencies are regarded as positive or
negative according to whether they are exercised towards electro-
negative or electropositive elements respectively, and it is supposed
that the actual exercise of valency implies the loss or gain of
electrons by the sheath of an atom, and further, that every atom
undergoing chemical combination tends to assume the condition of
an inert gas as regards its sheath, then the following ideas regard-
ing the valencies of the elements of the first short period follow.

Lithium, with 1 electron in its sheath, can assume the external
condition of an inert gas either by losing 1 electron, so as to
simulate the helium atom which precedes it, or by gaining 7
electrons, so as to have a sheath identical with that of a neon atom.
It is, however, much easier for an atom to lose 1 electron than gain
7; hence lithium invariably manifests a valency of +1 by becoming
a cation carrying one positive charge, rather than a valency of 7,
that is, an anion carrying seven negative charges.

Similarly beryllium loses 2 electrons on ionization, becoming a
bivalent cation, rather than gaining 6 electrons to become a sexi-
valent anion.

With carbon, however, the chances of losing or gaining electrons
are about equal; and with nitrogen the alternative of the loss or
gain of electrons also exists. Oxygen and fluorine, however, are
too electronegative ever to become cations by losing electrons, the
contravalencies, at any rate in the case of fluorine, being entirely
latent.

The same considerations apply to the next short period from
sodium to argon, the only difference being that the atoms of both
the inert gases, neon and argon, to the external configuration of
which the intervening elements tend to conform when they enter
into chemical union, are both alike in having 8 electrons in their
sheaths.

So far it appears that valency depends on the number of
electrons in the sheath of an atom; and whether that atom exercises
positive or negative valency depends upon whether it more easily
loses or gains electrons, so as to present a completed sheath on its
outer surface.

Langmuir, who, following G. N. Lewis, has developed this idea, 1
carried it further, and applied it to the whole of the periodic

1 J. Amer. Chem. Soc. t 1919, 41, 868.

THE MODERN VIEW OF THE ATOM

113

system, so that whilst neon and argon at the end of the two short
periods have each 8 electrons in their sheaths, krypton and xenon
have each 18, and radon has 32. Thus Langmuir postulates l that
" the electrons in atoms tend to surround the nucleus in successive
layers containing 2, 8, 8, 18, 18, and 32 electrons respectively".

In the long periods, however, it is not possible to connect valency
with the electronic content of the atomic sheath in the same simple
fashion as in the short periods. Consider the period of 18 elements
from K to Kr. This period begins and ends like the preceding
short period; i.e. K, Ca, Sc resemble Na, Mg, Al in valency, and
As, Se, Br similarly resemble P, S, 01; but with the intermediate
elements, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, more compli-
cated relations appear. In one sense the elements K to Mn resemble
a short period; e.g. with regard to the oxides

K 2 O, CaO, Sc 2 O 3 , TiO 2 , V 2 O 6 , CrO 3 , Mn 2 O r .

Thus Mn 2 O 7 resembles C1 2 O 7 , permanganates being isomorphous

with perchlorates, so that a similar arrangement of electrons on

the atomic sheaths of Cl and Mn may be inferred. The oxides

Cu 2 O, ZnO, Ga 2 O 3 , GeO 2 , As 2 O 6 , SeO 3 , ,

again, less perfectly reproduce the relations of the oxides of the
first short period. Moreover, a new phenomenon occurs in the
centre of this long period; this is reducibility of the higher com-
pounds with the loss of single units of valency, and the simul-
taneous appearance of coloured ions.

Thus salts corresponding with the following oxides have coloured

ions:

TiO
Ti 2 3
(Ti0 2 )

(V,0)

V 2 3
V 2 4

CrO
Cr 2 3

Cr0 3

MnO
Mn 2 O 3
MnO 2

MnO 3
Mn 2 O 7

FeO
Fe 2 3

FeO 3

CoO
Co 2 3

NiO

(Cu 2 0)
CuO

In view of these considerations, MendeteefFs division of the long
periods into the elements of the A and B sub-groups may be brought
forward again, thus:

Elements of A Subgroups K
Elements of B Sub-groups ! Cu

Ca
Zn

Sc
Ga

Ti
Ge

Mn
Br

FeCoNi

(DCO)

1 Science, July, 1921.

114

CHEMICAL THEORY

and so it may be pointed out that the triad (Fe Co Ni) appears to
function like a single element, i.e. like an inert gas at the end of a
period; and although Ni cannot be compared with an inert gas,
Pt may, because it is so inert. Thus if the non-valent, inert gases
are regarded as " standards of atomic stability ", Ni, Pd, Pt, or more
accurately, imaginary inert forms of these elements, are sometimes
regarded as " sub-standards of atomic stability ".

Now if Langmuir's postulate is examined in the light of all
these considerations, it will appear difficult to account for valency
by reference to the electrons supposed to be contained in the atomic
sheaths of the neutral atoms. Nickel, for example, would have
10 electrons, and require 8 to complete the sheath; copper would
have 11 electrons, and require 7 to complete the sheath, and so on.
And whilst the existence of Ni(CO) 4 would seem to justify a valency
of 8 for nickel, there are no properties of copper or the succeeding
elements to show such high valencies as would thus be attributed
to them.

Consequently attention may be drawn to other views regarding
the distribution of the electrons in an atom.

In 1921 Bury 1 modified Langmuir's theory by assuming that
the number of electrons in a completed sheath of an atom never
exceeds 8; and in the same year Bohr, from a consideration of the
spectra of the elements (see later), adopted a similar view.

The following table sets forth the atomic structures of the non-
valent gases according to Bohr. The numbers in brackets are the
atomic numbers of the elements, and therefore the numbers of
electrons external to the nucleus.

Orbits, numbered outwards
from Nucleus.

Helium (2)

Neon (10)

Argon (18)
Krypton (36)

2
2

8
8

8
18

Xenon (54)

Radon (86)

. 8

An advantage of this view is that it provides for the addition
of electrons either in the sheath of an atom or in an orbit or shell
below it. When the latter occurs there need be no change o

i/. Amer. C/tem. Soc. t 1921, 43, 1602.

THE MODEKN VIEW OF THE ATOM 115

valency in passing from one element to the next, as is shown, for
example, in the chlorides

VClg, CrCljj, Mild* Fe01 2 , CoCl 2 , NiCl 2 , CuCl* ZnCl*
The idea is specially helpful, however, in accounting for the 14
elements of the rare earths, all of which have the same valency.
Successive additions of electrons are here supposed to be made to
the electrons in the fourth orbit, so increasing these from 18 to 32.

A final question, so far as the present study of atomic structure
is concerned, is that of the activities of the electrons within the
atom. Regarding these activities the views of the physicist and
the chemist appear to be at variance. The physicist believes the
electrons to be revolving round the nucleus in their several orbits
as the planets revolve round the sun. To him an electron at rest
is as unthinkable as a planet in such a condition. The chemist,
however, is well content to think of a stationary electron; indeed
he seems to demand it by his ideas of valency and the constitution
of compounds. How can chemical compounds be formed without
points of attachment between the atoms, and how can points of
attachment be provided by swiftly revolving electrons? It is true
that an electrolyte like sodium chloride might exist; for the chemist
has learned to regard its atoms as held together not by bonds, but
by electrostatic attraction between oppositely charged ions. But
what is to be said about such compounds as methane and the host
of organic substances, concerning whose structure and stereo-
chemistry the chemist has such elaborate and satisfying ideas,
based upon the doctrine of bonds?

The physicist, however, needs to account for those beautiful
phenomena, the luminous spectra of the elements. It used to be
asked: how can the atom of iron vibrate in hundreds of ways at
once so as to give rise to the hundreds of lines in its luminous
spectrum? It is sufficient now to ask how the hydrogen atom,
consisting of one proton and one electron, can vibrate in various
ways so as to produce the various lines in its spectrum. We are
indebted to Bohr 1 for an explanation of this phenomenon, based
on Planck's Quantum Theory of Energy, which now finds general
acceptance.

Imagine an atom with revolving electrons which are radiating
energy into space. If this radiation were continuous, the electrons

1 Vide Tlit Theory of Spectra ami Atomic Constitution, by Niels Bohr : Cambridge University
Press, 1924.

116 CHEMICAL THEORY

would be continually losing energy, and in consequence continually
approaching the nucleus in a spiral path. Moreover, such continu-
ous radiation could not produce a discontinuous line spectrum. To
avoid the nemesis of the atom by the collision of planetary electrons
and nucleus, it is assumed that a revolving electron loses no energy
so long as it remains in a single orbit; that it is only change of
orbit which is accompanied by change of energy; a loss of a definite
amount of energy, the so-called quantum, will thus accompany the
fall of an electron from one orbit to that beneath it, i.e. nearer to
the nucleus; and a corresponding gain of energy will accompany
the restoration of the fallen electron to its former state. It is now
easy to understand, if there are numerous possible orbits, that each
kind of fall gives rise to a particular radiation which produces its
own line in the spectrum; and that the several lines occurring
simultaneously in the hydrogen spectrum are produced by corre-
sponding simultaneous falls from several different orbits in the
peripheries of hydrogen atoms with their electrons in several
different states, although each atom contains only one electron.

The apparently irreconcilable views of the physicist and chemist
may be expressed thus: to the physicist an atom is a hive of
activity, a home of swarming electrons; to the chemist it is an
abode, if not of rest, then of nothing more than vibratory motion of
electrons about their mean positions. Can these views be recon-
ciled? It is possible that they may be if a revolving electron can
be considered to be more in one place than any other, if there is
any point through which it passes very frequently whilst other-
wise tracing out divergent paths. This is impossible if an electron
describes circles in the same plane round the nucleus as centre, or
if its path is a simple ellipse, like the path of a planet, with the
nucleus at one of the foci of the ellipse.

If, however, the motion of an electron is compounded of a cir-
cular or elliptical motion, and a circular motion at right angles to
it, the path travelled will be precessional upon the surface of an
ellipsoid, 1 i.e. it will be represented by a series of curved lines
whose directions are constantly altering so as to cover the whole
surface of the ellipsoid very much as the coloured and twisted lines
on an ornamental glass marble cover its surface. The consequence
of such a motion will be that the rotating electron will pass, during

1 An ellipsoid is the solid figure formed by rotating an ellipse about its major axis, just as
a sphere is the solid figure formed by rotating a circle about its diameter.

THE MODERN VIEW OF THE ATOM 117

one cycle, many times through two points, which are at the ex-
tremities of the major axis of the ellipsoid, but only once through
every other point.

Such a conception, which is due to J. D. Main Smith, provides
for the localization of an electron, as well as satisfying some re-
quirements of the physicist. Whether, however, it will suffice to
account for both the physical and the chemical properties of the
atom cannot yet be said. Meanwhile the idea of stationary elec-
trons within or upon the surface of an atom is so very valuable
a contribution to the theory of chemical structure that it will be
adopted and developed in the next chapter, which deals with the
modern view of the molecule.

CHAPTER VI
THE MODERN VIEW OF THE MOLECULE

The student of chemical history is aware that two views have
been held regarding the structure of chemical compounds. The
first view was expressed in the electrochemical theory of Berzelius,
which postulated electricity as the binding force between atoms, so
that a molecule consisted of atoms held in electrical equilibrium by
mutual attractions.

There were two kinds of electricity, and each atom in a com-
pound possessed some of both kinds, but in unequal quantities, so
that a positive or negative charge preponderated, according to
whether the atom was metallic and electropositive, with a larger
positive than negative charge, or non-metallic and electronegative,
with a larger negative than positive charge. Thus it followed that
every molecule consisted of two parts, a positive and a negative
part; and these parts in turn might consist each of two smaller
positive and negative parts, and so on, down to the individual
atoms. For example, the double salt potassium-alum, apart from
its water of crystallization, would be accounted for somewhat in
this way:

KOSO S A1 2 O 3 3SO 3

+ +

KO S0 3 A1 2 3 3SO 3

+ - + - + - r + ~i

K O S 3 A1 2 3 3 [S 3 J.

This was the dualistic system] and it was successful in account-
ing for the structure of electrolytes, which are polar compounds,
but failed when applied to organic compounds, which are non-
electrolytes or non-polar compounds. Thus if every compound is
composed of positive and negative parts, in equilibrium, what, it
may be asked, are these parts in such a compound as CH 4 ; and
again, if + and parts are balanced in CH 8 COOH, how is it

118

THE MODERN VIEW OF THE MOLECULE 119

possible for electropositive hydrogen to be replaced by electro-
negative chlorine, so as to produce CC1 3 COOH?

In view of questions like these, Dumas propounded a second
view in his unitary system of chemical compounds, in which every
compound formed a complete whole, and did not therefore consist
of two opposite and balanced parts. He thus referred the pro-
perties of a compound to its type rather than to the properties of
its constituent atoms. The consequence was that unitary views
prevailed and dualism was discredited. When, therefore, the
doctrine of valency was developed, graphic formulas with " bonds "
were employed indiscriminately to represent the structure both
of electrolytes and non-electrolytes.

I/the electrolytic dissociation theory of Arrhenius, however,
marked a return in part to dualism; and that theory, together with
the properties of solutions to which it was related, emphasized the
real difference which exists between electrolytes such as sodium
chloride, and non-electrolytes such as chloroform. Since, however,
" bonds " were supposed equally to join the atoms of sodium and
chlorine in sodium chloride, and carbon, hydrogen, and chlorine in
chloroform, simple solution in water involved the breaking of bonds
in one case, but not in the other; and it was difficult to find a valid
reason for such an extraordinary difference of behaviour of differ-
ent compounds as they dissolved in water. ^

v'Now it has been seen that the electronic theory of valency, so
far as it was developed in the last chapter, appears to deal with
valency in electrolytes alone. This is true of Ramsay's idea of
valency, and also of the theory of transference of electrons during
the combination, say, of sodium and chlorine to form sodium
chloride; so that a conception of the molecule is reached resembling
that of Berzelius, because it represents atoms, or, more strictly,
charged ions, as held together by electrostatic attraction. There is
difficulty, however, in applying this simple theory to all molecules;
and to realize the difficulty it is only necessary to consider the
molecule C1 2 . Sodium and chlorine combine because of an electro-
chemical difference between these elements; the sodium atom loses
an electron which the chlorine atom, because of its different
chemical nature, readily takes up. No such reason can account
for the union of two chemically identical chlorine atoms to form
a molecule, so that there cannot be transfer of electrons in such
a case.y

120

CHEMICAL THEORY

This difficulty is met by a conception due to G. N. Lewis and
developed by Langmuir: the conception of covalency, as distinct
from electrovalency, which is the kind of valency hitherto con-
sidered. Now an atom of chlorine has 7 electrons in its sheath,
and requires 1 to complete the octet characteristic of the sheath of
an inert gas. Such an atom, however, cannot gain its required
electron from a similar neighbouring atom, and even if it did it
would become a chloride ion such as does not exist in chlorine gas.
It is possible, however, for two chlorine atoms, with identical
requirements to satisfy these requirements, mutually, by the sharing
of a pair of electrons, each chlorine atom providing one electron of
the pair. The accompanying figure makes this plain, i

/Thus two octets containing only 14 electrons between them, are
possible because two of these electrons are common to both octets.
The shared electrons held in common by both chlorine atoms,
which are shown in the figure within the rectangle, constitute a
duplet. This duplet is a unit of covalency, and is equivalent to a
single valency bond. According to Langmuir, it is the only sort of
bond, and it represents the kind of union which exists between
the atoms of compounds which are not electrolytes. \/

Thus, by this view, there are two kinds of valency: electro-
valency and covalency. The theory of electrovalency is the
modern equivalent of the dualistic theory of Berzelius; that of
covalency corresponds with the unitary theory of Dumas. /

Covalency may now be further illustrated. The molecule O 2
consists of a pair of atoms, each of which separately has six
electrons in its sheath, and therefore requires two electrons to
complete the octet. Two oxygen atoms can combine together to
produce a pair of octets if each atom shares two electrons with its
neighbour, producing a pair of duplets representing a double bond,
O = O, thus:

Similarly, carbon dioxide, O = C = O, can be represented by

THE MODERN VIEW OF THE MOLECULE

121

showing the carbon atom sharing two duplets with each oxygen
atom thus:

o + c + o =

The combination of hydrogen with oxygen, and with chlorine, to
form water and hydrogen chloride respectively, may now be con-
sidered. Since water is a non-electrolyte, the two hydrogen atoms
in each molecule are supposed to unite with the oxygen atom by
covalency, in a manner which may be represented thus:

H + O + H = HOH

Thus each hydrogen atom completes its sheath of two electrons,
and the oxygen atom its sheath of eight.

As regards hydrogen chloride, since in the anhydrous state this
compound is a non-electrolyte, its molecule may be represented
thus, the atoms of hydrogen and chlorine being united by covalency :

In presence of water, however, the covalency bond is broken, and
the hydrogen chloride is ionized, becoming hydrochloric acid. It is
now believed, however, that the hydrogen ion of hydrochloric acid
is hydrated, that, indeed, it is an unstable oxonium ion (OH 3 )'
analogous to the ammonium ion (NH 4 )', and formed thus: 1
OH 2 + HC1 = (OH 3 )' + Cl'.

This change would be represented structurally thus:

HCl

[Cl]'

The explanation of this change of hydrogen chloride from a
neutral, non-electrolytic substance to a powerful acid through the
assimilation of water is to be found in the principle of symmetry.

1 Vide Lowry, Chemistry and Industry, 1923, 46.

122

CHEMICAL THEORY

The molecular condition of non-ionized hydrogen chloride is an
unsymmetrical and therefore relatively unstable one, whilst the water
molecule requires two hydrogen nuclei to be attached to the four
remaining electrons to create perfect symmetry. Having attached
one hydrogen nucleus, however, the water molecule, which has now
become a positive [OH 3 ]" ion, is unable to take up another because,
being positively charged, it repels similarly charged hydrogen nuclei.
Hence the condition of things shown in the figure is arrived at.

Now when hydrogen chloride solution is mixed with ammonia,
neutralization takes place thus:

[OH 3 ]- + [Cl]' + NH 3 [NHJ- + [Oil + H 2 0,

the neutralization being due to the fact that the ammonia molecule
appropriates another hydrogen nucleus to form the highly sym-
metrical ammonium ion with greater force than the water molecule
retains it. Nevertheless, just as [OH 3 ]* is unstable, easily losing a
hydrogen nucleus in presence of the hydroxidion of an alkali to
leave a stable OH 2 molecule, so [NH 4 ]' is also unstable, and in
presence of much hydroxidion similarly loses a hydrogen nucleus
to provide hydrion and form water. That is why ammonium salts
evolve ammonia in presence of alkalis.

The student of organic chemistry is conversant with the theory
that the four valencies of the carbon atom are directed towards the
angular points of a regular tetrahedron; and he knows that this
theory has been most fruitful in elucidating the structure and
stereochemistry of carbon compounds. It is therefore necessary to
inquire how the stereochemical theory is related to the electronic
theory of valency, so far as this applies to the carbon atom.

In the preceding figures the atoms of all the elements except
hydrogen have been represented as having a cubical structure, or
rather, the eight electrons in the completed sheath of an atom have
been placed at the angular points of a cube. Now a regular tetra-
hedron is the hemihedral form of the cube; if, therefore, the eight
electrons draw together into four pairs, two pairs being produced
by movements at right angles to the movements of the other two,
a tetrahedral figure will be produced (see fig., p. 123). Thus it is

THE MODERN VIEW OF THE MOLECULE

123

believed that whilst the cubical form of the atom as regards the

distribution of its electrons is preserved when chemical union is by
electrovalency as in sodium chloride, union by
covalency involves the distortion of the cubical
form into the tetrahedral. So the structure of
carbon dioxide is represented by the following
figure. Thus the tetrahedral model of the
carbon atom is preserved, and union by single,
double, and triple bonds becomes union at an

angle, a side, and a face of the tetrahedron respectively, by one,

two, or three pairs of electrons.

It is troublesome though picturesque to represent atoms by
cubes, and molecules by numbers of united cubes. A simpler plan
is to use the ordinary atomic symbol surrounded by dots to re-
present electrons. Thus the molecule of chlorine, instead of being
represented as in fig. on p. 120, becomes

and other formulae are:
H 2 NaCl

H:H Na :01:

which are equivalent to
H-H Na- Or

C0 2

C 2 H 2
H:C:::C:H,

O=O O=C=O H-C==C-H.

So it appears that a pair of electrons, acting as a unit of covalency,
is equivalent to a single chemical bond. It is to be noted, more-
over, that with the exception of H, which has two electrons, and
Na, which is an ion with eight electrons in a lower layer, every
atom is represented as having or sharing eight electrons, since all
the electrons which bind two atoms together are shared by both
atoms concerned.

The question may be asked whether there is an absolute
distinction between electrovalency and covalency, or otherwise

124 CHEMICAL THEOEY

between polar and non-polar compounds; or whether one kind of
valency merges into the other in a series of compounds, which thus
show transition from polarity to non-polarity. Thus, while electro-
valency is the mode of union in sodium chloride, it may be asked
whether all chlorides, metallic and non- metallic, are constituted
similarly. Now electrovalency must be the mode of union in the
case of the chloride of a univalent atom, for such an atom, e.g. Na,
cannot share the two electrons which are necessary to a covalent
bond; and co valency must be the mode of union when two similar
atoms unite, as in the case of C1 2 , for there is no reason for
electronic transfer.

It may be questioned, however, whether electrovalency is the
mode of union of the atoms in CC1 4 . Rather may it be supposed
that this compound, which resembles methane in inertness, is simi-
larly constituted, in which case the two molecules may be repre-

sented thus:

H :C1:

H:C:H :C1:C:C1:.

H ":C1:"

With SiCl 4 , however, the case is different, for this compound is
distinguished from CC1 4 by its reactivity towards water. Possibly,
therefore, the state of union between the silicon and chlorine atoms
may be represented as something intermediate between covalency
and electrovalency, thus:

:C1:
:C1: Si :C1: ,

the eight electrons which in covalency the silicon atom would share
with the chlorine atoms being drawn towards these atoms, so that
the silicon atom becomes charged positively and the chlorine atoms
negatively, though these charges are not sufficiently free to consti-
tute the compound an electrolyte. This can be understood if it
is remembered that the electrons in the sheath of a silicon atom
are farther from the positively charged nucleus than those in the
sheath of a carbon atom, and so would be held less tenaciously by
the nucleus; or in other words, that the silicon atom is more electro-
positive than the carbon atom, and more ready to shed the electrons
of its sheath.

THE MODERN VIEW OF THE MOLECULE

125

The stability of CC1 4 , as well as of CH 4 , as compared with SiCl 4
and SiH 4 , is probably to be referred eventually, however, to the
inner structures of the atoms of the two elements. Between the
sheath of the carbon atom and its nucleus there is only the shell
of two electrons which constitute the sheath of the helium atom,
whilst in the case of the silicon atom the completed octet which
constitutes the sheath of the neon atom intervenes. Consequently
the eight electrons of the completed sheath of the carbon atom in
a compound such as CC1 4 can assume the tetrahedral position
without strain, whereas the neon octet in the case of the silicon
atom makes tetrahedral symmetry and resulting stability more
difficult of attainment.

The idea of gradations between polar and non-polar molecules
is due to G. N. Lewis. Thus Lewis accounts for the properties of
hydrogen chloride by supposing that the two electrons which are
shared between the hydrogen and chlorine atoms in the anhydrous
molecule may be displaced in the direction of the chlorine atom so
as to cause this to become negatively charged, whilst the hydrogen
atom becomes positively charged to the same degree, but that when
displacement is complete hydrion and chloridion result thus:

H:C1: H :C1: * [H] + -

Langmuir, on the other hand, believes that the electrolytic pro-
perties of hydrogen chloride solution are due to the hydrogen
nucleus uniting with the water molecule to form a hydrated hydro-
gen ion, or oxonium ion, according to the theory already given.

The octet theory that is, the theory that the completed sheath
of an atom contains eight electrons accounts for the frequency
with which four atoms are joined to a central atom, to form a
molecule if the product is electrically neutral, or an ion if it is not.

Examples are the molecules CH 4 and Os0 4 , the cation NH 4 ', and
the anions Si0 4 "", PO 4 '", S0 4 ", MnO 4 ", Mn0 4 ', CIO/, which are
formulated thus:

:6:

:O:

:6:

:0:Si:0:

:O:P:O:

:O:S:O:

:O:Mn:O:

:O:C1:O:

:0:

:O:

:0:

the valencies of the separate atoms, and of the ions, their algebraic
sum, being:

126

Si = +4 P = +5
O = -8 4O = -8

CHEMICAL THEOKY

S = +6 Mn = +6
4O = 8 40 = 8

Mn = +7 Cl = -f 7

SiO 4 = -4 PO 4 = -3 SO 4 = -2 Mn0 4 = -2 MnO 4 = -1 C1O 4 = -1

It is noteworthy that the covalent bond, i.e. the duplet, is the
same whether the uniting element is univalent like H or bivalent
like O. Consequently the practice of representing O as united
with other elements by means of double bonds, unless these are
double covalent bonds, disappears.

The way in which the ions represented above are built up may,
however, be elucidated a little further. Consider the SO 4 ion.
The sulphur atom, when neutral, had a valency of +6 with six
electrons in its sheath, and needed two more to complete its octet;
the four oxygen atoms, when neutral, similarly possessed six
electrons in their sheaths, so that each also required two more
electrons to complete its octet, eight in all being required, of which
the sulphur atom was ready to provide six for covalent union.
Two more electrons were therefore necessary, and these were pro-
vided by two neutral hydrogen atoms, which so became ions.
Thus H 2 SO 4 consists of two hydrogen ions each with one + charge,
because of the loss of an electron, and the S0 4 ion having two
negative charges owing to the gain of two electrons.

Actually, of course, H 2 S0 4 is produced by the union of H 2 O
and S0 3 , both of which are neutral molecules because their com-
ponent atoms mutually satisfy each other in covalent union. In
order that the sulphur octet in S0 3 may be completed, however,
one oxygen atom is represented as united with the sulphur atom

by a double covalent bond, thus: 0=S<^Q, so that the reaction

H 2 + SO 3 H 2 SO 4
becomes:

H:O:H

S:O:
6:

:O:S:O;

THE MODEKN VIEW OF THE MOLECULE 127

It is thus plainly seen that the two extra electrons provided by
the hydrogen are necessary because, owing to the opening out
of the double bond, an oxygen atom fully furnished with an octet
of electrons must be available to convert the molecule S0 3 into
the ion SO 4 .

It has been usual to show the constitution of sulphuric acid by
its derivation from sulphuryl chloride thus:

Cl HOH OH HC1

+ o=s=o +

Cl HOH OH HC1;

this reaction now becomes:
:C1: H:6:H

H::H

- H + :C1:
6:S:6: +

H+ :C1:

Thus the connection between the old and the new way in chemistry
is perceived. In this new way, or something like it, chemical com-
pounds will be formulated in the textbooks of the future.

The methods of representing the constitutions of inorganic salts
according to the older ideas of valency sometimes led to difficulty
on account of the isomorphism of compounds not related chemically.
Thus sodium nitrate, NaNO 3 , and calcspar, CaCO 3 , are isomorphous;
yet they are chemically unrelated, and were given constitutional
formulae to accord with their chemical properties, thus:

in which nitrogen was shown to be quinquivalent and carbon quadri-
valent.

Now in the modern method of formulation the valencies of the
individual atoms in a compound radicle disappear when the octets
of the constituent atoms are completed. Thus the nitrate and
carbonate ions are similar in constitution though they differ in the
electric charges they carry, the two salts being formulated thus:

It is thus an argument in favour of the electronic theory of valency

128 CHEMICAL THEORY

that these formulae accord with the fact of isomorphism, which is
obscured by the older formulae.

Moreover, it now becomes clear that Mitscherlich, who enunci-
ated the law of isomorphism, was right when he stated that
isomorphism depended primarily on the number and mode of
arrangement of the atoms in the molecule of a compound rather
than upon the chemical nature of those atoms.

This chapter on the modern view of the molecule would be
incomplete without reference to the question of the existence of
the molecule in the solid state, although this subject has already
been dealt with briefly under the subject of molecular association
(p. 46).

Molecular formulae have long been used in expressing the
reactions of solids without consideration whether they stand for
realities. There is no harm in their use, provided it is understood
that they represent, on an atomic basis, only the quantitative
relations of reacting substances. Formulae such as NaCl and
CaC0 3 , &c., suggest no more in their common use than the atomic
relations within the compounds they represent; and it is immaterial
for general purposes whether the true molecular formulae of these
compounds should be simple, or, say, (NaCl) m and (CaC0 3 ) n . Never-
theless it is desirable to form a mental picture of a solid, and if
possible obtain a true conception of its molecular state. When
a chemical compound is truly solid it is crystalline, for the amor-
phous state is really the state of super-cooled liquid. The arrange-
ment of the atoms in a crystal is revealed by X-ray spectrography,
according to the researches of Sir W. H. and W. L. Bragg 1 ; and
the method can be applied not only to obvious crystals, but also to
powders, 2 such as precipitated calcium carbonate, which are thus
shown to be essentially crystalline.

In connection with the arrangement of the atoms within the
crystal, the idea of a space-lattice has been introduced. Consider a
piece of wooden trellis- work, which can be opened to show a pattern
of diamond-shaped spaces. Such a pattern, indefinite in extent, but
with units which are similar parallelograms, is a lattice, i.e. a plane
lattice, a lattice in two dimensions; and all that is essential to form
it is two sets of parallel lines, the lines in each set being equi-
distant. Now extend the idea to three dimensions, and let there

1 X-rays and Crystal Structure (G. Bell & Sons).

2 Hull, J. Amur. Cfam. Soc. (1919), jl, 1168.

THE MODERN VIEW OF THE MOLECULE 129

be three sets of parallel planes which intersect. Thus a series of
identical units or cells is produced, each cell being a parallelopiped.
This is a space-lattice. Moreover, the pattern is preserved, whether
in two or three dimensions, if the lines or planes are obliterated,
provided the points of intersection of these lines or planes are
preserved. Such a pattern in three dimensions is a picture of the
disposition of the atoms in a crystal; for the lines, of course, are
only imaginary, and the atoms may be regarded as points in space.
Further, a lattice unit, whether in two or three dimensions, is
a single parallelogram or parallelopiped; similarly there is a crystal
unit or crystal cell, which is the smallest unit in which the essential
properties of the crystal, as regards the space disposition of the
atoms of the substance, are expressed without repetition.

The figure below depicts a crystal unit of sodium chloride as

revealed by X-ray spectrography, white
spheres representing sodium atoms, and
black spheres chlorine atoms. The space-
lattice, of which this is the smallest es-
sential part, is called the face-centred
cube lattice, because an atom in this
case a chlorine atom is at the centre

^ eac ^ ^ ace ^ ^ e cu be. It is imma-
terial whether a sodium or a chlorine
atom forms the face centre, for by
bisecting the cube parallel to a face
and adding to one of the halves another half cube a sodium face-
centred cube would be formed.

It is to be observed, however, that each sodium atom in such
a structure is surrounded by six equidistant chlorine atoms, as is
seen to be the case with the central atom in the figure; and
similarly, that each chlorine atom is surrounded by six equidistant
sodium atoms. The question may therefore be asked: what has
become of the molecule of sodium chloride? and the answer is that
no such molecule exists in the solid salt. Such an answer, more-
over, is quite in accordance with the electronic theory of valency
as applied to sodium chloride. It is ions, however, and not neutral
atoms of sodium and chlorine, which are packed together in solid
salt; and when the salt disintegrates in water these ions wander
freely in the solvent without existing as NaCl molecules.

It is therefore quite impossible to write a molecular formula for

(D60) 10

130 CHEMICAL THEORY

solid sodium chloride, and therefore the simple formula NaCl for
the salt serves every useful purpose.

The electrostatic attraction which binds the ions of sodium and
chlorine together in sodium chloride is the cause why this com-
pound is a solid at ordinary temperature and not a gas.

There are no molecular boundaries, and all the ions in a mass
of the salt are fastened together in one bundle by a pervasive
force or field of electric attraction, which hitherto has been called
cohesion. When, however, such a mass is fused, and so strongly
heated as to be converted into vapour, ions of sodium and chlorine
pair off and exist as NaCl molecules. It could hardly be otherwise.
Oppositely charged ions of sodium and chlorine could remain
separate in the state of vapour only if they possessed such high
velocities on account of elevated temperature that the attractive
force between them was rendered ineffective. That would be
thermal dissociation, such as occurs even when pairs of identical
atoms have been united by covalency, as with I 2 .

In view of the conception of the crystal unit as the smallest
portion of the solid necessary to represent completely the properties
of the crystal, it may be asked whether this unit is identical with
the chemical molecule. It has been seen that in the case of
crystallized sodium chloride no chemical molecule can be said to
exist; yet molecules of organic compounds, that is, compounds
whose atoms are united by covalent bonds, are believed to be
present even when these compounds are solid. Nevertheless, ac-
cording to Sir William Bragg, 1 there is no reason why the crystal
unit should be identical with the chemical molecule; and X-ray
analysis has shown that this unit generally consists of two, three,
or four molecules; e.g. the crystal unit of naphthalene consists of
(C 10 H 8 ) 2 , and that of a-naphthol, where the symmetry is reduced by
the introduction of an OH group, of (C 10 H 7 OH) 4 . Similarly two
molecules of benzene, C 6 H , constitute a crystal unit, but four of
benzoic acid, C 6 H 5 COOH.

Consider the two solid and related elements carbon and silicon.
When carbon burns it forms an oxide which is a gas, but when
silicon burns the oxide formed is a solid. What is the reason for
this difference? The answer is that carbon dioxide consists of
molecules of three atoms compactly joined together by covalency,
each forming a self-satisfied system with very little outside in-

i "The Significance of Crystal Structure", Tram. Chem. Soc., 1922, 121, 2766.

THE MODERN VIEW OP THE MOLECULE 131

fluence, so that only at low temperature and high pressure do the
separate molecules unite to produce liquid and then solid carbon
dioxide; but that it is otherwise with silica. Silica has long been
recognized to consist of polymerized molecules, and it is now
known that as quartz its crystal unit is (SiO 2 ) 3 . It may well be
doubted, however, whether Si 3 O 6 molecules, existing side by side
with little attractive force between them, would be competent to
produce an inert and non-volatile solid like quartz. Bather it
would seem that a cohesive force of great strength exists between
the Si 3 O 6 molecules, to overcome which very high temperature
is necessary; or otherwise that the mode of union pertaining
between silicon and oxygen in silica is analogous to that between
the elements in TiO 2 , SnO 2 , and other related oxides. Indeed,
as compared with other dioxides it is carbon dioxide which has
exceptional properties, not silica; and, as was said earlier in this
chapter, it is the inner structure of the carbon atom that fits it
for covalent union, which then confers volatility on the compounds
carbon forms with other elements.

Nevertheless, it is remarkable that silicon tetrachloride (B.P. 57)
is so volatile compared with silica. This may be because the
silicon atom in the tetrachloride is unable to exercise much attrac-
tive influence on its similar neighbours owing to the cordon of four
chlorine atoms with which it is surrounded; whereas the oxygen
atoms of silica are not numerous enough to exercise this influence.
It may be noted, moreover, that SO 3 , C1 2 O 7 , and even OsO 4 are
volatile.

Compounds such as PC1 5 and SF 6 call for comment. If it is
believed that the halogen atoms are attached to the other atom
in these molecules by covalency, then ten and twelve electrons
respectively are concerned in the process; but it seems unlikely
that the octet of electrons is exceeded in the sheaths of atoms of
such low atomic number as phosphorus and sulphur. The alterna-
tive is to regard the halogen atoms as united by electrovalency, in
which case the five and six electrons originally present in the
sheaths of phosphorus and sulphur atoms respectively will have
left these atoms to become attached severally to the halogen atoms.
In this case the stability and volatility of SF 6 , as of other poly-
fluorides, is to be attributed to the simplicity of internal structure
of the fluorine atoms which allows them to come very close to a
sulphur atom; whilst the dissociation of PC1 6 vapour into PC1 3 and

132 CHEMICAL THEORY

C1 2 , with subsequent oxidation of PC1 3 , may be represented as
taking place in the following way:

.. : 9) : .. :C1: :C1: :C1: :O :C1:

:C1: p :C1: _^ ' + :P:C1:, and :P:C1: -^':O:P:O:;

which shows two chlorine atoms returning an electron each to the
phosphorus atom, and then forming a chlorine molecule by co-
valent union; whilst the other three chlorine atoms also enter into
union with the phosphorus atom, so that PC1 3 becomes a covalent
compound, which may be subsequently oxidized to phosphoryl
chloride in the manner shown.

An alternative view, however, regarding PC1 5 is that two of the
chlorine atoms are united to the phosphorus atom by single electrons

the singlet bond thus:

C1:P:C1,

ci ci

so that the octet rule is not departed from; and there is experi-
mental support for this view.

These illustrations suffice to show the trend of the modern
theory of the molecule; and they leave no doubt that the kinds of
formulae which have embellished our textbooks for a generation
must soon give place to formulae of another kind. The present
position of the theory, however, must be regarded as a phase, for
many of the ideas are speculative; but so long as chemical science
remains alive and active, so long must its theories continue to
undergo modification.

PART II
STATES OF MATTER

p ART HSTATES OF MATTER

CHAPTER VII
THE PROPERTIES OF GASES

The states of matter which are commonly recognized are three
solid, liquid, and vapour or gas. Matter in these three physical states
is composed of molecules, the inter-relations of which determine the
state of the matter. A solid is characterized by a volume which is
but slightly affected by changes in temperature and pressure, and by
a shape of its own, which may be naturally assumed or artificially
induced. By reason of the heat energy they possess, the molecules
or it may be ions of a solid are in motion; but, inasmuch as
the volume of the solid does not tend to change spontaneously,
this motion does not affect the distance apart of the molecules;
and, since the shape of the solid is permanent, the molecules do
not change their positions relatively to each other. The motions
ofJJie molecules of a solid, therefore, are restricted to vibrations
about mean fixed positions, and the molecules themselves are re-
tained in these positions by the exercise of the force of cohesion,
which in solids is at its maximum. Even in some solids, however,
cohesion does not prevent vaporization or diffusion, as in the case
of some solid metals. A liquid, like a solid, has a definite volume,
little affected by changes of temperature and pressure; it has no
definite shape of its own, however, but when placed in a containing
vessel assumes for the time being the shape of the vessel. It is
thus the property of a liquid to flow, to spread itself out, and
adapt itself to external conditions, while maintaining its volume
unchanged; consequently the molecules in their motions must keep
the same mean distances apart while the liquid changes its shape.
Thus the force of cohesion in a liquid, although weakened so that
the molecules are not maintained in a fixed mean position as in a
solid, is yet sufficient to confine these molecules within a definite
volume so long as they remain part of the liquid.

134

STATES OF MATTER AND PROPERTIES OF GASES 135

The characteristic of a gas is its power of indefinite expansion,
so that it will distribute itself uniformly over whatever space is
at its disposal. Thus a gas has neither definite shape nor definite
volume, and it must be confined within an impervious envelope
or it will be lost in space. The latter, of course, is true to a
limited degree of most liquids and some solids; that is to say, they
slowly disappear when left exposed in the air on account of
evaporation.

The reason that a gas tends thus to expand, and so exerts a
pressure upon the envelope designed to keep it in bounds, is that
its molecules are not held together by cohesion, but are indepen-
dent of each other, and free to move in accordance with their
inherent kinetic energy. Consequently they move in straight lines,
according to Newton's first law of motion, until they encounter
other molecules or the sides of the containing vessel, when the
direction of their path is changed.

Ideal gases are entirely devoid of cohesion between their mole-
cules, whilst these molecules are wide enough apart to behave
towards one another like points in space. In so far as gases
conform to these conditions, they behave similarly, and indepen-
dently of their chemical composition, under changes of temperature
and pressure. In this respect they differ from liquids and solids,
which show individuality of behaviour under such changes. Con-
sequently the two fundamental gas laws connecting the volume of
a gas with its temperature and pressure are of universal applica-
tion; and, although in no case rigidly true, are only seriously
departed from at high pressures and low temperatures, when the
molecules of a gas are being brought into closer relations with
one another.

i. The Gas L^ws

The Law of Boyle (or Marriotte),

The volume of a gas at constant temperature is inversely pro-
portional to its pressure.

p oc - or pv = constant,

or, since density and volume are reciprocal, ike density of a gas
varies directly as its pressure.

The Law of Charles (or Gay-Lussac).

The volume of a gas, at constant pressure, increases by

136 CHEMICAL THEORY

(= 0*00367) part of its value at C. for every degree rise of
temperature.

Thus 273 vol. of gas at become 274 vol. at 1.

283 10.
263 ,,-10,&c.,

and if the law held for all ranges of temperature, the volume of
a gas would vanish at 273. Although no such event takes
place, because conditions are modified at extremely low tempera-
tures, 273 is called the absolute zero of temperature, and
(t + 273) is absolute temperature, or T. Consequently the law
of Charles may be stated in another way the volume of a gas
varies directly as its absolute temperature; and to correct the
volume of a gas for change of temperature it is only necessary
to add 273 to each temperature concerned, and multiply or divide
by one or other absolute temperature, according to whether expan-
sion or contraction is taking place. Thus, for example, a volume
V at 10 becomes fffV at 15*. '

It may be added that, if the volume of a gas is kept constant,
while its temperature changes, the pressure of the gas varies directly
as its absolute temperature. This follows from Boyle's law.

The Gas Equation. v

Since pv is constant at constant temperature, and p or v varies
directly as the absolute temperature, while the other remains con-
stant, it follows that pv varies directly as the absolute tempera-
ture, or

pv = RT where R is a constant.

Thus, if, with a given quantity of gas, p, v, and T undergo
change, their values must be so related that R remains constant, or

This is a useful form of the gas equation which may be
employed in correcting the volume of gas for simultaneous changes
of temperature and pressure; for

EXAMPLE. 100 cu. cm. of air at 15 at 750 mm. pressure are

STATES OF MATTER AND PROPERTIES OF GASES 137

heated to 30, while the pressure becomes 770 mm. What is the
new volume?

= 100 X 760 X 303

1 770 X 288

The problem may, however, be solved quite easily without any
formula by considering whether the changes in temperature and
pressure, respectively, will cause increase or decrease of volume,
and then arranging the data accordingly.

2. Diffusion of Gases

Owing to their powers of indefinite expansion, gases mix when
free to do so. Thus, if a jar of ammonia, hydrogen sulphide, or
other strongly- smelling gas is opened in a room, the odour of the
gas will soon be perceived at some distance from the jar. This
process of mixing is called gaseous diffusion. It may be demon-
strated by placing a jar of hydrogen closed with a glass plate
above a similar jar of the brown gas nitrogen peroxide, and care-
fully withdrawing the glass plates which separate the gases. The
brown gas, although much heavier than hydrogen, will be seen
to be rising in the upper jar, and after a time the contents of
both jars will be uniformly brown. Or a jar of carbon dioxide
may be opened beneath one of hydrogen, and very soon the presence
of carbon dioxide in the upper jar may be proved by lime-water.

These effects, taking place in spite of gravitation, which tends
to keep the heavier gas in the lower cylinder, must be due to the
intensely active motion of the molecules of the diffusing gases.
Nevertheless the process takes a little time; it is not instantaneous,
as the process of free .expansion of a gas into a vacuum appears
to be. A mental picture of the process would show the molecules
of one gas rapidly threading their way through the obstructing
molecules of the other gas, with consequent and frequent deflections
and hindrances in their course, which are the cause of the time
consumed before the mixing process is completed.

The fact that mixed gases do not separate again on account
of their different densities was known to Priestley; but it was
Dalton, in 1803, who proved that a lighter gas cannot rest per-
manently upon a heavier, as oil upon water. Dobereiner, in 1823,
observed that hydrogen escaped through a crack in a glass flask
containing it, so that the gaseous pressure inside the flask

138

CHEMICAL THEORY

diminished; and Graham, in 1832, examined systematically the
phenomena of gaseous diffusion, and established the law regarding
them. Graham employed a glass tube closed at one end with a
plug of plaster of Paris. This material, as well as unglazed porce-
lain, or a thin plate of artificial graphite, is porous, i.e. it will
allow a gas to pass through it by diffusion whilst preventing its
escape in bulk. Consequently an alteration of gaseous pressure
may take place inside such a tube on account of diffusion; and
this when measured will indicate the extent to which diffusion
is taking place. Graham confined various gases in this tube
over water, and observed sometimes a rise, sometimes a depression
of the water, corresponding to a diffusion of the gas from the tube
into the air at a faster or slower rate than that at which the
air diffused into the tube. Then, by an analysis of the gas re-
maining in the tube, data were obtained by means of which the
following law of gaseous diffusion was established.
V The rates of diffusion of different gases are inversely proportional
to the square roots of their densitiesv '

For example, oxygen is 16 times as heavy as hydrogen; so
hydrogen diffuses ^/16 = 4 times as quickly as oxygen.

The law is true whether diffusion takes place into another
gas or into a vacuum. 1 \/

The following results of Graham's experiments substantiate
the law.

Gas.

Density :
Air = 1.

V density '

Velocity of Dif-
fusion : Air = 1.

Hydrogen

0-06949

3-7935

3-83

Methane

0-559

1-3375

1-344

Carbon monox

0-9678

1-0165

1-1149

Nitrogen

0-9713

1-0147

1-0143

Ethylene

0-978

1-0112

1-0191

Oxygen

1-1056

0-9510

0-9487

Hydrogen sulphide
Nitrous oxide

1-1912
1-527

0-9162
0-8092

0-95
0-82

Carbon dioxide

1-52901

0-8087

0-812

Sulphur dioxide

2-247

0-6671

0-68

*/ The statement of the law may be modified by saying that the
times of molecular passage of gases through a porous septum are
directly proportional to the square roots of the densities of the gases;

1 Probably the same law always applies to the rates at which different gases expand into
A vacuum under comparable conditions ; but the speed is ordinarily too great for measurement.

STATES OF MATTER AND PROPERTIES OF GASES 139

and this aspect of the law is illustrated by the following results of
Graham.^

Gas.

Time of Molecular
Passage into Air
at 100 mm. Pressure.

V Density :
Oxygen = 1.

Time of Mole-
cular Passage into
a Vacuum.

Hydrogen
Air
Oxygen
Carbon dioxide ...

0-2472

1-0000
1-1886

0-2509

1-0000
1-1760

0-2505
0-9501
1-0000
1-1860

The phenomena of diffusion may be illustrated in an interesting
manner by employing a large glass U-tube partly filled with
coloured water, and having a cylindrical porous pot attached to
one end of it by means of a tightly-fitting rubber bung.

When an inverted beaker of hydrogen is brought over the
porous pot (fig. 11), the water in the adjacent limb of the U-tube is
depressed, because hydrogen enters the pot quicker than air escapes
from it, and conse-
quently an excess of
pressure is set up
within the pot. When,
however, the beaker
is taken away, the hy-
drogen now within
the pot diffuses out
quicker than air can
enter, so that not
only does the water
rise again in the ad-
jacent limb, but it goes
beyond its original
level, showing a tem-
porary pressure within the pot less than that of the atmosphere.
Finally, the water slowly falls till it attains the same level in both
limbs.

If the apparatus is now modified so that the pot is attached
to the U-tube the other way up (fig. 12), a beaker of carbon dioxide
may be brought round it so that the pot is immersed in this heavy
gas. In this case the water will rise in the adjacent limb, because
carbon dioxide diffuses into the pot more slowly than air diffuses

Fig. 11

Fig. 12

140

CHEMICAL THEORY

out of it; and when the beaker is removed the pressure will
gradually be restored again, as the carbon dioxide escapes from
the pot and air takes its place.

The first part of the experiment may be modified so as to
cause a fountain to play (fig. 13), or the displaced liquid may be
caused to complete an electric circuit and so to ring a bell. The
presence in mines of methane, which is lighter than air, may be
indicated by an alarm given in this manner.

When gases pass through a minute aperture (^j- in. diameter)
in a thin metallic plate, they obey the law of diffusion. This

phenomenon was called by
Graham the effusion of
gases.

Diffusion or effusion may
be employed (a) to deter-
mine the relative densities
of gases, or (6) partially to
separate the constituents of
a mixture of gases. The
latter process is called at-
molysis.

(a) The density of ozone
relative to that of chlorine
was determined by Soret,
rig. is who showed that 227

volumes of chlorine dif-
fused in the same time as 271 volumes of ozone; or that the rates
of diffusion of the two gases were as 0-8376 : 1. If the density of
chlorine is 35-46, then that of ozone is consequently 24-9; for

1 : 0-8376 : : V35^46 : V24^9.

Similarly, Ladenburg found that a mixture of oxygen and
ozone containing 86-16 per cent of ozone required 430 seconds to
effuse, while pure oxygen required 367-4 seconds. Putting the
density of oxygen = 1, then

367-4 : 430 : : VI : Vl-3698;

and if the density of this mixture of 86-16 per cent ozone with
13-84 per cent oxygen is 1-3698, that of pure ozone may be
calculated to be 1-429 (0 = 1).

STATES OF MATTER AND PROPERTIES OF GASES 141

(6) Atmolysis takes place when electrolytic gas, i.e. a mixture
of 2 volumes of hydrogen with 1 of oxygen, passes through
an unglazed earthenware tube, such as the stem of a "church-
warden" tobacco-pipe; the gas which is collected does not explode,
but, owing to the escape of hydrogen by diffusion, contains a
sufficient proportion of oxygen to ignite a glowing wood splint.

Also, when ammonium chloride is vaporized in a glass tube
through the centre of which an unglazed pipe-stem passes, the
products of dissociation are partially separated by atmolysis; the
lighter ammonia, passing through the stem, may be blown out
against red litmus paper, which it turns blue, while the denser
hydrogen chloride, remaining in the glass tube, shows its presence
by reddening blue litmus paper.

Atmolysis has been employed partially to separate heavier
argon from lighter nitrogen derived from air, but the process is
imperfect and of little practical use.

From the standpoint of dynamics, diffusion may be attributed
to the velocity of the molecules of a gas, and the different rates
of diffusion to differences in molecular velocities.

Now the kinetic energy of a particle of mass m, moving with
a velocity v, is mi; 2 . Consider, therefore, two gases whose mole-
cular weights are m x and m 2 ; when these two gases ^are in thermal
equilibrium, that is at the same temperature, the following relation
will hold:

or
or

This relation, however, expresses the law of diffusion if v l and
v 2 stand for rates of diffusion. It may therefore be concluded
from the kinetic theory of gases, to which the above argument
belongs, that the rate of diffusion of a gas depends directly upon
the velocity of its molecules.

3. Pressure of Gaseous Mixtures

v Dalton's Law of Partial Pressures states that the total pressure
of a mixture of gases is equal to the sum of the partial pressures
of the individual gases.

Otherwise regarded, the law states that the different kinds
of molecules present in a gaseous mixture do not interfere with

142 CHEMICAL THEORY

each other, so that each gas exercises a pressure proportional to
the concentration of its molecules, as if that gas alone occupied
the whole space. A priori, there is no reason why this should
not be true; and, indeed, it must be true if Boyle's law is true
for all the gases in the mixture, and these gases have no chemical
influence on each other. For, to use a well-worn analogy, gases
occupy a space as the soldiers of an army might
occupy a country, not in massed battalions but in
widely separated units, so that the soldiers of an-
other army might equally occupy the country if
they were evenly distributed between the men of
the first army.

The law may be examined in this way. Suppose
a rectangular box divided by removable partitions
into several spaces, I, II, III, &c. ; let the volume of these spaces
be respectively v v v 2 , v 3 , &c., and let

= V.

Let different gases, A, B, C, &c., which have no chemical action
on one another, be contained in these spaces, and let them all
be at the same pressure, p. Now let the partitions be withdrawn,
so that the gases mix by diffusion and their molecules become
evenly distributed over the whole space. The partial pressure of

gas A will then be, according to Boyle's law, ^p, that of B -^p, and

so on; the law then states that the sum of these partial pressures
will equal the original pressure p:

i.e. -ip + p + p + tec. = p.
v v *v

Again, it may be seen that the validity of this law depends on
the validity of Boyle's law, for the above statement is reducible
to this: v l + v 2 + v 3 + &c. = v; and' so it may be said that
the volume occupied by a mixture of gases is equal to the sum of the
volumes occupied by its constituents under the same conditions of
temperature and pressure.

The truth of this statement may be tested in the following
way. From the densities of air, oxygen, and atmospheric nitrogen
the percentage composition of air by weight may be calculated;
it ma} also be determined from the volumetric composition, if
the density of air is left out of account, by multiplying the

STATES OF MATTER AND PROPERTIES OF GASfiS 143

relative volume of the gases by their respective densities, and
converting the resulting values into percentages, By calculation
air has thus been found to contain 23-21 per cent by weight of
oxygen and by estimation 23-18 to 23-23 per cent.

Dalton's law underlies the correction which is made for the
pressure of water vapour when a gas is measured over water,
and is therefore saturated with water vapour. The pressure of
water vapour at the temperature of the gas, which is subtracted
from the pressure under which the gas is measured, was estimated,
originally, in vacua \ and yet the correction is applied to a gas
at atmospheric pressure, under the justifiable assumption that
the presence of the air makes no difference to the pressure the
water vapour exerts. For example, if a quantity of gas is
measured over water, and the barometric pressure is 765 mm.,
whilst the room temperature is 14 C., the pressure of the dry
gas will be 765 12 = 753 mm., since the pressure of water
vapour at 14 C. is 12 mm.

4. Deviations of Gases from Boyle's Law

Boyle's law is true only for perfect gases, that is for gases
whose molecules are independent units, behaving like points
moving in space without attracting one another, and rebounding
after impact with unabated energy. The law is most nearly true
for gases at comparatively low pressures, and far from their
liquefaction temperatures. It is therefore truer at ordinary
temperatures for difficultly than for easily condensable gases. The

following table shows the values of ^ at atmospheric tem-
peratures, where p = 0-5 atmosphere and P = 1 atmosphere
for a series of gases arranged in the order of their condensability. 1

For a perfect gas ^ = 1.

Gas.

B.-P. Atmospheric
Pressure.

Temperature.

Hydrogen

(-253 C.)

10-7

0-99974

Nitrogen

(-196 C.)

14-9

1-00015

Air

(about -190C.)

11-4

1-00023

Carbon monoxide

(-190 C.)

13-8

1-00026

Oxygen

(-183 C.)

11-2

1-00038

Nitrous oxide

(-89-8 C.)

11-0

1-00327

Ammonia

(-33-5 C.)

9-7

1 -00632

1 Lord Rayleigh, Proc. Roy. Soc. t 1905, 74, 446.

144

CHEMICAL THEORY

Ideal gas

It will be observed that hydrogen at atmospheric pressure is
actually less compressible than accords with Boyle's law; it is
therefore sometimes spoken of asra more than perfect gas. The
deviations of the other gases from Boyle's law are very slight
because they are far removed from their liquefaction temperatures.
With a more condensable gas, like, carbon dioxide, the deviation
is much greater.

The accompanying diagram (fig. 14) expresses the deviations of
gases from Boyle's law graphically. Along the horizontal axis

pressures in atmospheres are
plotted, and along the vertical
axis values of pv. The state of
an ideal gas at constant tem-
perature is therefore represented
by a horizontal straight line.

It will be seen that, except
at very low pressures and up
to 330 atmospheres, hydrogen
at atmospheric temperature is
always less compressible than
an ideal gas, and nitrogen at
atmospheric temperature is
more compressible than an ideal
gas at moderate pressures, but
less compressible at high pres-
sures; and the same is true in
a much exaggerated degree of

carbon dioxide, which as an easily liquefiable gas departs widely
from Boyle's law. An interpretation of the shape of the carbon
dioxide curve for 35-1 may now be sought. The downward trend
of the curve at first shows that the value of pv rapidly diminishes
when that of p is increased. If the curve descended vertically,
liquefaction would be taking place, i.e. continuous diminution of
volume at constant temperature and pressure. Therefore the early
part of the curve shows an approach to the process of liquefaction.
The latter part of the curve, on the other hand, shows a difficulty
of compressibility approaching that of a liquid. This foreshadow-
ing of the condition and properties of a liquid by a substance still
a gas suggests that these two states of matter -are not always
sharply divided; and, as will be seen directly, there is a condition

Pressure in atxns.

330

Fig. 14

STATES OF MATTER AND PEOPERTIES OF GASES 145

in which the two states merge into one and become indistinguish-
able. Meanwhile it may be noted that another curve, representing
conditions of carbon dioxide at 0, contains a vertical portion ending
in an angle whence it diverges sharply upwards. It need hardly
be pointed out that true liquefaction is represented as taking place
in the part of the curve descending vertically, and that when this
is completed the upward turn of the curve represents the com-
pression of liquid carbon dioxide.

The physical properties of carbon dioxide may be expressed,
perhaps more simply, by a diagram in which pressures are plotted
directly against corresponding volumes. The isothermals, i.e. lines
of equal temperature, then take the form shown in the following
figure:

90.
85.
80.
75.
70
65
CO.
55

if'ig. 15

A is the isothermal of a perfect gas, which accords with Boyle's
law that pv = constant. Whatever point is chosen on A the
product of the numerical values of p and v corresponding to this
point is always the same.

With the isothermals of carbon dioxide for temperatures down
to 31 there is a diminution in the value of pv at moderate pressures

(DOO) tl

146 CHEMICAL THEORY

and an increase in its value at high pressures. This is shown
in the bulging of the isothermals, the curves being flattened in the
middle, where the value of pv diminishes owing to the increased
compressibility of the gas, and steepened towards their upper ends,
where on account of greatly diminished compressibility the value
of pv rapidly increases. If the curve in any part became horizontal,
that would represent liquefaction. Such is the case with the
isothermal for 21-5, which consists of three distinct parts: firstly,
an upward sloping part, BC, representing compression of carbon
dioxide gas; secondly, a horizontal part, CD, representing diminu-
tion of volume at constant pressure which is the condition of
liquefaction; thirdly, a steeply sloping, almost vertical, part, DE,
representing the slight compressibility of liquid carbon dioxide.
Liquefaction of carbon dioxide is possible only at temperatures
whose isothermals consist in part of horizontal straight lines. The
dotted line xyz marks the upward boundary of liquefaction. At
temperatures whose isothermals pass at all within the region thus
outlined, liquefaction will take place if sufficient pressure is applied;
at other temperatures liquefaction is impossible at any pressure.

Attention may now be drawn to the isothermal which touches
the top of the curve xyz. Its temperature is 31 or 304 absolute.
This is the highest temperature at which liquefaction can take
place, or the lowest at which it can be avoided. It is evidently
a critical temperature, and the carbon dioxide is in a critical state.
The temperature is called the critical temperature for carbon
dioxide, and the state the critical state. But what is to be seen
there? The experiment of Dr. Andrews answers this question.

Carbon dioxide gas at 21-5 was compressed in a strong glass
tube until it was seen to liquefy at a pressure of 60 atmospheres.
Thus two layers were observed in the tube, the dense gas above
and the light liquid beneath. When the temperature was raised,
the liquid expanded and at the same time partially evaporated,
so that, since the total volume was kept constant, the gas became
denser because of what it received from the liquid, while the liquid
became specifically lighter because of expansion. Thus the densi-
ties of the two phases, liquid and gas, approached each other. At
31, under a pressure of 73 atmospheres, these densities became
identical, the line of demarcation between liquid and gas dis-
appeared, and the tube became filled with a homogeneous fluid
hardly describable as liquid or gas.

STATES OF MATTER AND PEOPERTIES OF GASES 147

When the temperature was slightly lowered, a flickering was
seen like the air currents faintly visible above a heated object;
and then the liquid appeared again, only to disappear in the same
way if the temperature was slightly raised. These changes were
brought about, it is to be observed, by slight variations of
temperature under constant pressure; they could not be induced
by slight variations of pressure at constant temperature; i.e. if the
temperature was maintained slightly above 31, so that no liquid
appeared in the tube, it was impossible by increasing the pressure
at this temperature to liquefy the carbon dioxide.

The fluid state, above described, which is common ground, so to
speak, between liquid and gas, and is reached from either side by
an approximation of density, is the critical state of carbon dioxide;
the temperature 31 C. is the critical temperature of this substance;
73 atmospheres is its critical pressure] the density of the fluid
at the critical temperature and pressure, which is 0*464, is its

critical density, and the specific volume (-3 rpj its critical

volume.

These terms may now be defined as follows:

The critical temperature of a gas is the highest temperature
at which the gas can be liquefied by compression.

The critical pressure of a gas is the pressure under which it is
liquefied at its critical temperature.

The critical state of a substance is the state in which its gaseous
and liquid phases merge into one homogeneous fluid.

The critical density is the density of the fluid in its critical
state; and

The critical volume is the specific volume of the fluid in its
critical state.

One possible misunderstanding must be avoided. It is not
necessary to realize the critical pressure in order to liquefy a gas;
the critical pressure is not the lowest pressure under which a gas
can be liquefied. A gas may be liquefied under a pressure of
one atmosphere or less if the temperature is low enough.

5. The Liquefaction of Gases

The student has already learned that the two conditions neces-
sary for the liquefaction of a gas are the lowering of its tempera-
ture and the increase of its pressure. He is also acquainted with

148

CHEMICAL THEORY

the fact, unknown to the earlier experimenters in this region, that
no amount of pressure can compensate for insufficient lowering of
temperature; that unless a gas is cooled below its critical tempera-
ture it cannot be liquefied.

The history of the liquefaction of gases is the history of the
gradual elucidation of fundamental physical principles, and of their
ingenious experimental application; and the subject well repays
careful study.

It will be well first to set forth a list of the commoner gases
with their critical temperatures and pressures, and boiling-points
under atmospheric pressure.

Gas.

Sulphur dioxide
Chlorine

Critical
Temperature.

157-2
141

Critical Pressure.
Atmospheres.

77-7
83-9

B.-P. under 760 mm.
Pressure.

-10-1
:i3-7

Ammonia*

132-9

112-3

-i3-46

Hydrogen sulphide

100-4

89-1

<;i -8

Hydrogen chloride

51-4

. .. . 81-6

33-0

Nitrous oxide

36 5

71-7

89-8

Acetylene

35-5

61-5

sublimes at 83-6

Carbon dioxide

72-9

Ethylene

103-5

Methane
Nitric oxide
Oxygen
Carbon monoxide
Nitrogen
Hydrogen
Helium

82-8
-93-5
-118-8
139-5
-145
-241
-267-75

45-6
71-2
50-8
35-5
33-6
19-4
2-75

-164
-149-9
-182-9
-190
-195-5
-252 -5
-?68-7

It will be seen that the gases in the above table, as far as
ethylene, might be liquefied by the application o sufficient pressure
with little or no cooling; but that the remaining gases could not
be liquefied by the earlier experimenters, because they had not
the means of securing sufficiently low temperatures, even if they
had recognized the necessity for them. Consequently, these gases
were called " permanent gases ".

The following are the methods which have been employed for
the liquefaction of gases:

i. The Method of Simple Compression, with or without Cooling.

Monge and Clouet liquefied sulphur dioxide before the year
1800 " by extreme artificial cold and a strong pressure exerted at

STATES OF MATTER AND PROPERTIES OF GASES 149

the same time"; and in 1805-6 Northmore liquefied sulphur di-
oxide, chlorine, and hydrogen chloride. The most important and
interesting of the earlier experiments, however, were those carried
out by Faraday under the direction of Sir Humphry Davy*

Faraday sealed up solid chlorine hydrate, C1 2 8H 2 O, in a glass
tube, and immersed the tube in boiling water. The hydrate melted,
and produced two liquids, one of which was chlorine water, and
the other liquid chlorine. Faraday generally employed, however,
a sealed glass tube bent at right angles. At one end of the tube
were placed the ingredients from which the gas could be generated
by heat; whilst the other end, which was empty, was immersed
in a freezing-mixture. The pressure produced by the confined gas
then sufficed to liquefy it in the cooled end.
In this way Faraday liquefied all the com-
mon gases, except methane and those that
follow it in the above list; these latter
he thus considered to be "permanent".

Faraday's results were communicated
to the Royal Society in a series of papers
during the years 1823-45; and in 1834 ' ' Fi 16

Thilorier liquefied carbon dioxide on a
large scale by decomposing sodium carbonate with sulphuric acid
in a strong copper vessel connected with another copper vessel in
which the gas was condensed. He also obtained carbon dioxide
snow by the rapid evaporation of the liquid.

The results of Andrews's experiments on carbon dioxide were
first published in 1863, and formed the subject of his Bakerian
lecture on " The Continuity of the Gaseous and Liquid States " in
1869. Thenceforth the necessity of cooling a gas below its critical
temperature in order to liquefy it was recognized; and this prin-
ciple was kept in view by all subsequent experimenters.

ii. The Cascade Method of Cooling, u

Sulphur dioxide can be liquefied at atmospheric temperature by
moderate compression; and it may then be evaporated again at
a lower temperature under reduced pressure. The liquid sulphur
dioxide, thus cooled, may serve to cool the carbon dioxide, so that
the latter gas can be liquefied in turn by suitable compression.

By rarefaction the liquid carbon dioxide may be made to
evaporate quickly, so that its temperature falls while it solidifies;

150

CHEMICAL THEORY

and so a " permanent gas ", cooled below its critical temperature by
means of this solid carbon dioxide, may be liquefied by compression.
This is the principle of the cascade method, by which Pictet lique-
fied oxygen in 1877.

Liquid sulphur dioxide was made to boil under reduced pressure
at 65 in a cylinder surrounding an inner tube containing carbon

Fig. 17

dioxide, which was liquefied at this temperature under a pressure of
4 to 6 atmospheres (fig. 17). The liquid carbon dioxide then flowed
into another cylinder, which surrounded oxygen highly compressed
in a steel tube connected with a steel retort in which it was being
generated. By evaporation under low pressure the carbon dioxide
was here cooled to a temperature of 140, which is well below
the critical temperature of oxygen.

The pressure under which the oxygen was contained in the
steel tube was 320 atmospheres; which was much in excess of
what was necessary. Since, therefore, oxygen gas had been cooled
below its critical temperature, under adequate pressure, it was
liquefied. V

STATES OF MATTER AND PROPERTIES OF GASES 151

ill. The Method of Adiabatic l Expansion.

If a force / moves through a space 8, work is performed which
is numerically equal to the product fa] and if a gas expands by
a volume v against a pressure p, it performs work equivalent to
pv. If the expansion is adiabatic, the only source of energy is the
heat of the gas; consequently a gas is cooled
by adiabatic expansion on account of the per-
formance of external work.

Cailletet applied this principle to the cool-
ing of gases in 1877, and succeeded in lique-
fying methane, nitric oxide, oxygen, carbon
monoxide, air, and nitrogen. The apparatus
employed (fig. 18) consisted of a strong capil-
lary glass tube (T) to contain the gas over
mercury. This tube was immersed in a closed
wrought -iron mercury bath, into the upper
part of which water (W) could be forced by
hydraulic pressure applied by a screw. Thus
the gas in the capillary tube could be highly
compressed, and then by the sudden releasing
of the pressure made to expand rapidly against
the pressure of the mercury. The work per-
formed in this expansion deprived the gas of
sufficient energy to cool it below its critical
temperature, whilst sufficient pressure re-
mained to liquefy it. A mist seen inside the tube was evidence
of liquefaction.

The principles underlying the methods of Pictet and Cailletet
have^been subsequently employed by Wroblewski, Olszewski, and
Dewar, but a different principle forms the basis of more recent
work upon the liquefaction of gases.

iv. The Method of Self-intensive Refrigeration.

When a perfect gas expands into a vacuum it does no kind
of work, and therefore does not change its temperature. But most
gases are less than perfect; they are more compressible at moderate
pressures than corresponds with Boyle's law, because their mole-
cules are not quite independent of one another. Intermolecular

i " Adiabatic " means allowing nothing to pass through. The term signifies the condition
of imperviousness to heat.

Fig. 18

152 CHEMICAL THEORY

attractions thus constitute a force to overcome which energy must
be supplied. Consequently, when a compressed gas expands into
a vacuum it is cooled slightly by reason of the internal work it
performs in overcoming the attractive force between its
own molecules. Thus a lowering of temperature is observed
when a gas expands through a small orifice; this effect,
which is called the Joule-Thomson effect, is, however, quite
small, amounting, for air at C., to only 0*29 C. for a
fall in pressure of 1 atmosphere. The effect has, however,
been made cumulative, and of use for gaseous liquefaction
by the following ingenious device.

Suppose a compressed gas within the tube A (fig. 19)
escapes through the small orifice B, and passes back at

/Tl

reduced pressure up the outer tube C, in which the inner

tube is enclosed. Whilst the gas is expanding, the orifice
is cooled a little, and in consequence serves to cool the
L gas as it passes through. This cooled gas, returning
Fig. 19 by C, then cools the gas passing down A. Thus the
temperature of the gas delivered through the central tube
continuously falls until at length the critical temperature is passed,
and liquid drops from the orifice through which the compressed
gas has hitherto been driven. The temperature at which this
takes place depends upon the internal pressure upon the gas
whilst the rate of the liquefaction process depends
chiefly upon the efficiency of the cooling. The mean-
ing of the term " self -intensive refrigeration" should
now be plain.

This principle was successfully applied to the
liquefaction of air by Linde and Hampson in^,1895.
The apparatus consists essentially of a long copper
spiral tube, surrounded concentrically by a larger tube
and packed in insulating material. The gas, either
compressed specially "to 150 atmospheres or more, or
Fig. 20 delivered from a gas cylinder directly into the appa-
ratus, passes down the inner tube, through the valve
at its base, and back at a low uniform pressure through the an-
nular space between the two tubes. The liquid is collected in a
"Dewar vessel" (fig. 20), which is a double- walled glass flask with
the space between the walls completely evacuated. In consequence,
convection of heat from the air to the contents of the vessel does

STATES OF MATTER AND PROPERTIES OF GASES 153

not take place, and the heat which penetrates by radiation is only
about one-sixth what would otherwise be conveyed. If the glass
is silvered, the heat entering is reduced to about one-thirtieth part
of what would enter through a single-walled glass vessel. The
well-known thermos-flask is constructed on the same principle as
the Dewar vessel.

Liquid air, which is now commonly used for low-temperature
research, was prepared in large quantities by Dewar, who employed
liquid ethylene, boiling in vacua at 152, as a cooling agent.

Air is now generally liquefied, however, by the self-intensive
process of Linde or Hampson, which depends on the performance
of internal work, and takes advantage of the Joule-Thomson effect,
or by the Claude process, in which cooling is secured by the per-
formance of both internal and external work.

In the Linde or Hampson process air carefully freed from
carbon dioxide and water vapour expands from a pressure of 150-
200 atmospheres within the coil of the apparatus to a little more
than atmospheric pressure outside the nozzle. About 5 per cent
of the air that passes through the apparatus is liquefied, and about
H litres of liquid air are obtained per hour. Since oxygen is more
condensable than nitrogen, liquid air contains a larger proportion
of the former element than gaseous air.

The cooling effect of expansion of a compressed gas with per-
formance of external work is much greater than that of free
expansion; -and the Claude process combines these two effects, in
which air is compressed to 40-50 atmospheres, and then divided
into two parts. One of these parts is allowed to expand in a
cylinder, so as to perform external work, and the air so cooled
then surrounds the other compressed part of the air, which is thus
sufficiently cooled to liquefy at once by free expansion. Con-
siderable economy of power is effected by this device.

The Liquefaction of Hydrogen and Helium.

Pictet imagined that he had liquefied hydrogen, and described
this substance as a steel-blue liquid. This, however, was a mistake.
Wroblewski, in 1884, first converted hydrogen into- a "liquide
dynamique " having the appearance of an instantaneous froth, by
cooling the gas to the boiling-point of the oxygen, and then
suddenly expanding it, according to Cailletet's method, from
100 atmospheres to 1 atmosphere pressure; and, Olszewski, in 1891,

154

CHEMICAL THEORY

performed similar experiments, but failed to obtain hydrogen as a
" static liquid" showing a meniscus, because no liquid was available
having a boiling-point lower than that of oxygen, by meana of
which the hydrogen gas might be cooled. Dewar in 1898 overcame
this difficulty by employing the self -intensive process, to cool the
gas. It will be remembered, however, that hydrogen is a "more
than perfect" gas at atmospheric temperatures; it is less compres-
sible than accords with Boyle's law, and is actually warmed by free
expansion. Nevertheless, below 805 C., which is called its
" inversion temperature ", hydrogen becomes, like other gases, loss
compressible than Boyle's law requires, and like them is cooled by
free expansion. Consequently, when cooled by liquid air below
its inversion temperature, hydrogen was liquefied in quantity by
self-intensive refrigeration.

Helium is even more difficult to liquefy than hydrogen, but,
in 1908 Kamerlingh Onnes liquefied this gas by the self -intensive
process, after first cooling it to the temperature of liquid hydrogen.

Practical Applications of Liquefied Gases.

In addition to their use for low-temperature research, liquefied
gases find practical application in several ways. Liquid ammonia

is employed in the manufacture of
ice. In the oldest freezing -ma-
chine, that of Carre*, concentrated
solution of ammonia is employed
in the following way.

The iron vessel A -contains
water saturated with ammonia at
C. When this vessel is heated,
ammonia gas distils by the pipe C
and condenses between the walls
of the cup-shaped vessel B. Con-
densation takes place at atmos-
pheric temperature when the pres-
sure of the gas reaches about 7

atmospheres. After most of the ammonia has thus been liquefied
in B, A is immersed in cold water, and gaseous ammonia at once
begins to redissolve in the water contained in A so as to re-form
the original solution. Consequently liquid ammonia evaporates
from B, and the heat necessary for this evaporation is drawn

Fig. 21

STATES OF MATTER AND PROPERTIES OP GASES 155

from the surroundings; so water, placed within the cup, may be
frozen.

In more modern ice -machines use is made of anhydrous
ammonia, which is liquefied by pressure, and then evaporated again,
the latent heat of evaporation as before being drawn from water,
which is thereby frozen.

When a kilogram of liquid ammonia evaporates at 10 C.,
322*3 calories 1 are absorbed. It must be remembered, however,
that this same amount of heat is evolved when an equal quantity
of ammonia gas is liquefied by compression; and provision must be
made for the removal of this heat. In practice water is frozen,
not by the direct contact of vessels containing it with evaporating
ammonia, but by their immersion in brine cooled to 10 by the
ammonia which is evaporating in coils that pass through it.

The parts of the machine are represented diagrammatically as
follows:

Expansion

coil in
Brine Tank

Condensing
coil

Expansion Valve
Fiff. 22

Gaseous ammonia is compressed by the pump A, and liquefied
in the condensing-coil B, the heat generated during condensation
being removed by water. The flow of the liquid ammonia is
regulated by the expansion -valve C, and this liquid evaporates
in the expansion -coil immersed in the brine -tank D, which also
contains rectangular cans in which is the water to be frozen. The
evaporated ammonia returns to the pump, again to go through
the same cycle of condensation and evaporation.

Chlorine, carbon dioxide, and sulphur dioxide gases are liquefied

1 A calorie is the amount of heat necessary to raise 1 grm. of water through 1 0.

156 CHEMICAL THEORY

for convenience of storage and portability. Liquid chlorine is
employed in large quantities for chlorination in the manufacture
of organic chemicals; and its use in recent warfare is notorious.
Carbon dioxide is stored in cylinders in the liquid state; for, unlike
oxygen and hydrogen, this gas is liquefied at atmospheric tem-
perature by moderate pressure; consequently cylinders containing
as much as 28 Ib. of the liquid can be supplied. When such a
cylinder is tilted, and the valve opened so that liquid carbon
dioxide can escape into the air, some of it is solidified and appears
in the form of carbon dioxide snow, on account of the heat removed
by the portion which evaporates. Liquid carbon dioxide is also
used in the manufacture of aerated waters, as a refrigerant in place
of ammonia, and for other purposes. Liquid sulphur dioxide is
supplied in steel cylinders or in thick glass siphons, convenient
for laboratory use. It is employed in bleaching, for antiseptic
purposes, and as a refrigerant.

The liquefaction of air, and the fractional distillation or
rectification of the liquid, are now employed as a means of pre-
paring oxygen and nitrogen gases from the air.

If liquid air is caused to evaporate as fast as it is produced,
the more volatile gas which escapes from it is nitrogen containing
about 7 per cent of oxygen; while the liquid that remains, and
can be drawn off and evaporated, is practically pure oxygen.

By a process of rectification, also, that is by allowing a
sufficiently cooled air-current to pass upwards through a down-
flowing stream of liquid air, nitrogen and oxygen may be
separated. This is because the more condensable oxygen in the
up-flowing air-current is liquefied, while the heat thus liberated
serves to gasify some of the nitrogen of the down-flowing liquid
air. By liquefying the escaping nitrogen, and repeating the
process, gaseous nitrogen may be obtained practically free from
oxygen.

The uses of oxygen are well-known. Nitrogen, obtained from
the air through liquefaction, is employed in large quantities for
the manufacture of calcium cyanamide and synthetic ammonia.

Hydrogen is now obtained from water-gas which consists of
this gas and carbon monoxide in approximately equal proportions
by fractional condensation of the carbon monoxide, which leaves
the hydrogen nearly pure.

STATES OF MATTER AND PROPERTIES OF GASES 157

SUMMARY

THE GAS LAWS: Law of Boyle. The volume of a gas at
constant temperature is inversely proportional to its pressure.

p oc -, or pv = constant.

r v

Law of Charles. The volume of a gas at constant pressure

increases by -L ( = 0-00367) part of its value at C. for
2 7 3

every degree rise of temperature; or

The volume of a gas varies directly as its absolute temperature.

The Gas Equation: PV = RT. (T = absolute temperature,
V = the gramme-molecular volume.)

Law of Gaseous Diffusion. The rates of diffusion of different
gases are inversely proportional to the square roots of their
densities.

Daltoris Law of Partial Pressures. The total pressure of a
mixture of gases is equal to the sum of the partial pressures of
the individual gases.

DEFINITIONS OF THE CRITICAL STATE. The critical temperature
of a gas is the highest temperature at which a gas can be liquefied
by compression.

The critical pressure of a gas is the pressure under which it is
liquefied at its critical temperature.

The critical state of a substance is the state in which its gaseous
and liquid phases merge into one homogeneous fluid.

The critical density is the density of the fluid in its critical
state.

The critical volume is the specific volume of the fluid in its
critical state.

METHODS OF LIQUEFACTION OF GASES:

i. The method of simple compression, with or without cooling
ii. The cascade method of cooling.

iii. The method of adiabatic expansion.

iv. The method of self -intensive refrigeration.

CHAPTER VIII
THE PROPERTIES OF LIQUIDS

It has already been seen how liquids differ from gases in the
relations of their molecules. An outcome of this difference is
that liquids possess more highly individualized properties than
gases. Perfect gases behave alike towards changes in pressure
and temperature; liquids are unequally though very slightly
compressible, and have separate co-efficients of thermal expansion.

This is because whilst the molecules of perfect gases are so far
removed from one another that they behave as points in space,
and their chemical composition has no influence on their physical
properties, the molecules of liquids approximate to each other,
with the result that their chemical composition influences their
physical behaviour.

The relation between the liquid and gaseous states has been
studied in the former chapter*, and it has been seen that liquids
and gases appear to be sharply distinguished except in the critical
state. There is no such sharp distinction between liquids and
solids, and there is no question of critical state. Liquids are
distinguished as mobile and viscous, or, in simpler words, as thin
and thick, and a liquid may be so viscous as to be scarcely
distinguishable from a solid. Thus, while water is a mobile liquid,
and ether is still more mobile, glycerine is comparatively viscous,
and pitch and beeswax are so extremely viscous as scarcely to
flow at all, and therefore frequently to be regarded as solids.
Experience shows that a viscous liquid increases in viscosity as
it is cooled, and becomes more mobile when heated. Pitch, for
example, which at low temperatures is brittle and may be broken
with a hammer, softens gradually . and becomes definitely liquid
when sufficiently heated; but it does not possess a melting-point,
that is to say, there is no fixed temperature at wl^ich change
from solid to liquid takes place.

The viscosity of a liquid is an important physical property,

158

THE PROPEBTIES OF LIQUIDS 159

which is measured by observing the rate of flow of the liquid
through a capillary tube under standard conditions.

Of the various other physical properties of liquids those which
will be considered here are: density, and its reciprocal value
specific volume; and vapour pressure, together with boiling-point

Density.

Different liquids differ much in density. Water, being the com-
monest liquid, provides the standard of 'density; but since changes
of temperature affect the densities of liquids on account of thermal
expansion, a definite temperature must be chosen at which the density
of water is reckoned to be unity; and similarly a definite tempera-
ture at which the liquid under consideration is to be examined.

Water assumes its maximum density at 4 C., from which
temperature it expands and becomes specifically* lighter whether
warmed or cooled. Consequently the density of water at 4 9 C.
is usually regarded as the standard, and said to be unity. Th'us
the densities of other liquids may be compared with that of water
at 4 C. For example, the density of chloroform at 11-8 C.

1 1 ft
has been found to be d * = 1-5039, whilst at 15, compared

1^
with water at 15, d~ = 1-5009. ^

The reciprocal of density or specific gravity, 1 -= - TT-, is

specific volume. Thus the specific volume of chloroform is
.. - O OQ = 0-665; and the product of specific volume and molec-

ular weight is molecular volume, so that the molecular volume
of chloroform is 0-665 x 119-4 = 79-4.

The relationships between molecular volumes of related com-
pounds have been found to be a subject worthy of study and to
yield results of interest. Thus the molecular volumes of successive
members of series of analogous carbon compounds, i.e. of homol-
ogous series of such compounds, are found to stand in arithmetical
progression. For example, the fatty acids:

Molecular Volumes. Differences.
Formicacid ...... H-COOH ...... 41-4

Acetic acid ...... CH 3 .COOH ...... 63-7 **'*

Propionic acid ...... C 2 H 5 -COOH ...... 85-4 * l 'l

Butyric acid ..... C 3 H r -COOH ...... 107-1 21 ' 7

1 The terms density and specific gravity are generally used as synonymous by chemists.

160 CHEMICAL THEOEY

Thus, for an atomic difference of CH 2 , there is a difference of
molecular volume of approximately 22 units; or, for every addition
of CH 2 to the molecule there is an addition of 22 to the molecular
volume. On this account molecular volume is said to be an
additive property. It is even possible to go further and attribute
a volume-effect to each atom within the molecule. Thus, from the
work of Kopp, the following values in the case of acetic acid,
CH 3 -COOH, are obtained:

2C = 11 X 2 =22
4H = 5-5 X 4 = 22
0"(inCO) 12-2

0' (in OH) 7-8

64-0 calculated.
63-7 observed.

In the following table are given the densities of a number of
liquids:

Liquid. Density.

Liquid hydrogen 0-0763 (-259-9)

Liquid helium 0-154 (cir. 269)

Liquid methane, CH 4 - 415 (- 164, B. P.)

Pentane, C 6 H 12 0-627(14)

Ethyl ether, (C 2 H 5 ) 2 O 0-72(17.4)

Ethyl alcohol, C 2 H 6 OH 806 (0)

Paraffin oil 0-8 to 0-85

Benzene, C G H 6 0-874

Carbon disulphide, CS 2 1-292 (0)

Carbon tetrachloride, CC1 4 1 595 (0)

Ethyl iodide, C 2 H 6 I 1-944

Methyl iodide, CH 3 I 2-293

Bromof orm, CHBr 3 2 9 (15)

Bromine 3-188(0)

Methylene iodide, CH 2 I 2 3-293

Mercury 13-595(0)

From these figures it is seen that the range of density is very
great; thus, mercury is about 178 times as heavy as liquid hydro-
gen. Heavy liquids, such as bromoform and methylene iodide,
are used in mineralogy to separate mineral fragments of different
densities, the lighter of which will float and the heavier sink in
a liquid of intermediate density.

Mercury, the only liquid metal, has a density far exceeding
that of any other liquid. Were it not for this, the manufacture
of accurate barometers of reasonable dimensions would not be

THE PROPERTIES OF LIQUIDS 161

Vapour Pressure and Boiling-point.

Liquids differ much in volatility or rate of evaporation at
atmospheric temperature and pressure. Thus a little ether, poured
upon a flat surface, will soon evaporate and disappear, whilst a
similar quantity of water will be little diminished. On the other
hand, sulphuric acid and glycerine when left exposed to the air
actually increase in bulk, because they do not evaporate but are
hygroscopic, and absorb water vapour from the air.

The idea of vapour pressure will be made plain by the
following consideration. Suppose a liquid is contained in a closed
evacuated vessel, which it does not fill. At a given temperature
a certain quantity of the liquid will evaporate into the space
above it until the pressure of the accumulating vapour reaches
a certain value, when a state of dynamic equilibrium will be
established in which evaporation and condensation take place
at equal rates. The pressure thus established is called the vapour
pressure of the liquid at the given temperature.

The vapour pressures of water at different temperatures are
shown in the following table:

Temperature C. Pressure mm. Hg. Temperature 0. Pressure mm. Hg.

4-6 80 354-3

5 6-5 90 525-5

10 9-2 100 760-0

15 12-7 111-7 1140-0

20 17-4 120-6 1520-0

30 31-5 127-8 1900-0

40 54-9 133-9 2280-0

50 92-0 144-0 3040-0

60 148-8 180-3 7600-0

70 233-1 213-0 - 15200-0

The relationships between these temperatures and the correspond-
ing pressures are shown graphically in the diagram (fig. 23) on
the following page, in which pressures are recorded on the vertical
and temperatures on the horizontal axis.

The vapour pressure curves for other liquids take a similar
form.

It will be observed that the vapour pressure at 100 C. is
760 mm. Now, 100 C. is said to be the boiling-point of water,
and 760 mm. is normal atmospheric pressure. A liquid boils
when its vapour pressure has become equal to atmospheric

(D60) 13

162

CHEMICAL THEORY

pressure; so that the various temperatures in the table on p. 161
represent the boiling-points of water under the corresponding
pressures. Thus on a mountain, where the barometer registers
525 5 mm., which is at a height of about 8500 ft, water will
boil at 90 C.; and on a higher mountain, where the pressure is

1700
1600
1500
1400
1300
1200

lOOO
&

& 900

.9 800

I 70
600
500
400
300
200

10 20" 30 40 50 60 70 80" 90" 100*110 120" 130*
Temperature.

Fig. 23

only 354-3 mm., at an altitude of 14,600 ft., it will boil at 80 C.
When a gas is measured over water, as is frequently done in the
estimation of gases, it is saturated with water vapour, and its
pressure is that of the dry gas plus the pressure of water vapour
at the observed temperature. The pressure of the dry gas is
therefore found by subtracting the pressure of water vapour at
the temperature of the experiment from the barometric pressure,

THE PROPERTIES OF LIQUIDS

163

provided the level of the water is the same inside and outside
the measuring-tube.

Not all liquids can be distilled at atmospheric pressure, for
some are decomposed by heat before they can reach a boiling-
point under these conditions. Pure hydrogen peroxide, for
example, decomposes when heated above 70 C., but does not
boil at this temperature under atmospheric pressure. It has
already been seen, however, that the boiling-point of a liquid is
lowered by lowering
the superincumbent
pressure; and the
application of this
principle permits of
the distillation and
purification of various
liquids which could
not otherwise be puri-
fied.

Distillation under
reduced pressure can
be carried out on a
small scale by means
of two distilling-
flasks fitted together
as shown in the figure.
The liquid to be dis-
tilled is contained in
the flask A, which
is closed by a tightly-fitting cork, through which a thermometer
passes. The delivery-tube of this flask passes through a cork
which fits tightly into the neck of a second flask B; and to the
delivery-tube of this flask is attached a rubber tube which leads
to an air-pump capable of reducing the pressure inside the flasks
to about 15 mm. of mercury. The liquid in the distilling-flask is
liable to bump when boiled; this is due to overheating whilst
boiling is delayed, followed by a sudden ebullition, which is liable
to carry the liquid bodily over into the collecting-flask. To prevent
this, and steady the boiling process, a few fragments of pipeclay
or unglazed earthenware are placed in the liquid. Air escapes from
the pores of the earthenware before the liquid boils, and vapour

To Pump

Fig. 24

164

CHEMICAL THEORY

of the liquid is generated at the surface of the earthenware, which
thus helps to steady the boiling by preventing overheating. In
this, or a similar apparatus, pure hydrogen peroxide can be dis-
tilled at 69*2 C. under a pressure of 26 mm.

Relation between Boiling-points of Liquids in Homologous Scries.

The boiling-point of a substance depends upon its molecular
weight and chemical constitution. This is well seen in the various
homologous series of carbon compounds.

Saturated
Formula.

C r H 16

dnUoo

normal Hydrocarbons,

B.-P. C. Diff.
37

Norm
Formula.

CH 3 .QH
C 2 H 5 .OH
C 3 H 7 .QH
C 4 H fi -OH
C 6 H n .OH
CJI,.. OH

al Monohydric
Alcohols
B.-P. 0.
66

Diff.

22
19
20
21
19

27
25
23

125
150

117

138

173

157

It is seen that in the case of the hydrocarbons the differences
between the boiling-points of successive members of the series
diminish as the series is ascended, whilst with the normal mono-
hydric alcohols these differences remain approximately constant.

Further, it is interesting to notice that constitution as well as
molecular weight affects boiling-point, for the more compact the
molecules of a liquid are the lower is the liquid's boiling-point.
Thus there are three pentanes which have the following constitu-
tions and boiling-points:

Normal pentane

CH 3 -CH 2 -CH 2 -CH 2 -CH 3
B.-R 37 C.

Isopentane

/CH,
CHg-CHj-CH

\CH,
B.-P. 30 C.

Tetrarn ethylmethane.

CH 3
CH 3 .C-CH 3

CH 3
B.-P. 9-5C.

So it seems that molecules consisting of simple elongated chains
are more difficult to disengage from the state of liquid than mole-
cules of equal weight which are* branched, and consequently not
so elongated.

THE PROPERTIES OF LIQUIDS 165

SUMMARY

The specific volume of a liquid is the reciprocal of its density.

Molecular volume is the product of specific volume and mole-
cular weight.

The molecular volume of a liquid is an additive property; it
is the sum of the atomic volumes of the constituent atoms of its
molecule.

A vapour pressure curve represents the relation between the
vapour pressures of a liquid at different temperatures, or otherwise
the temperatures at which the liquid will boil under different
pressures.

The boiling-points of successive members of a homologous series
of chemical compounds rise regularly with successive increments of
molecular weight

CHAPTER IX
THE PROPERTIES OF SOLIDS

It has already been seen that solids differ from liquids by
possessing a shape of their own, and from gases by possessing,
in -addition, a definite volume at a given temperature, which is
little affected by changes of pressure; and, further, that these
properties depend upon the fact that the molecules of solids possess
only vibratory, and not to any extent translatory, motion.

The following topics will be considered in this chapter: The
manner of formation of solids from a state of (a) vapour, (6) fusion,
(c) solution; and then the following properties: physical form,
including the crystalline and amorphous states; isomorphism, poly-
morphism, and allourtfpy.

The Formation of Solids.

(a) SOLIDIFICATION OF VAPOURS

Most vapours, when sufficiently cooled under atmospheric pres-
sure, liquefy, but there are some which solidify without passing
through the liquid state. Examples of these are ammonium
chloride and arsenious oxide. There are other substances, such
as iodine, sulphur, and water, which may either liquefy or solidify
from the state of vapour, according to circumstances. Atmospheric
.water vapour, for example, becomes rain or snow, according to
temperature; and sulphur, when distilled, will yield flowers of
sulphur or liquid sulphur according to the rate of cooling.

When solid iodine is kept in a bottle at atmospheric tempera-
ture it slowly vaporizes and produces crystals on the sides of
the bottle. If, however, it is gently heated in a test-tube, it is
observed to melt and to produce a violet vapour simultaneously;
but when this vapour becomes cooled in the upper parts of the

100

THE PROPERTIES OF SOLIDS

167

tube it yields crystals of iodine without passing through the liquid
state.

In the former case the iodine behaves thus:

Solfd

liquid

vapour

in the latter case thus:

Solid

liquid

vapour

The former is a true case of sublimation, in which no liquid
appears at all. In the latter case, although a sublimate of iodine
appears, the melting of the iodine during heating precludes the
description of the whole process as one of sublimation only.

A solid evaporates, or volatilizes, because it possesses a vapour
pressure, and this vapour pressure increases with rise of tempera-
ture till the melting-point of the solid is reached. After the sub-
stance has melted, the vapour pressure, which continues to increase
with rising temperature, is, of course, that of the liquid, and not
of the solid.

The vapour pressures of many solids are very small, even near
their melting-points, but those of iodine are considerable. In the
following table are given the vapour pressures of iodine at different
temperatures.

From these figures it appears that much iodine could be
sublimed without melting at, say, 100 C., since the vapour pres-
sure at that temperature amounts to more
than 40 mm. of mercury.

It will be observed that 114'3 is the
melting-point of iodine, so that the vapour
pressures above this temperature are those
of the liquid element.

Now, suppose that the pressure of the
vapour above heated iodine is kept below
90 mm., the consequence will be that the
temperature of the iodine will never reach
114'3, its melting-point. It will therefore
be impossible to melt iodine under these conditions, and heat
energy supplied to it will cause it to sublime without melt-
ing. And, conversely, if the heated vapour of iodine is quickly

VAPOUR PRESSURES OP
IODINE

Temp. C. Press, mg. Hg.

58-1 .

4-9

75-2 .

11-5

91-9 .

29-6

102-7 .

50-7

114-3(m.-p

) 90-0

120-4 .

113-4

125-5 .

135-8

166-6

475-0

177-6 .

630-3

184-35 (b. p

) 760-0

168 CHEMICAL THEORY

cooled, so that its pressure falls below 90 mm., liquid iodine will
not appear, but crystals of the solid instead. Thus it is easy
to see that the true sublimation of iodine will necessarily occur
when it is so heated that its vapour pressure does not reach
90 mm.

These considerations will explain why, for example, arsenious
oxide, when heated, sublimes without melting. If this substance is
heated under ordinary conditions in the air, its vapour pressure
does not rise to a value corresponding with its melting-point.
When, however, this compound is enclosed in a space which i&
continuously and uniformly heated, so that instead of condensing
the vapour accumulates, fusion eventually takes place under the
increased pressure.

(6) SOLIDIFICATION OF LIQUIDS MELTING-POINTS OF SOLIDS

Water freezes and ice melts at C. under a pressure of

1 atmosphere. Alteration of pressure affects the melting-point of
ice to a very small but perceptible extent. The change amounts
to about 0-0076 C. per atmosphere, so that under a pressure of

2 atmospheres, for example, ice melts at 0-0076 C. Thus
increase of pressure causes ice to melt without the supply of heat
energy; consequently when the pressure is removed the water
formed by the melting of the ice freezes again. This phenomenon
is called regelation.

Now water expands on freezing, or ice contracts on melting^
1 volume of water at C. becomes 1-09082 volumes of ice at
the same temperature; or the density of ice at C. is 0-91674,
if that of water at C. is unity.

The lowering of freezing-point of water by compression and the
contraction of ice on melting are related phenomena. Ice melts
when compressed, because in so doing it accommodates itself to
the pressure by changing into water, which occupies a smaller
volume.

Most solids, however, e.g. mercury and phosphorus, expand when
liquefied; consequently compression promotes the formation of solid
rather than liquid; in other words, increase of pressure raises the
melting-point.

The influence of pressure upon melting-point is shown for ic&
and phosphorus in the following table:

THE PROPERTIES OF SOLIDS 169

ICE PHOSPHORUS

Pressure in Atmospheres. Melting-point. Pressure in Atmospheres. Melting-point.

1 0C 1 44-10 C.

336 2-5 50 ... . 45-50

615 5 100 47-00

1155 10 150 48-45

1625 -15 200 49-85

2042 -20 300 ... . 52-80

When a pure liquid is cooled, its temperature may fall below
the freezing-point of the liquid, or melting-point of the solid,
without the separation of solid. If this takes place, the liquid is
said to be super-cooled, 1 and is in a metastable state, 2 a state which
is disturbed by the addition of a crystal of the solid substance.
Thus, when water is being frozen in the Beckmann apparatus for
molecular weight determination by Raoult's method (q.v.), the
thermometer is often observed to fall continuously, and then to
rise again quickly to a fixed point directly any ice is formed. The
point to which the thermometer rises, and at which it remains
stationary, is the freezing-point of the water. The lower tempera-
tures show a state of supei -cooling.

Melting-points.

It has been seen that some solids, when heated, sublime instead
of melting. Many other solids, which are compounds, decompose
before a melting-point is reached. Chalk, for example, cannot be
melted, for when heated it decomposes into lime and carbon-
dioxide gas. Some salts, which contain much water of crystalliza-
tion, and are at the same time very soluble in water, melb in their
water of crystallization at comparatively low temperatures, pro-
ducing solutions. . If such a salt is said to have a melting-point, the
limitations of the case it presents should be clearly understood.
The temperature at which the salt melts will depend upon the rate
of heating, for if heated slowly it will lose much of its water
during the process, and will melt eventually at a higher tempera-
ture, or will not melt at all if too much water has been lost; and,
further, if kept at its melting-point for a long enough time, it will
solidify owing to the evaporation of the water. Soda crystals,
Na 2 C0 3 -10H 2 O, and Glauber's salt, Na 2 S0 4 -10H 2 O, show a further
peculiarity when heated. The former melts at 35-1, forming not

1 That is, excessively cooled.

9 This condition is comparable with that of supersaturation (q.v.).

170

CHEMICAL THEORY

a clear but a turbid liquid, which consists of crystals of the mono-
hydrate, Na 2 CO 3 H 2 O, suspended in a saturated solution of the
salt; the latter melts at 33, forming a solution in which crystals
of the anhydrous salt, Na 2 S0 4 , are suspended.

If it does not sublime or decompose, a pure solid substance melts
when sufficiently heated; i.e. it possesses a melting-point, viz. the
temperature at which it remains during melting, so that the heat
supplied, instead of raising the temperature of the substance, be-
comes its latent heat of fusion. So far as the chemical elements
are concerned, few of them sublime, and there is no question of
their decomposition; consequently they possess melting-points,
which, however, range between wide limits.

MELTING-POINTS OF ELEMENTS

Element.

Melting-point,
Centigrade.

Element.

Melting-point,
Centigrade.

Helium be

ow 271

Aluminium

658-7

TTvrl rofifpn

259

Calcium

810

Fluorine

223

Arsenic

850

Oxvcren

218

under pressure.

Nitrogen

210

Silver

960-5

Argon

-188

Gold

1063

Chlorine

-101-5

Copper

1083

Mercury

- 38-87

Manganese

1230

Bromine

- 7-3

Silicon

1420

Phosphorus

Nickel

1452

Potassium

62-3

Cobalt

1480

Sodium

97-5

Iron

1530

Iodine

113-5

Chromium

1615

Sulphur (rhombic)

112-8

Platinum

1755

Sulphur (monoclinic)

119-2

Titanium

1800

Tin

231-9

Boron

2200 8 -2500

Bismuth

271

Iridium

2350

Cadmium

320-9

Molybdenum

2550

Lead

327-4

Osmium

2700

Zinc

419-4

Tantalum

2900

Antimony

630

Tungsten

3400

Magnesium

651

Carbon

above 3600

The melting-points of the principal elements given in the above
table are interesting for several reasons.

First, it may be noticed that, with the exception of boron,
carbon, and silicon, the non-metals have low melting-points; whilst
the melting-points of the metals are higher, and cover a very wide
range.

THE PROPERTIES OP SOLIDS 171

If the melting-points of the elements are plotted on a vertical
axis against the atomic weights on a horizontal axis, a succession
of curves is obtained which show that the melting-points are
periodic functions of the atomic weights. Moreover, this melting-
point curve, taken as a whole, is roughly an inversion of the
atomic-volume curve, the elements with the highest melting-points,
such as iron and chromium and the metals of the platinum group,
having the smallest atomic volumes, and vice versa.

The useful applications of the various metals in the arts depend,
among other properties, upon their melting-points. Especially, the
employment of tantalum and tungsten for the metallic filaments of
incandescent electric lamps is attributable to the great infusibility
as well as electric resistance of these metals.

In the study of organic chemistry melting-points are of much
importance. A large proportion of solid organic compounds melt
at temperatures which can be registered by means of an ordinary
mercurial thermometer, and the observation of melting-point
is therefore frequently employed to characterize or identify an
organic compound, as well as to ascertain its purity. Since the
solution of a solid substance in a solvent lowers the freezing-point
of the solvent, the presence of an impurity will generally lower the
melting-point of a substance. The highest melting-point observed
after successive recrystallizations of a substance is therefore re-
garded as the true melting-point of that substance.

There is a connection between the melting-point and chemical
constitution of a compound, and in a homologous series increase in
molecular weight involves rise in melting-point. For example, the
primary alcohols have the following melting-points:

Alcohol.
CH 3 OH
CH,CH 2 OH
CH 3 (CH 2 ) 2 OH
CH 3 (CH 2 ) 3 OH

Melting-point. Difference for CH S
-94 18 o
-H2 - *
-127 -

199 ' 5
^^ \ iii

CH/CH^OH
CH 3 (CH 2 ) r OH
CH 3 (CH 2 ) 8 OH
CH 3 (CH 2 ) 9 OH
CH 3 (CH 2 ) 10 OH
CH 3 (CH 2 ),,OH
CH 3 (CH 2 ) 12 OH

36-5
- 17-9
- 5
+ 7
19
25
30-5

+ 18-6
12-9
12
12
6
5-5

The first few members of the series show exceptional relation-

172 CHEMICAL THEORY

ships; then a fairly regular melting-point difference for every in-
crement of CHo appears, but this diminishes when the higher
members of the series are reached.

When aqueous solutions of silver nitrate and sodium chloride
are mixed together, silver chloride is formed as a white precipitate
which is apparently amorphous; nevertheless, silver chloride can
occur crystalline. Again, if solutions of calcium chloride and
ammonium carbonate are mixed, a white, flocculent precipitate of
calcium carbonate is formed ; but when this precipitate is heated in
the liquid in which it has been produced, it becomes granular and
denser, having been converted, according to temperature and other
circumstances, into either anigonite or calcite. Hydrated magnesium
ammonium phosphate, MgNH 4 PO 4 6H 2 O, presents a similar pheno-
menon. When first precipitated from moderately concentrated
solutions, this salt is flocculent, but on standing it becomes crystal-
line. If, however, an attempt is made to precipitate it from a
quite dilute solution, no solid separates at once; but crystals of
perceptible size are formed after a time. This case can hardly be
distinguished from one of ordinary crystallization from aqueous
solution; and from the other facts here mentioned it may be
inferred that the amorphous state is not a permanent state, but
a condition assumed when a solid is suddenly formed, so that the
molecules have not had opportunity to marshal themselves in the
order which pertains to crystals.

There is, however, another kind of precipitate, which is gelati-
nous, and never becomes crystalline. Examples of such precipitates
are hydrated alumina and silica. These are colloidal substances,
consisting of large and complex molecules, and are probably not
true solids at all.

Another example of an amorphous substance is glass, which
appears solid enough. When glass is melted, it consists of a solu-
tion of various oxides, and this solution solidifies as a whole, when
cooled, without crystallizing. Glass may therefore be regarded as
a solid solution. Since, however, glass is slightly flexible, and
gradually becomes more flexible as it is heated, it may be ques-
tioned whether this substance should be regarded as a solid or
a super-cooled liquid solution. Support is lent to the latter view
by the tendency of glass to devitrify, that is, for crystals to

THE PROPERTIES OP SOLIDS 173

separate within it, just as crystals may separate within a liquid
solution.

With regard to the general question of amorphous solids, the
extreme view is sometimes taken that no pure chemical substance
exists in the solid amorphous condition, but that all amorphous
substances are in reality super-cooled liquids^

Solidification of Mixtures Cryohydrates Eutectic Mixtures.

Liquid mixtures are solutions, and when solutions are cooled
sufficiently for solids to separate from them a variety of pheno-
mena may occur. It has already been seen that when a dilute
aqueous solution of a substance is cooled, ice separates from that
solution just below C., and
that the more concentrated the
solution the lower is the tem-
perature at which the water ~|
freezes. On the other hand, from ,

a warm concentrated solution it

is the dissolved solid, i.e. the
solute, rather than the solvent _ 40

that separates on cooling, pro- c

i j XT r . , in % of Sodium Chloride

vided the tormer is less soluble

in cold than in hot water. Now

since the separation of ice from an aqueous solution increases
the concentration of the solution, and the separation of solute
diminishes its concentration, it is clear that two solutions of the
same substance from which ice and solute separate respectively on
cooling continually approach each other in concentration, so as
eventually to become identical. This condition of identity has
been approached from either side by continuous cooling, and the
curves connecting concentrations and temperatures at which solid,
whether ice or solute, separates, approach each other and eventu-
ally meet in a point, shown in the case of sodium chloride in the
accompanying figure.

The curve AC represents the temperatures at which ice
separates according to the percentage of salt in the solution;
similarly, the curve EC represents the temperatures at which
salt separates from the cooling solution that it may remain
saturated. Between +0-15 C. and - 22 C. it is the hydrate
NaCl-2H 2 O which is formed in the latter case. These two

10 20 30

174

CHEMICAL THEORY

approaching curves meet at the point C, which corresponds to
22 C. and 23-6 per cent of sodium chloride. Under these con-
ditions the solution solidifies as a whole, for changes in concen-
tration owing to separation either of ice or of salt are at an end.
This state, moreover, is reached whatever the concentration of
the original solution may have been; for change in original con-
centration will affect only the point on either curve at which
separation of solid begins. The solution in the state shown at C
is a liquid with a definite solidifying-point, and the solid produced
from it is one with a definite melting-point. These seem to be
the properties of a chemical compound rather than of a mixture;
but the solution is a mixture, for a 23-6 per cent salt solution
cannot be regarded as a chemical compound; and the solid formed
from this solution is a mixture of salt and ice, as may be revealed
by the microscope.

The phenomena here described were first observed and studied
by Guthrie, who called the solid produced when a solution solidifies
as a whole a cryohydrate. 1 Since, however, the cryohydrate of a
particular substance has a constant composition, Guthrie regarded
such a product as a chemical compound. Clearly, however, he was
wrong.

In the following table are given a few data regarding cryo-
hydrates:

Salt.

Cryohydric Point,
Centigrade.

Percentage of
Anhydrous Salt in
Cryohydrate.

Sodium chloride...

23-60

Potassium iodide

-22

52-07

Sodium nitrate ...

-17 -5 r

40-80

Ammonium sulphate .
Ammonium chloride

-17
-15

41-70
19-27

Magnesium sulphate .
Sodium sulphate

- 5
- 0-7

21-86
4-55

The point C in the cryohydric diagram (fig. 25) marks the lowest
temperature at which solidification takes place, and the solid
mixture formed at that temperature is the easiest melted of
any possible mixture. The point is therefore called the eutectic 2
pointy and the mixture which solidifies is a eutectic mixture.

1 1.6. ice-hydrate, because of the low temperature at which it is formed.
2 Eutectic = easiest melted.

THE PROPERTIES OF SOLIDS

175

The phenomena here studied are characteristic of other mixtures
besides aqueous solutions; they are of particular importance in the
case of alloys which are mixtures of metals. First, it may be noted
that, in accordance with principles with which the student is now
familiar, mixtures of metals melt at lower temperatures than their
constituents. A mixture of sodium and potassium may be made
which is liquid at atmospheric temperature; and Wood's fusible
metal is composed and melts as follows:

1 part tin

? par * le * d .
1 part cadmium

4 parts bismuth

WOOD'S FUSIBLE METAL
M.-p. 232 C. ^

?J I Alloy melts at 6Q.5' 0.
.. 321 C. j

.. 271 C. J

A molten mixture of zinc (melting-point 419 C.) and cadmium
(melting-point 321 C.) yields, by solidification of one or other
metal during cooling, a eutectic mixture which solidifies at 270 C.
and contains 73 per cent of cadmium. These facts are expressed
diagrammatically thus:

419

10096
Zn

321

270

<3 "0
Cd

100,0
Cd

Fig. 26

Crystals and Crystallography.

According to popular ideas, a crystal is a clear and sparkling
solid. 1 These ideas are true as far as they go. A well-formed
crystal is transparent because its parts are packed in a close and
orderly manner 2 without interspaces; and it often sparkles because

1 Gr. KptfaroAAo*, ice.

2 The influence of close-packing is well seen by comparing together ice and snow. Both
consist of crystals, but those of ice, which in small quantity is colourless and transparent, are
close-packed without interspaces, whilst those of snow, which is opaque as well as white, lie
loosely upon one another with air between.

176 CHEMICAL THEORY

its faces are flat, and reflect incident light without scattering it
in innumerable directions.

To these ideas, however, must be added a third, also popularly
recognized: that of geometrical form. Consequently a crystal is a
solid body, formed naturally from a state of vapour, liquid, or solution,
which is bounded by plane surfaces, pairs of which intersect at edges
which are straight lines, and three or more in solid angles or coigns.

Whilst this definition serves for obvious crystals, it does not
suffice to describe the crystalline state. A piece of rock-crystal, 1
for example, may plainly show a crystal form, but parts of it
may be found which are irregular in shape; a lump of ice or
rock-salt may be amorphous in external shape, whilst it is un-
doubtedly crystalline in internal structure. This may be shown
in the case of rock-salt by breaking a lump with a hammer;
among the fragments may be found almost perfect cubes; so
the crystalline structure is revealed by fracture. When, however,
a lump of rosin is broken by a hammer, or a pane of glass by
means of a stone, the kind of fracture produced is dependent on
the nature of the blow and not on the internal structure of the
solid; for these are truly amorphous substances. Internal structure,
however, as shown by fracture, is not alone a criterion of crystal
nature; for slate and coal, for example, have a laminar fracture
owing to being formed in strata by compression,

A fundamental distinction between crystalline and amorphous
substances is well illustrated by the different behaviour of glass
and rock-crystal towards light. Light travels through glass at
the same rate in all directions; it is similarly refracted from
whatever point it enters the glass. This, however, is not so with
quartz; for if a ray of light enters a quartz-crystal in any direction
except one it is doubly refracted, being thus divided into two plane-
polarized rays; whilst in one direction only, parallel to what is
called the optical axis, the ray is not doubly refracted. 2

Thus crystals differ in optical as well as in other physical
properties in different directions; and this fact is a criterion of
crystal structure which is independent of external form.

It must be added, however, that crystals belonging to one sys-
tem the regular system show similar properties in all directions.

1 Rock-crystal, because of its icy appearance, was originally thought to have been formed
by cold amid the Alps.

8 Modern work on crystal structure is referred to on p. 128.

THE PROPERTIES OF SOLIDS

177

There are, on the other hand, liquids which differ in physical
properties in different directions; and consequently the existence of
liquid crystals has been recognized.

For the scientific classification and study of crystals there must
be some system, and crystallographers have adopted the geomet-
rical system of axes round which the crystal is supposed to be
built up.

Consider, for example, a crystal such as that shown in fig.
27, which is the form in which the alums crystallize.

It is a double square pyramid, that is, a pyramid on a square
base ABCD, together with what would appear as its image if it

Fig. 28

stood upon a mirror. The equal diagonals of this square, AC and
BD, bisect each other at 0, and through this point, which is called
the origin, passes the line EOF, which is equal to AC and BD, and
is also bisected at 0. The double-pyramid has eight faces, and is
consequently called an octahedron. These faces are equilateral
triangles, and they are all equal; moreover, the six angular points
of the solid figure are equidistant from the centre or origin, O;
consequently the figure is exceedingly symmetrical. Any of the
other five angular points might take the place of the point F
without altering the appearance of the figure. For instance, A
might take the place of F; BEDF would then become the central
plane, but this figure is also a square with the equal diagonals,
BD and EF, bisecting one another at the point O. The lines AC,
BD, EF are the crystallographic axes around which the regular
octahedron appears to be built up.

It is possible, however, for a quite different figure to be con-
structed around the same crystallographic axes. This is the figure

(l>00) 13

178

CHEMICAL THEORY

of a cube (fig. 28). Here it is seen that the three equal axes emerge
from the crystal at the centres of the six square faces of the cube.

These two forms, the double-pyramid and the prism (a cube
is a square prism), are fundamental, and occur with modifications
in other crystal systems.

Crystals are seldom as perfect as these figures represent them
to be. Frequently some faces are developed
by unequal growth at the expense of others,
so that the form appears distorted. Thus
alum may grow irregularly, and produce,
instead of a regular octahedron, a form
such as that shown in fig. 29, where the
smaller faces are farther removed from the
centre of the crystal than the larger ones.
There can be no doubt, however, as to the
type to which such a crystal belongs, for
the existing faces are parallel to the corresponding faces of the
perfect octahedron, being inclined to one another at the same angles.
Sometimes alternate faces are so developed as to obliterate the
other faces; the crystal consequently possesses only half the number
of faces of the original form. This is seen in the following illus-
tration (fig. 30), which shows a regular tetrahedron, that is, a figure

Fig. 29

Fig. 30

Fig. 31

with four equal faces in relation to the octahedron from which it
is supposed to be derived.

Sometimes the other four faces of the octahedron begin to-
appear at the corners of the tetrahedron, as is shown in fig. 31.
If these were so far developed as to yield faces equal to the
faces of the tetrahedron reduced by their development, the octa-
hedron would be regenerated.

The octahedron, with the full complement of faces, is said to be

THE PROPERTIES OF SOLIDS

179

a holohedral form, to which the tetrahedron is the corresponding
hemihedral l form.

Very frequently, also, actual crystals are found to combine two
forms, e.g. the pyramid and prism. The follow-
ing figures show the ideal form of quartz, which
combines the hexagonal prism and pyramid, and
a form in which quartz more often occurs, show-
ing distortion by unequal growths in different
directions.

The following crystallo-
graphic systems are recognized.
Unfortunately one of several
names may be employed for
each system.

Fig. 32

Crystal Systems.

i. REGULAR SYSTEM (Cubical,
Isometric, Monometric, Tesseral). 3 axes, all equal, and at right
angles (fig. 33).

Examples. Alum, salt, diamond, fluor-spar, various metals.

Fig 33

rig. 34

ii. QUADRATIC SYSTEM (Square Prismatic, Dimetric, Tetragonal).
3 axes, 2 equal, all at right angles (fig. 34).

Examples. Zircon (ZrSiO 4 ), tin-stone (Sn0 2 ), potassium dihy-
drogen phosphate, potassium ferrocyanide.

i Holohedral = full number of faces ; hemihedral = half number of faces.

180

CHEMICAL THEORY

iii. RHOMBIC SYSTEM (Trimetric, Orthorhombic). 3 unequal
axes, all at right angles (fig. 35).

Examples. Sulphur, nitre, potassium sulphate, magnesium

sulphate, barium sulphate, potassium
perchlorate, potassium permanganate.

Fig. 35

Fig. 36

iv. MONOCLINIC SYSTEM (Monometric, Monosymrnetric, Clino-
rhombic). 3 unequal axes, 2 at right angles, the third inclined
(fig. 36).

Examples. Gypsum, felspar, borax, soda crystals.

v. TEICLINIC SYSTEM (Asymmetric, Anorthic). 8 unequal axes,
all inclined, there being no right angle
between them (fig. 37).

Examples. Copper sulphate, potas-
sium dichromate, boric acid.

Fig. 37

Fig 33

vi. HEXAGONAL SYSTEM. 4 axes, 3 equal, in one plane, forming
a regular hexagon; the fourth perpendicular to the plane of the
other three (fig. 38).

Examples. Quartz, beryl, apatite, lead iodide.

THE PROPERTIES OF SOLIDS 181

Hemihedral forms: calcite, corundum, graphite.

NOTE. The hemihedral forms of the hexagonal system are
often classified as a separate system, the Bhombohedral System
(Trigonal).

Crystallization.

It has already been seen that crystals may be produced from
a substance in the state of vapour, as when snow is formed from
water vapour, or iodine and naphthalene by the cooling of their
respective vapours; also, that a pure liquid substance may solidify
in the crystalline form, as when water freezes, or molten sulphur
becomes solid. More frequently, however, crystals are obtained by
the separation of a substance from solution in water or other
solvent.

Inorganic salts are crystallized from aqueous solution; and if
a salt is more soluble in hot water than cold, advantage is taken
of this fact by preparing a hot saturated solution of the salt, and
cooling the solution to atmospheric temperature. The excess of
the salt, which can be dissolved in hot water over what can remain
in solution at atmospheric temperature in contact with the solid
salt, then separates in crystals, the size of which depends upon their
rate of formation.

An experiment with copper sulphate will illustrate the forma-
tion of crystals. Suppose a hot saturated solution of this salt is
rapidly cooled. As the solution approaches atmospheric tempera-
ture a mass of minute crystals separates, and if these are quickly
filtered off, the nearly cold solution will slowly deposit more
crystals as it stands. These crystals, moreover, will be larger, and
will consequently appear deeper in colour than those which compose
the crystal-meal that separated at first. Indeed, the more slowly
a crystal forms the larger will it be; large crystals, like large trees,
cannot be grown in a short space of time. On account of the long
time available for the natural growth of crystals, mineral sub-
stances often occur in large crystals such as cannot possibly be
made artificially. For example, calcium carbonate and barium
sulphate occur naturally in very large transparent crystals, whilst
artificially produced crystals of these substances are little more
than microscopic in size.

The largest and most perfect crystals of a salt are best produced
artificially by allowing its saturated solution to evaporate spon-

1& CHEMICAL THEORY

taneously at atmospheric temperature in contact with some small,
well-formed crystals of the salt. The rate of deposition of solid
is commensurate with that of the evaporation of the water, since
the solution must remain saturated; and if the growing crystals are
turned over from time to time, to prevent disproportionate growth,
large, well-formed crystals are eventually obtained. The mother-
liquor, however, should not be allowed to evaporate completely, or
the crystals will be contaminated with any impurity present in
solution.

Recrystallization,

A crystalline salt may be purified by recrystallization, i.e. by
dissolving it in hot water, filtering the solution if necessary,
and allowing it to cool, so that crystals separate again. The
soluble impurity will then probably remain in the mother -liquor,
especially if it is present in relatively small amount, or is more
soluble in water than the salt that is being recrystallized. Successive
recrystallizations enhance the purity of the crystallized substance,
but necessarily reduce its amount very seriously, since much of it
remains in the successive mother -liquors. When salts are isomor-
phous, they tend to crystallize together, and their separation by
recrystallization is therefore difficult. Copper sulphate and ferrous
sulphate thus form isomorphous mixtures; but in this case the
iron may be oxidized by means of a little nitric acid to non-
crystallizable ferric sulphate, and the copper sulphate subsequently
be obtained in crystals free from iron.

Polymorphism and Allotropy,

The phenomena of isomorphism have already been studied,
and the law of isomorphism has been stated thus: the molecules
of isomorphous substances contain equal numbers of atoms, which
when not identical are analogous. Polymorphism has also been
noticed, and it has been seen that polymorphous substances are
those which, whilst preserving their chemical identity, can assume
two or more crystalline forms under different conditions. Thus
calcium carbonate is dimorphous, crystallizing as calc-spar in the
hexagonal system, and as aragonite in the rhombic system. Titanic
oxide, Ti0. 2 , is trimorphous, crystallizing as rutile and anastase
in two different forms of the quadratic system, and as brookite
in the rhombic system. Ammonium nitrate is tetramorphous,

THE PROPERTIES OP SOLIDS 183

crystallizing in the regular system, the hexagonal system, and
two forms of the rhombic system.

Conditions of temperature and state of solution determine
which crystalline form is assumed; thus which of the two rhombic
forms of ammonium nitrate separates from solution depends on
whether the temperature is above or below 32 C.; and, moreover,
crystals of sodium chloride, which are cubical when grown from
water only, are octahedral when they separate from an alkaline
liquid; and crystals of arsenious oxide consist of regular octahedra
when formed by sublimation or from a solution in hydrochloric
acid, and of rhombic prisms when produced from a caustic soda
solution.

Various elements, and particularly some non-metals, resemble
compounds in their polymorphism. Thus solid sulphur exists
in two well-defined crystalline forms, belonging to the rhombic
and monoclinic systems; carbon in two crystalline forms as well
as in the amorphous state, and so on. Certain elements, however,
exist in different forms in the liquid or gaseous state, and to these
forms the term polymorphism cannot be applied. Thus there are
two forms of liquid sulphur, one of which is yellow and mobile,
and the other brown and viscous; these forms are observed
successively when sulphur is carefully heated in a test-tube. And
there are two forms of gaseous oxygen, ordinary oxygen, O 2 , and
ozone, O 3 ; and two forms of gaseous nitrogen, inactive and active
respectively.

When an element can exist in more than one physical form, these
forms are said to be allotropic 1 ; they are allotropes, and the phenomenon
is called allotropy or allotropism.

Thus ozone is said to be an allotropic form of oxygen; rhombic
and monoclinic sulphur are allotropic forms of that element, and
so forth. The term allotropy was given by Berzelius, in 1841, to
the relation between different forms of an element to which it
was thought the term isomerism could not well be applied. Since
it includes such diverse phenomena as those presented by gaseous
oxygen and liquid as well as solid sulphur, the term is in one sense
an inclusive one, perhaps too inclusive to be scientific; and, more-
over, since it excludes the phenomena presented by mercuric
iodide, for example, because this substance is a compound, its
scientific value is further discounted; for sulphur and mercuric

IAA<K, another; rpoirdf, habit.

184 CHEMICAL THEORY

iodide are dimorphous substances which, as will appear, closely
resemble one another in their transformations.

On the other hand, since the term polymerism denotes atomic
condensation within the molecule, so that substances which possess
the same quantitative composition differ in molecular weight and
chemical properties (e.g. C 2 H 2 and C 6 H 6 ), there is no valid scientific
reason why the ozone-oxygen relationship should not be regarded
as one of polymerism.

In view of these considerations, attempts to give a scientific
definition of allotropy are not very successful. Ostwald directs
attention to the energy aspect of the question in the following
words: "Elements which by reason of different energy content
possess different properties are allotropic". This conception is of
value, and is particularly applicable to the ozone-oxygen relation-
ship in connection with which it is developed, because individual
and simple molecules are the units in which the phenomenon is
displayed; but in the case of the dimorphism of sulphur the con-
ception of the molecule is subsidiary to that of the crystal, and
to refer the phenomenon to a difference of energy content alone,
whether molecular or otherwise, is scarcely satisfying.

Again, a criticism has been passed on this definition that it
would class as allotropic the physical states of solid, liquid, or
vapour in which an element may exist, since the different pro-
perties of a substance in these states depend upon differences of
energy content.

So that, whilst the term allotropy serves usefully to draw
attention to the fact that various elements can exist in different
forms, it cannot be defined with scientific accuracy, since the
category it creates is imperfect by reason both of what it includes
and of what it excludes.

It will now be profitable to study the conditions of change
of one solid allotrope into another, or more properly of one form
of a dimorphous substance into the other. The case of rhombic
and monoclinic sulphur furnishes a useful example.

Rhombic or octahedral sulphur crystallizes from a solution
of sulphur in carbon disulphide, and is otherwise formed at or near
atmospheric temperature. Monoclinic or prismatic sulphur is
produced at elevated temperature, as, for example, under the
crust formed upon molten sulphur as it cools, and after some of
the liquid has been drained away. It is significant, however, that

THE PROPERTIES OF SOLIDS 185

when kept at atmospheric temperature, prismatic sulphur gradually
changes into the octahedral variety; and that naturally occurring
sulphur is always octahedral. So it is concluded that whilst
prismatic sulphur is the more stable form at certain elevated
temperatures which favour its formation, octahedral sulphur is
the more stable form at atmospheric temperature. Investigation
has shown that a definite temperature exists above or below
which the prismatic or octahedral form is respectively the more
stable. This temperature is 95-5 C. Above 95-5 C. the prismatic
form is the more stable, below this temperature the octahedral
form; at this temperature both forms are equally stable. Conse-
quently 95 5 C., which is the limiting temperature for the
stabilities of the two forms of sulphur, is called the transition
temperature, and the relationships with regard to this temperature
are expressed thus:

95-5C.
Octahedral sulphur -^-^ prismatic sulphur.

Across this temperature transition takes place from one form
to the other, and the change, like many chemical reactions, is
reversible. Such a change, with reference to dimorphous substances,
is called an enantiotropic change.

Mercuric iodide presents phenomena closely analogous to those
of sulphur. It exists in scarlet and yellow forms, the scarlet form
consisting of quadratic (tetragonal), the yellow of rhombic crystals.
When the scarlet crystals are gently heated, they turn yellow, melt,
and yield a yellow sublimate, which, when cold, reverts to the red
form, reversion being hastened by scratching. The transition
temperature is 127 C., and the change is enantiotropic, thus:

127 C.
Scarlet (quadratic) mercuric iodide -~^ yellow (rhombic) mercuric iodide.

It is possible for yellow mercuric iodide to exist below 127,
though its existence is transitory. Thus, when formed by heating
the red variety, it does not immediately become red when cooled
below 127; also, if the precipitation of mercuric iodide by the
addition of potassium iodide to mercuric chloride solution is
carefully observed, it will be seen that the precipitate, when first
formed, is yellow, but that it quickly becomes red and eventually
scarlet. And if mercuric iodide is dissolved in hot alcohol, and the

186 CHEMICAL THEORY

solution is then poured into water, the precipitate when first formed
is pale yellow, but on keeping gradually turns red. When a
substance persists in a state after the conditions for its stability
have been exchanged for those under which another state of the
substance is stable, the substance is said to be in a meta-stable
state.

The relationship between yellow and red phosphorus is ap-
parently different from the foregoing. Yellow phosphorus changes
very slowly into the red variety when kept at atmospheric tem-
perature, and the rate of change is accelerated by heating, becoming
fairly rapid at 240-250 C.

Evidently, then, the red variety is the more stable at compara-
tively low temperatures. It is formed from the white variety with
the evolution of 27,300 calories per gramme-atom (=31 grm.);
thus

nP 4 + 4 P n + 27,300 X 4?i calories.

If, however, red phosphorus is heated above 250, out of contact
with air, in order to reach a temperature at which it may pass into
the yellow variety, it is found to sublime at about 290, producing a
vapour which condenses into yellow phosphorus. This phenomenon
cannot, however, be regarded as the transformation of red into
yellow phosphorus, since vapour intervenes; such transformation
does not take place under ordinary conditions. Inasmuch as
the change in the phosphorus takes place in one direction only,
it is said to be monotropic.

It may be questioned, however, whether the failure to reverse
the change in the case of phosphorus is not due to a failure to
realize the right conditions. Chemical changes sometimes appear to
be irreversible from this cause. That this is the case here is shown
when red phosphorus is heated under pressure in a sealed tube.

Being unable to sublime, the red phosphorus forms a yellow
liquid at 610, from which, on cooling, red particles begin to
separate at 580, whilst the whole turns red at 570. Thus the
change from yellow to red phosphorus is seen to be reversible
under pressure at about 600 C., which may be regarded as the
transition temperature, so that under these conditions the change
is enantiotropic.

In the following table are gathered together the chief examples
of allotropy among the elements.

THE PROPERTIES OF SOLIDS

187

Element. Forms.

Oxygen ... Oxygen, 2 ^^ Ozone, 3 .
Boron Amorphous (brown powder), density 2 45 (impure l ).

Crystalline (adamantine), density 2*34.
Carbon . . . Amorphous, density 1 45-2 0.

Graphite, density 2-2*25.

Diamond, density 3*5.
Silicon . . . Amorphous, density 2 35.

Crystalline, density 2-49.

Sulphur ... Rhombic (octahedral) -^-^ monoclinic (prismatic), 95- 5.
density 2 06 density 1 96

Also amorphous, colloidal, and liquid S (2 forms).
Phosphorus ... Crystalline (yellow) * amorphous (red),
density 1 83 density 2 106

Reversible at 600 under pressure.

" Metallic " phosphorus (impure).
Arsenic ... " Metallic " (steel grey), density 5*727.

Black or mirror arsenic, density 4-713.

Yellow arsenic, density 3*88.
Antimony ... Metallic, black and yellow forms, comparable with those

of arsenic.
Tin Grey tin ^~^ tetragonal tin, 20 C.

Tetragonal tin ^^ rhombic tin, 161 C. 1
Iron a-iron * /3-iron, 770.

j8-iron * 7-iron, 890.

The following compounds present analogous phenomena:

Ammonium nitrate

Mercuric iodide ...
Potassium nitrate
Silver iodide

a-rhombic ^
j8-rhombic ^
rhombohedral
tetragonal ^
rhombic ^-^
hexagonal ^

/3-rhombic, 32-2*.
rhombohedral, 83.
^= regular, 125-6.
rhombic, 127.

rhombohedral, 129*5,
regular, 146-147.

There are certain metallic oxides which are well known to
undergo temporary change of colours when heated. Amongst
these are HgO, ZnO, Pb 3 4 . Such changes are enantiotropic,
though the transition temperatures are not in all cases known.
Red lead, for example, blackens when heated sufficiently, but, if the
compound is not decomposed, the original colour is restored on
cooling. The transition temperature has in this case been found
to be 580 C.

1 Since " amorphous boron " contains 4 to 5 per cent of oxygen as suboxide, it can hardly
be regarded as an allotropic form of the element.

CHEMICAL THEORY

CRYOHYDRATE; EUTECTIC. A cryohydrate is a solidified mix-
ture of solute and solvent (water), having the same composition as
the solution.

A eutectic mixture, or eutectic, is a solidified mixture of solute
and solvent which has the lowest melting-point. A cryohydrate
is a special case of a eutectic.

A eutectic point is the point on a graph where the curves
representing separation of solute and solvent intersect. The tem-
perature thus represented, being the lowest temperature at which
solid can be formed, is the eutectic temperature, and the mixture,
the composition of which is also represented, is the eutectic mixture.

CRYSTALLOGRAPHY. A crystal is a solid body, formed naturally
from a state of vapour, liquid, or solution, which is bounded by
plane surfaces, pairs of which intersect at edges which are straight
lines, and three or more in solid angles or coigns.

CRYSTALLOGRAPHIC SYSTEMS. 1. Regular System. 3 axes, all
equal, and all at right angles.

2 Quadratic System. 3 axes, 2 equal, and all at right angles.

3. Rhombic System. 3 unequal axes, all at right angles.

4. Monoclinic System. 3 unequal axes, 2 at right angles,
the third inclined.

5. Triclinic System. 3 unequal axes, all inclined, there being
no right angle between them.

6. Hexagonal System. 4 axes, 3 equal, in one plane, forming
a regular hexagon, the fourth perpendicular to the plane of the
other three.

ALLOTROPY. When an element can exist in more than one
physical form, these forms are said to be allotropic, i.e. they are
allotropes, and the phenomenon is called allotropy or allotropism.

CHAPTER X
SOLUTIONS

In the broadest sense a solution is a homogeneous mixture of
two or more substances; it may therefore be a gaseous, liquid, or
solid mixture. Thus a mixture of gases such as the atmosphere
is a gaseous solution; a mixture of solids such as an alloy or a
glass is a solid solution. The term solution is, however, frequently
restricted to liquid mixtures, and in this sense it will be employed
here.

A solution, then, is a homogeneous mixture of a liquid with
a gas, another liquid, or a solid. The dissolving liquid is called
the solvent, and the dissolved gas, liquid, or solid the solute, the
mixture of solvent and solute thus becoming the solution.

That a solution is merely a mixture may not be credited without
demur, for there are numerous examples of chemical change accom-
panying the act of solution. A solution of sulphur dioxide gas
in water, for example, is a solution containing sulphurous acid, and
not merely a mixture of sulphur dioxide and water; and a solution
of calcium phosphate in hydrochloric acid contains calcium chloride
and phosphoric acid rather than the original salt. It is not, how-
ever, the process of solution that is being considered, but the product;
not the change necessary to produce the mixture, but the mixture
produced.

There is, however, a further objection to the idea that a solution
is only a mixture. Is there no chemical union between the solvent
and the solute? The act of solution of a substance in water is
generally accompanied by lowering of temperature, due to the
absorption of heat of solution incident upon the change in the
physical state of the solute; but when sulphuric acid is mixed with
water, much heat is evolved, and this probably results from some
kind of chemical union between the acid and the water. Again, from
a concentrated solution of copper sulphate, crystals of the hydrated
salt CuS0 4 5H 2 separate. It may be believed that this compound

180

190 CHEMICAL THEORY

exists in solution just previous to its separation as crystals, but a
different view of the nature of a dilute solution of copper sulphate
appears justified by the properties of such a solution. Therefore
some kind of chemical change seems to take place during the
making of a dilute solution of this salt from the solid and water.
The further consideration of the phenomena presented by the
dilute solutions of salts may, however, be deferred till a little later.

Solutions of Gases in Liquids

Aqueous solutions of gases are familiar. Drinking-water con-
tains dissolved air; effervescing beverages contain dissolved carbon
dioxide; chlorine, ammonia, and hydrogen sulphide solutions are
common in the laboratory.

A gas does not dissolve to an unlimited degree in a liquid, and
the extent to which it dissolves depends upon four conditions:

i. The nature of the solvent.

ii. The nature of the gas.

iii. The temperature of the solvent,

iv. The pressure of the gas.

The solubility or coefficient of solubility of a gas in a liquid is
defined as the volume of the gas which is dissolved by unit volume
of the liquid. The temperature must be stated, since this influences
the amount of gas dissolved, but the question of pressure does not
enter into the definition for a reason which will appear later.

i. THE NATURE OF THE SOLVENT.

The influence of the solvent is illustrated by the following
figures relating to carbon dioxide:

Vol. C0 2 dissolved by 1 vol.
Solvent. Solvent at 25 O.

Water 0-8256

Ethyl alcohol, 97 % 2-706

Chloroform 3 430

Methyl alcohol 3-837

Acetone 6-295

ii. THE NATURE OF THE GAS.

Gases differ widely in their solubility in water or other solvent.

The solubility of a gas may be measured by the apparatus
known as an absorptiometer, and shown in the accompanying
figure.

The pure, dry gas is contained over mercury in the graduated

SOLUTIONS

191

tube A, and its pressure is regulated by means of the tube B
connected with A by thick rubber tubing. At the top of A is a
3-way tap leading to the atmosphere, or to the flexible tube C
made of lead, and connecting the measuring-tube with the absorp-
tion vessel D. This vessel is at first completely filled with air-
free water or other solvent for the gas. Some of this solvent is
then made to flow out through the tap F, whilst an equal volume
of gas enters the absorption vessel from the
top, until the liquid level stands at E, the
known volume of the remaining solvent being
that which is to be saturated with the gas.
The gas is agitated with the liquid repeatedly
until no further diminution of volume in A is
observed. The remaining gas is then taken
back into A by opening the tap F under mer-
cury, and so adjusting the pressure as to cause
mercury to enter, and the liquid consequently
to rise to the top of D, so as to displace all
the gas. The volume of gas which has been
dissolved by the known volume of liquid in
D is then ascertained by observing the volume
remaining in A at atmospheric pressure. If
desired, the absorption vessel may be surrounded by a water-jacket
to keep its temperature constant.

In the following table are given the solubilities in water of some
of the commoner gases, together with their boiling-points:

Fig. 39

Volume of Gas dissolve 1

B.-p. of Liquefied

Gas.

by 1 volume

Gas under

of Water at C.

Atmospheric Pressure .

Hydrogen
Nitrogen

0-021
0-0235

-252 -^
-195-7

Carbon monoxide

0-0328

-190

Oxygen

0-0489

-182-9

Nitrous oxide

1-305

- 89-8^

Carbon dioxide

1-713

- 80

Hydrogen sulphide
Chlorine

4-37
4-610

- 61-8
- 33-7

Sulphur dioxide

79-789

- 10-1

Hydrogen chloride

503-0

- 83

Ammonia

114S-0

- 33-5

It will be noticed that these gases fall conveniently into three

192 CHEMICAL THEORY

categories. First, there are the gases hydrogen, nitrogen, carbon
monoxide, and oxygen, which are but slightly soluble in water, so
that they may be collected over water without appreciable loss,
whilst for rough experimental purposes their solubility may be
ignored. Secondly, there are nitrous oxide, carbon dioxide, chlorine
and hydrogen sulphide, the solubilities of which need to be taken
account of in the ordinary experiments of the laboratory. Nitrous
oxide is generally collected over warm water to reduce the amount
lost by solution; carbon dioxide is generally collected by air dis-
placement, although it may be collected conveniently over warm
water; chlorine and hydrogen sulphide are rather too soluble in
water to be collected over that medium without serious loss. There
remain sulphur dioxide, hydrogen chloride, and ammonia, which are
very soluble in water, and must always be collected by air displace-
ment, or over mercury or some other liquid in which they do not
dissolve.

Two interesting questions arise from the consideration of these
figures. The first is as to how far solubility in water is connected
with condensibility; and the second is as to the connection between
solubility and power to combine chemically with water.

A consideration of the boiling-points of the liquefied gases shown
in the table reveals some connection between solubility and con-
densibility. The least soluble gases are the least condensible; but
there is no quantitative connection between the two properties, for
chlorine and ammonia, for example, have practically the same
boiling-point, but widely different solubilities.

No doubt chemical union between the gas and water will have
some influence on solubility, though the precise nature of this
influence is not apparent. Of the above gases carbon dioxide,
sulphur dioxide, and ammonia combine with water to form car-
bonic acid, sulphurous acid, and ammonium hydroxide respectively,
but the proportion of combined to uncombined gas is in each
case relatively small; in the case of carbon dioxide it is less than
1 per cent. It is now believed that hydrogen chloride forms with
water the compound OH 3 C1.

iii. THE TEMPERATURE OF THE SOLVENT.

It is the temperature of the solvent, rather than that of the gas,
which needs to be considered here. If the temperatures of the liquid
and gaseous phases are initially different that of the gaseous will
soon become adjusted to that of the liquid phase, so that the only

SOLUTIONS

193

question concerns the condition of equilibrium between a liquid and
a gas, both at the same temperature.

The solubility of a gas in water or other solvent invariably
diminishes with rising temperature, though there is no simple con-
nection between temperature and solubility. As a rule, when a gas
has been dissolved in water, boiling the water serves to expel from
solution all the dissolved gas. Thus dissolved air is completely
expelled from water by boiling the water for a short time.

The following figures and accompanying curve (fig. 40) show the
solubility of carbon dioxide in water, volume for volume, at different

temperatures:

CARBON DIOXIDE

Temperature,
Centigrade.

Solubility.

1-713

1-424

1-194

1-019

0-878

0-759

Temperature,
Centigrade.

Solubility.
0-665

0-592

0-530

0-479

0-436

0-359

There is a notable difference between hydrogen chloride and
ammonia as regards effect of temperature on solubility.

HYDROGEN

CHLORIDE

AMMONIA

Temperature,
Centigrade.

Solubility.
0-825

Temperature, <
Centigrade. *

0-783

0-742

0-700

0-665

0-633

0-603

0-575

Solubility.

0-875
0-713
0-582
0-474
0-382
0-307
0-244
0-186

This difference is seen when concentrated solutions of the two
gases are boiled. Whilst ammonia is completely expelled from
water by boiling, this is not the case with hydrogen chloride.

When a concentrated aqueous solution of hydrogen chloride is
boiled under atmospheric pressure, it loses hydrogen chloride faster
than water vapour; when a dilute solution is boiled, it loses water
vapour faster than hydrogen chloride. Since, therefore, a concen-
trated solution becomes weaker by boiling, and a dilute solution
more concentrated, the two solutions approach each other in strength

(DOG) 14

194

CHEMICAL THEOKY

when boiled, and eventually become identical. The solution ob-
tained by continued evaporation of any aqueous solution of hydro-
gen chloride contains 20-24 per cent of hydrogen chloride, and
distils unchanged at 110 C. under a pressure of one atmosphere.

A constant boiling-point at a certain pressure is characteristic of
a chemical compound, and the constant boiling solution might on

ID
14
13
12
11
10
9

"3 <

6
5
4
3
2
1

(

^">

*N.

) 10 20 30 40 50 60

Temperature.
Fig. 40. Solubility of Carbon Dioxide in Water

this account be supposed to be a chemical compound of hydrogen
chloride and water. A chemical compound, however, whilst it
distils at temperatures which vary with varying pressures, yields
always the same distillate. But this is not the case with hydro-
chloric acid, for the composition of the distillate from a constant-
boiling acid depends upon the pressure under which the boiling-point
is reached. For example, the distillate from a constant-boiling acid
under 100 mm. pressure contains 22-9 per cent of hydrogen chloride,

SOLUTIONS 195

whilst that obtained under 2500 mm. contains only 18-0 per cent of
this substance.

It must be remembered that when heat is continuously applied to
a liquid or solution, the temperature must rise until it becomes con-
stant or all the liquid has evaporated. In the case of hydrochloric
acid, owing probably to chemical union (see p. 192) between hydro-
gen chloride and water, the gas is retained in quantity sufficient to
cause the boiling-point of the solution to rise to 110 C., whilst in
the case of ammonia combination must be feebler since all the gas
has left the water by the time the temperature reaches 100 C., so that
eventually, instead of a constant boiling mixture, only water distils.

iv. THE PRESSURE OF THE GAS.

That increase of pressure increases the amount of a gas dissolved
by a liquid is well known. Soda-water, for example, contains carbon
dioxide dissolved under pressure, and the escape of this gas from
solution when the pressure is released is the cause of effervescence.

The following is the law of Henry:

The quantity of a gas dissolved by a given volume of a liquid is pro-
portional to the pressure of the gas.

Since, according to Boyle's law, the volume of a gas is inversely
proportional to its pressure, it follows that

The volume of a gas dissolved by a given volume of a liquid is
independent of the pressure.

The law may be otherwise stated as follows:

The ratio of the concentration of the gas dissolved in the liquid
to its concentration in the atmosphere above is constant. Since this
constant depends upon the nature of the gas and of the solvent, it
expresses the solubility of the gas.

Now, there are deviations from Boyle's law, which are the
greater the higher the pressure and more condensible the gas. And
there are similar deviations from Henry's law, which are the greater
the higher the pressure and the more soluble the gas.

Hydrogen chloride, for example, departs widely from Henry's
law; for 1 grm. of water at C. dissolves 0-856 grm. of hydrogen
chloride when under 1000 mm. pressure, and 0-657 grm. instead of
0-0856 grm. of the same gaa under 100 mm. pressure.

The law is approximately true, not only for slightly soluble
gases, but also for more soluble gases at high temperatures. For
example, the law is true for ammonia at 100 C. and for sulphur
dioxide above 40 C.

196

CHEMICAL THEORY

Solubilities of Gaseous Mixtures.

Henry's law applies to the solution of the constituents of
gaseous mixtures, the pressures of the individual gases being exer-
cised according to Dalton's law of partial pressures, which has
already been expressed as follows: The total pressure of a gaseous
mixture is the sum of the partial pressures of the individual gases.

A combination of these two laws sets forth the behaviour of
a gaseous mixture in contact with a solvent. It is Dalton and
Henry's law.

The quantities of the various gases dissolved from a gaseous
mixture by a given volume of liquid are proportional to the solubilities
and partial pressures of the gases.

Consequently these quantities are expressed relatively by the
products of the partial pressures and solubilities, as the following
example shows.

To find the composition of the gas dissolved from air by water
at C.

Air.

Percentage
Composi-
tion.

Solubilities.

Product.

Percentage
Composition
of Gas
Dissolved.

Nitrogen
Oxygen
Carbon dioxide ...

79
21
0-04

0-024
0-049
1-71

1-896
1-029
0-0684

63-3
34-4
2-3

The products of the partial pressures, as indicated by percent-
age composition, and solubilities show the relative amounts of
nitrogen, oxygen, and carbon dioxide dissolved in the water, and
recoverable from it by boiling. These may be regarded as the
partial pressures of the recovered gases at the same volume, or,
what is equivalent, the partial volumes at the same pressure. The
figures in the fourth column, which are those in the third column
brought to a percentage, thus give the percentage composition
by volume of air boiled out of water. The greater solubility of
oxygen, and particularly of carbon dioxide, as compared with
nitrogen, accounts for the enrichment of air by these gases
through the process of solution in water; and it will be seen
that if carbon dioxide were first removed it would be possible
to obtain a small quantity of practically pure oxygen by repeating
the processes of solution and boiling a number of times.

SOLUTIONS 197

Solutions of Liquids in Liquids

Pairs of liquids may be divided into three categories, according
to their mutual solvent action.

i. LIQUIDS WHICH Mix IN ALL PROPORTIONS. Water and
alcohol, alcohol and ether, ether and olive oil, are examples of
pairs of liquids which mix in all proportions; there is no limit
to their mutual solubility; there is no question of a coefficient of
solubility. Changes of volume as well as changes of temperature
occur, however, during mixing. For example, 539 volumes of
ethyl alcohol mixed with 49-8 volumes of water, both at C.,
produce 100 volumes of mixture instead of 103*7 volumes.

Similia similibus solvuntur: like are dissolved by like. Thus
water, H-OH, dissolves alcohol, C 2 H 6 -OH, and acetic acid,
CH 3 COOH; a hydroxylic solvent dissolves hydroxylic compounds.
When, however, the radicle, combined with the hydroxyl (OH),
group is large, the alcohol or acid is not soluble. For example,
cetyl alcohol, C 16 H 33 OH, and stearic acid, C 17 H 35 COOH, do not
dissolve in water.

Further, one hydrocarbon will dissolve another; natural petro-
leum is a mixture of open-chain and other hydrocarbons; the
products of tar distillation are a mixture of closed-chain hydro-
carbons. Acquaintance with organic chemistry will reveal other
examples to the student.

ii. LIQUIDS THE MUTUAL SOLUBILITY OF WHICH is LIMITED.
When ether is added little by little to water, and the mixture is
shaken, at first the liquids mix, but soon the limit of solubility is
reached, and ether is seen to form a separate layer on the surface
of the water. When this takes place, the liquids are saturated
solutions of water in ether and ether in water respectively. At
15 C., 8- 2 grm. of ether dissolve in 100 grm. of water, whilst 1-16
grin, of water dissolve in 100 grm. of ether. The consequence of
this is that ether which has been washed with water is " wet ", and
if it needs to be dried must stand in contact with anhydrous calcium
chloride, phosphoric oxide, or metallic sodium, and then be redis-
tilled. The presence of ether in the water can be shown by freezing
the latter, and then attempting to set fire to the ice. The ether
will vaporize from the surface of the ice and burn.

A distinction must be drawn between a solution of two liquids
and an emulsion. The former is necessarily clear, the latter turbid

198 CHEMICAL THEORY

or opaque, since it consists of minute drops of one liquid suffused
through the other. When aniline or other organic liquid is dis-
tilled in presence of water, the distillate may be turbid owing to
the presence of a little water which forms an emulsion with it.
Contact with a drying agent removes the water, and clarifies the
distillate.

The distribution of a substance between two solvents is an
interesting case which sometimes arises; and in this case the
available material is distributed throughout the two solvents in
contact according to its relative solubilities in them. Thus the
ratios of the concentrations of the dissolved substance in the
solvents remains the same at a given temperature, whatever the
relative volumes of the solvents and the consequent states of
dilution of the solutions may be.

In three experiments the amounts of succinic acid found dis-
solved in unit volumes of water and ether respectively were as
follows:

Succinic acid per unit volume of

Water. Ether. Ratio = Constant.

0-024 grm 0-0046 grm 5-2

0-070 0-013 5-2

0-121 0-022 5-4

This constant ratio is called the coefficient of distribution, or
partition coefficient.

There is a close analogy between this case and that of a gas
dissolving in a solvent; for in the latter case the ratio between
the concentration of the gas in the solvent and in the space above
it is constant when equilibrium has been reached. Indeed, the
solubility of a gas in a liquid may be regarded as a partition
coefficient.

An illustration of the distribution of a substance between two
solvents arises in qualitative analysis, when bromine or iodine,
liberated from a bromide or iodide by chlorine, is removed from
water by means of carbon disulphide. The halogen is so much
more soluble in carbon disulphide at ordinary temperature than
in water that nearly all of it passes into the former solvent, leaving
little in the water. A complication arises as regards the partition
coefficient in the case of iodine, since the molecular state of this
substance differs in the two solvents.

iii. LIQUIDS WHICH ARE IMMISCIBLE. Under this category are

SOLUTIONS 199

included water and hydrocarbons such as petroleum and benzene,
and water and oils.

A practical advantage in the use of petroleum or benzene as
a solvent over ether is the fact that the former is essentially
anhydrous, whilst ether may be "wet". Even oils, however, are
not absolutely insoluble in water, for essential oils, when shaken
with water, impart their odour to the water.

The Distillation of Mixed Liquids Fractional Distillation.

A liquid boils when its vapour pressure has attained to the
pressure of the superincumbent atmosphere, and this is true
whether the liquid is a single substance or a mixture. Consider
the case of a mixture of ethyl alcohol boiling at 78 C., and water
boiling at 100 C.

Ethyl alcohol alone would exert a pressure of 760 mm. at 78,
whilst the vapour pressure of water at 78 would be 327 mm.
Neither of these liquids, however, can exert its full vapour pressure
in the mixture, for the vapour pressure of a solvent at any tem-
perature is lowered by a dissolved substance; and so alcohol and
water mutually lower each other's vapour pressure. If the pro-
portion of water to alcohol is small its effect in lowering the vapour
pressure of the alcohol will be proportionately small, whilst the
larger proportion of alcohol with which it is mixed will have a
proportionately great effect in reducing the vapour pressure of
the water. Consequently the mixture will begin to boil slightly
above the boiling-point of alcohol, and the first portions of the
distillate will be nearly pure alcohol. The boiling-point, however,
will gradually rise, and the proportion of water in the distillate
correspondingly increase until the boiling-point is practically 100
and the distillate nearly pure water. If, however, the distillation
is stopped before the alcoholic distillate becomes much diluted with
water, and if the residue is then rejected and the distillate partially
redistilled, a distillate still richer in alcohol will be obtained. So
by repeated distillations the alcohol, though reduced in quantity,
may be nearly freed from water.

The process of fractional distillation is generally carried out as
follows: Fractions A, B, C, D, &c., are collected at small tempera-
ture intervals, and the residue is put on one side. A is then redis-
tilled up to, say, 1 above the boiling-point of the more volatile
constituent, thus yielding a distillate A x ; to the residue of A, B is

200

CHEMICAL THEORY

then added, and the distillation continued up to the same tem-
perature, the distillate Bj being collected. By a similar procedure
distillates G l and D t may be obtained, and the process repeated
if desired with A a , B 1? C v D r Thus a fairly efficient separation of
the more from the less volatile constituent of the mixture is
effected. The residues from the successive
distillations, if mixed and distilled, will yield
the less volatile constituent nearly pure after
the rejection of a small intermediate fraction
containing what remains of the more volatile
constituent.

A more perfect separation results if a
dephlegmator or fractionating-column is em-
ployed. This consists of a tube fixed between
the distilling-flask and the condenser, in which,
by means of disks and rods or other means, a
large surface is provided on which the vapour
may be cooled. Some of the vapour of the
less volatile constituent is thus continuously
condensed, and runs back into the distilling-
flask, the ascending vapour passing through
a certain amount of liquid in the column;
consequently only the most volatile part of
the vapour passes forward into the condenser.
The difficulty of obtaining pure alcohol
from a mixture of alcohol and water is in-
creased when the proportion of alcohol to
water is small. The vapour of the alcohol is
held back by the water, and at the same time
the alcohol has only a small effect on the
vapour pressure of the water. Hence the
temperature at which the liquid first boils is
considerably above the boiling-point of alcohol, and the distillate
at once contains much water. On the other hand, if it is desired
to get rid of the alcohol, leaving behind only water, this may be
effected by heating the mixture on the steam bath for a time.

Solutions of Solids in Liquids

Water is the solvent almost invariably used for inorganic sub-
stances; for organic substances alcohol, ether, petroleum - ether,

Fig 41

SOLUTIONS 201

ethyl acetate, chloroform, benzene, and other solvents may be em-
ployed. Similia similibua solvuntwr again applies. Water dis-
solves acids, bases, and many salts, with or without water of
crystallization, as well as some organic compounds containing OH
groups, such as the lower alcohols, acids, glycerine, and sugars;
petroleum dissolves other hydrocarbons, ether other neutral organic
substances, and so forth.

In dealing with the solution of solid substances, the specifically
chemical action of the solvent may be disregarded for the present.
For instance, the hydrolysis of phosphorus pentachloride by water
and the consequent solution of the hydrolytic products is irrelevant
to the present study. Attention will here be confined to the action
of water as a solvent.

Solubility.

The solubility of a solid substance in water is defined as the number
of grams of the substance which will dissolve in 100 grm. of water.

This definition involves the idea of saturation. When a sub-
stance is brought into contact with water it gradually dissolves up
to a certain point, beyond which no further increase in concentra-
tion of the solution takes place. A solution is said to be saturated
when it contains as much dissolved substance (solute) as possible,
and remains in equilibrium with some of the solid substance; i.e.
no spontaneous change takes place under constant conditions in the
concentration of the solution or the amount of substance remaining
undissolved. A solution is unsaturated when it contains less of
the solute than a saturated solution under the same conditions;
that is to say, when the act of solution is incomplete because either
sufficient of the substance is unavailable, or there has not been
sufficient opportunity for the state of saturation to be reached.
Consequently an unsaturated solution will always dissolve more
of the substance when opportunity occurs.

In contrast to the unsaturated state is the supersaturated state.
A solution is supersaturated when it contains more of the solute
than a saturated solution in contact with the solid could contain
under the same conditions. A supersaturated solution can exist
only in the absence of solid matter; it is in a meta-stable state
which is disturbed by a slight impetus from without. Its produc-
tion depends upon the fact that most substances are more soluble in
hot water than cold, and that if a hot, saturated solution, free from

202 CHEMICAL THEORY

solid, is cooled, the separation of crystals may be postponed until a
fragment of crystal enters the solution. Then rapid separation of
crystals takes place until equilibrium in contact with the solid is
secured, that is, the solution becomes saturated. The phenomenon
illustrates the inertia of matter, that is, the tendency of a substance
or mixture to persist in the same physical state, although the con-
ditions for the stable existence of that state have ceased, until an
external stimulus causes a rapid adjustment to the conditions of
true equilibrium. The supercooling of a liquid, i.e. its persistence
in the liquid state at a temperature below that at which solidifica-
tion normally occurs, is an analogous phenomenon.

Supersaturation is well illustrated by sodium sulphate solution.
If a hot, saturated solution of this salt is prepared and filtered
into a flask, the neck of which is then plugged with cotton-wool,
the solution may be cooled to atmospheric temperature without
separation of solid. If the cotton-wool is then removed and a
fragment of the crystallized salt dropped into the flask, rapid
crystallization takes place with evolution of heat, until a solid
cake of crystals is produced.

Whether a solution of a solid substance is unsaturated, saturated,
or supersaturated at a given temperature is properly determined
with reference to the substance, as follows: If the solid is added
to an unsaturated solution, solid will dissolve, and the solution
become more concentrated; if it is added to a saturated solution
the solution will remain unaltered; if it is added to a supersaturated
solution solid will separate, and the solution become more dilute.
All three solutions will attain to the same concentration, because
they attain to a state of true equilibrium in contact with the solid.
Moreover, the state of equilibrium is dynamic rather than static.
The concentration of the solution remains constant, not because
nothing is taking place, but because molecules of the solute are
entering and leaving the solution at the same rate. The three
states may be represented graphically as follows:

solute

Unsaturation : solid * solution.

solute

Supersaturation: solid - solution.

solute
Saturation: solid -^-^ solution.

The state of equilibrium between a solid and a solution is
thus analogous to the state of chemical equilibrium between the

SOLUTIONS 203

factors and products of a reversible chemical reaction, and to the
state of thermal equilibrium between two bodies at the same
temperature.

The Process of Solution.

A readily soluble solid is easily dissolved by being stirred or
shaken with the solvent. A saturated solution may be prepared
either by keeping the solid in contact with the unsaturated solvent
until saturation is attained, or, more rapidly, by heating the solvent
with excess of the solid, and then cooling the solution to the desired
temperature in contact with the solid. The more finely divided a
solid is the more rapidly will it dissolve, on account of the larger
surface exposed. It is therefore wise to pulverize a solid before
attempting to dissolve it, unless it is quite easily soluble.

The physical state of a solid, apart from mere fineness of sub-
division, also determines its solubility. Thus yellow mercuric oxide
is about 14 per cent more soluble in water, and at the same time
more active chemically than the red, crystalline form of this com-
pound; this difference is probably due to an allotropic difference
in the two forms. Again, precipitated calcium carbonate is more
soluble in ammonium salt solutions than the crystalline form
produced by heating the precipitate with the liquid. Thus, when
calcium carbonate is precipitated by ammonium carbonate added
to calcium chloride solution in presence of ammonium chloride,
the precipitate is at first flocculent and amorphous, but becomes
crystalline and less bulky when the liquid containing it is boiled;
and, further, boiling serves to precipitate in a crystalline form
traces of calcium carbonate that remain in solution in the cold
liquid. These phenomena illustrate the influence of physical state
on solubility.

Influence of Temperature on Solubility.

Rise of temperature generally, but not always, increases the
solubility of a solid in water.

The solubility of potassium chlorate, for example, is much in-
creased by rise of temperature, that of potassium chloride shows a
less increase, whilst that of sodium chloride is little affected by
change of temperature. On the other hand the solubility of calcium
hydroxide, as well as that of certain organic salts of calcium such as
the acetate, decreases with rise of temperature, so that from a cold

204

CHEMICAL THEORY

saturated solution solid separates on warming; whilst the solubility
of gypsum, CaSO 4 -2H 2 O, reaches a maximum at about 40 C. and
then diminishes again.

TABLE OF SOLUBILITIES

Tempei a-
ture.

KC10 3 .

KC1.

Nad.

Ca(OH) 2 .

Tempera-
ture.

CaS<V
2 H 2 0.

10 J

4-40

31-0

35-7

0-176

0-2439

7-22

34-0

35-8

0-165

0-2550

9-26

37-0

36-0

0-153

0-2631

13-31

40-0

30 -3

0-141

0-2644

17-95

42-0

36-7

0-128

0-2651

23-42

45-5

37-1

0-116

0-2653

29-16

48-3

37-5

0-106

0-2541

36-93

51-1

38-0

0-094

65-3

0-2444

46-11

54-0

38-5

0-085

0-2337

100

55-54

56-7

39-1

0-077

100

0-2048

These results are shown more clearly when expressed as curves
in which solubilities are ordinates and temperatures are abscissae
(fig. 42.) The values for Ca(OH) 2 and CaSO 4 2H 2 O are multiplied
by 100 to bring them into the figure according to scale.

10" 20 30 40 50 60 70 80 90 100

Temperature.

Fig. 42

The solubilities of crystallized sodium sulphate appear to be
anomalous. They are as follow:

SOLUTIONS

SOLUBILITIES OF SODIUM SULPHATE

205

Temperature.

Solubility.

Temperature.

Solubility.

9-0

50-6

13-4

48-8

19-4

45-3

28-0

43-7

40-8

100

42-5

32-75

50-65

The curve (fig. 42) shows a sharp break at 32 5, and indeed it
is clear that the solubility is represented by two curves which inter-
sect at that point. These curves must correspond to two distinct
states, either solid or in solution, the point of intersection indicating
an abrupt change from one state to the other. It must be remembered
that the numerical values represented by the curve are the amounts
of the anhydrous salt found in the saturated solutions at the corre-
sponding temperatures. Why does the power of water to dissolve
sodium sulphate so suddenly change at the indicated temperature?
Inquiry shows that the nature of the solution does not change; it is
the nature of the solid in contact with the solution that changes.
The fact is that the decahydrate Na 2 SO 4 10H 2 exists in contact
with the solution only up to 32-5; above that temperature this
hydrate loses all its water, so that it is the anhydrous salt that
remains in contact with the solution. Thus the first part of the
curve represents the solubility of Na 2 SO 4 10H 2 O in terms of
Na 2 SO 4 , the second part that of anhydrous Na 2 SO 4 . The change
from the crystallohydrate to the anhydrous salt at 32 5 C. may
be seen if the solid hydrate is gently heated. It will melt at this
temperature, which is the transition temperature, and give rise to
a meal of anhydrous crystals in contact with a saturated solution
of the salt in the available water derived from the hydrated crystals.
It follows from this that if a solution of sodium sulphate is made
to crystallize above 32 5 C. the anhydrous salt will separate, whilst
below this temperature the decahydrate will be formed.

Sodium carbonate shows analogous phenomena, but in this case
it is Na 2 OO 3 H 2 O which is formed from Na 2 C0 3 '10H 2 O, not the
anhydrous salt.

Relation between Chemical Composition and Solubility.

The relationship between chemical composition and solubility is
a large question of much interest and importance. The operations
of qualitative analysis familiarize the student with parts of the

206 CHEMICAL THEORY

subject, and a thoughtful study of a table of solubilities leads to
certain generalizations.

For example, all nitrates and chlorates, and nearly all chlorides
and sulphates, are readily soluble in water; all carbonates and
phosphates, except those of the alkalis, are practically insoluble;
moreover, the solubilities of sulphides in water and dilute acids and
alkalis are important for analytical separations.

The solubilities of metallic compounds in relation to the periodic
system are significant. Thus it is only the most powerfully basic
hydroxides that dissolve in water; and the alkaline earth hydroxides
which follow those of the alkalis in basic strength increase in solu-
bility with increase of basic strength from calcium to barium. Thus
100 parts of water at 10 C. dissolve

Ca(OH) 2 Sr(OH) 2 Ba(OH) 2

0-176 0-566 2-48

Gradations of solubility of related compounds in a natural group
are shown particularly in the case of almost insoluble salts. The
sulphates of the alkali metals, the sulphides of zinc and cadmium,
which differ sufficiently in solubility in dilute acid to be precipitated
in different analytic groups, and the chloride, bromide, and iodide of
silver furnish examples.

SOLUBILITIES OF SLIGHTLY-SOLUBLE SALTS
Grams per litre :

Atl8C. CaSO 4 2-016 SrSO 4 0-114 BaSO 4 0-0023

At20C. AgCl 0-0016 AgBr 0-000084 Agl 0-0000028

It is interesting to note that on account of solubility differ-
ences bromide solution converts precipitated AgCl into AgBr, and
iodide solution AgCl and AgBr into Agl. A similar phenomenon
is observed in the titration of neutral chloride solution by standard
silver nitrate in presence of chromate. Owing to the greater solu-
bility of chromate there is no permanent precipitation of silver
chromate until all the chloride in solution has been precipitated.
The relationships of solubility between simple and double or com-
plex salts are worthy of notice. Compare, for example, the alkali
sulphates and the alums:

K. Kb. Cs.

Sulphates 9-22 42-6 173-1

Alums 7-60 1-81 0-49

grams per 100 grm. water at 10 C.;

SOLUTIONS 207

as well as the simple chlorides and platinichlorides:

K. Rb. Ca.

Chloride 31-0 84-4 174-7

Platinichloride 0-90 0-154 0-050

grains per 100 grm. water at 10 C.

These figures strikingly illustrate the fact that the solubilities of
double and complex salts of the alkali metals stand in a reverse
relation to those of the corresponding simple salts. It will be
remembered, moreover, that sodium alum and sodium platinichloride
are freely soluble in water.

SUMMARY

SOLUBILITY OF A GAS. The solubility, or coefficient of solubility,
of a gas in a liquid is defined as the volume of the gas which is dis-
solved by unit volume of the liquid.

LAW OF HENRY. The quantity of a gas dissolved by a given
volume of a liquid is proportional to the pressure of the gas. Or
(from Boyle's law) the volume of a gas dissolved by a given volume
of a liquid is independent of the pressure.

LAW OF DALTON AND HENRY. The quantities of the various
gases dissolved from a gaseous mixture by a given volume of liquid
are proportional to the solubilities and partial pressures of the
gases.

SOLUBILITY OF A SOLID SUBSTANCE. The solubility of a solid
substance in water is defined as the number of grams of the
substance which will dissolve in 100 grm. of water.

CHAPTER XI
THE PROPERTIES OF DILUTE SOLUTIONS

The properties of solutions, and particularly of dilute solutions,
have been studied in great detail during recent years, and the facts
and theories connected with this subject have aroused keen interest.

An aqueous solution of silver nitrate yields a precipitate of
silver chloride with a similar solution of hydrogen chloride, or any
metallic chloride, but not with an aqueous solution of chloroform,
CHC1 3 , in spite of the large proportion of chlorine this compound
contains. What is the reason for this difference? A plain answer,
correct, so far as it goes, would be that silver nitrate solution is a
reagent, not for chlorine, but for a chloride, and chloroform is not
a chloride. A further difference, however, between solutions of a
chloride and of chloroform relates to electrolysis, that is, the decom-
position of a substance in aqueous solution by means of a current
of electricity. A chloride solution is capable of electrolysis, a
solution of chloroform is not. Consequently the former is an
electrolyte, the latter not. Since the properties of dilute solutions
are elucidated by the phenomena of electrolysis, that subject will
first receive attention. 1

Electrolysis.

The beaker shown in fig. 43 contains a solution of copper
sulphate into which dip two pieces of platinum foil. These are
connected by lengths of copper wire preferably silk-coated, except
at the ends, for the purpose of insulation with an electric battery
of, say, three cells, shown at the bottom of the figure. An electric
current flows from the positive pole of the battery through the
wire to one platinum plate, then through the solution and back
to the battery through the other platinum plate and wire; so the
circuit is completed. The platinum plates are called the electrodes,

1 The contents of this chapter were partly anticipated when the modern view of the
molecule was discussed. Nevertheless the chapter stands in proper sequence here.

208

THE PKOPERTIES OF DILUTE SOLUTIONS

206

Fig. 43

i.e. the ways of the current, the entrance to and exit from the

solution undergoing electrolysis; they are named, rather pic-

turesquely, the anode and the cathode, i.e. the way up and the way

down; the positive (+) electrode is the

anode, the negative ( ) electrode the

cathode. The electric current, generated

in the battery hy chemical action, travels

in two modes throughout this circuit

through the wires without chemical

change, through the solution with chem-

ical change. For electrolysis takes place

when a current passes through a solution

of copper sulphate, which is an electro-

lyte; metallic copper is deposited on the

cathode, which thus becomes electro-

plated, and a quantity of oxygen chemi-

cally equivalent to the copper is evolved from the anode, whilst

round about the anode the solution develops acidity, such as would

be derived from the reaction:

S0 4 + H 2 = H 2 S0 4 + 0.

It is believed, therefore, that atoms of copper charged with
4- electricity travel to the cathode, and that there they give up
their charges, which pass back through the wire to the battery,
whilst the copper remains behind on the cathode; and that similarly
SO 4 radicles, with corresponding charges of negative electricity,
travel to the anode, there causing oxygen to be evolved from
water in the manner shown above. The charged Cu atoms and
SO 4 radicles are called ions, because they go to the electrodes;
the former are cations, because they travel to the cathode, the
latter anions, because they travel to the anode. The S0 4 ion is
also sometimes called the sulphion.

If the anode is of copper instead of platinum, there is no
evolution of oxygen gas, but copper is dissolved by the influence
of S0 4 , so that the concentration of the solution remains unchanged.
while copper is in effect transferred from the anode to the cathode.
This kind of action takes place in silver-plating. The article to
be plated is made the cathode and a bar of pure silver the anode,
the electrolyte being a solution of potassium silver cyanide,
KAg(CN) 2 . Metallic silver then dissolves from the anodic bar at

) GO;

210 CHEMICAL THEORY

the same rate as it is deposited on the article forming the cathode.
If dilute sulphuric acid is electrolyzed, hydrogen appears at the
cathode, and an equivalent quantity of oxygen at the anode, whilst
the acidic strength of the solution remains unchanged. The action
may be represented thus:

Cathode. Anode.

2H,

2H 2 !SO 4 2SO 4
: + 2H 2 O

= 2H 2 SO 4

'2-

Since the water alone undergoes permanent change, being de-
composed into its constituent gases, 2 volumes of hydrogen and 1
volume of oxygen, it is customary to call this process the electro-
lysis of water, the acid being added to enable the water to conduct
the current. If the current is led into pure water no action is
observed to take place, for the water is practically a non-conductor,
a non-electrolyte.

Why diluted sulphuric acid should be an electrolyte, and water
not, is a question which may well be asked a question to which
an answer must be found.

One further example of electrolysis may, however, first be
noticed. The electrolysis of sodium sulphate gives sodium ions
at the cathode and S0 4 ions at the anode. The S0 4 , reacting with
water, produces sulphuric acid and oxygen as before, but the sodium
at the cathode causes hydrogen to be evolved from the water and
sodium hydroxide to be formed in solution, thus:

Cathode Anode

2 Na 2 :S0 4

4 Na 2 S0 4

+ 4 H 2 -f 2 H 2 O

= 4 NaOH -f =2 H 2 SO 4 +

So in this case acid and alkali appear in equivalent quantities
at the anode and cathode respectively, together with oxygen and
hydrogen gases respectively.

If wires from the terminals of a battery are laid side by side
on a piece of porous paper soaked in a solution of sodium sulphate,
with phenol-phthalein as an indicator, a crimson stain will appear
beneath the wire attached to the negative pole of the battery.
The cathode is thus indicated through the development of alkali
by electrolysis. Such paper is called pole-finding paper.

THE PROPERTIES OF DILUTE SOLUTIONS 211

Electrolytic Dissociation.

There now remains the question as to the nature of electrolysis
and the constitution of electrolytes; why, for example, hydrogen
chloride solution is an electrolyte and not chloroform, and why
pure water is not an electrolyte.

And it is a gain at once to recognize that it is only substances
reactive in aqueous solution, which are capable' of reactions of
neutralization or double decomposition, that are electrolytes; in
a word, it is acids, bases, and salts that belong to this category.

In connection with electrolysis there are three points to be
considered :

i. The separation of the electrolyte into its constituent ions,
ii. The migration of these ions to the electrodes.

iii. The separation of the electrolytic products at the electrodes.

Now ii and iii are plainly connected with the passage of the
current; the ions are conveyed by the current to the electrodes,
and there appear as or give rise to the products of electrolysis.
But upon the question as to whether the current separates the
electrolyte into its ions before directing them to the electrodes
further light is needed.

A priori, it would appear that when a substance like sodium
chloride dissolves in water its molecules dissolve entire. The
alternative is to suppose that they decompose or dissociate into
sodium and chlorine; and this seems absurd to anyone acquainted
with the properties of these two elements. But then the question
recurs: What is the difference between the chlorine in sodium
chloride and in chloroform, which accounts for the different be-
haviour of these substances towards silver nitrate? or why are
the component parts of electrolytes also reactive in other ways,
whilst the components of non- electrolytes are inert? Evidence
apart altogether from electrolysis can be found in answer to this
question.

The student will remember that the molecular weights of dis-
solved substances can be determined by cryoscopic and ebulliscopic
methods; i.e. by observing the depression of freezing-point or rise
of boiling-point of a solvent caused by a known weight of dissolved
substance. He may also have noticed that the substances generally
chosen to illustrate these methods are non-electrolytes; sugar, for
example, in aqueous solution, or some organic substance in an

212 CHEMICAL THEORY

organic solvent. The reason for this choice is that electrolytes
give anomalous results which are best avoided in an elementary
consideration of the subject. These anomalous results, however,
throw much light on the matter now in hand. A given weight
of sodium chloride in dilute solution, for example, causes practi-
cally twice the depression of freezing-point or rise of boiling-point
caused by an equimolecular proportion of cane-sugar. Therefore
there are twice as many molecules, or of what behave as molecules,
in a solution of sodium chloride as in a solution of sugar calculated
to be equimolecular. This fact supports the apparently impossible
conclusion as to the dissociation of sodium chloride molecules into
sodium and chlorine atoms; and further support for this conclusion
is derived from the phenomena of osmotic pressure.

The osmotic pressure of a substance in solution is the pressure
which causes it to diffuse into the pure solvent. It is an expression
of the motions of the molecules of the solute, and is measured by
means of a " semi-permeable membrane " through which the solvent
can pass but not the solute. 1 Now, there is a close analogy between
osmotic pressure and gas pressure. The osmotic pressure of a dis-
solved substance varies directly as its concentration, or inversely as
its volume (cf. Boyle's law); it also varies directly as the absolute
temperature (cf. Charles's law); and the numerical value of the con-
stant R in the equation pv = RT is the same whether p stands
for gaseous pressure or osmotic pressure. Further, equimolecular
solutions of different substances have the same osmotic pressure
(cf. Avogadro's theory).

Now, when a gaseous substance, such as N 2 4> undergoes dis-
sociation without increase of volume its pressure increases, and
is doubled when every molecule has produced two molecules.
Similarly, the osmotic pressure of dilute sodium chloride solution is
double what might have been expected when it is compared with
that of a non-electrolyte such as sugar. So, again, there is direct
evidence of the dissociation in aqueous solution of sodium chloride
molecules into the only possible parts, sodium and chlorine.

The student is now prepared for the theory of electrolytic dis-
sociation or ionization. This theory, originated by the Swedish
physicist Arrhenius, in 1887, has revolutionized modern thinking
upon a vast number of chemical phenomena. It asserts that when

1 Such a membrane is produced by precipitating cupric f errocyanide within the walls of a
porous cell.

THE PKOPERTIES OF DILUTE SOLUTIONS 213

an electrolyte, such as sodium chloride, is dissolved in a suitable
solvent, such as water, it undergoes spontaneous dissociation into its
ions, sodium and chlorine, but that these ions have upon them charges
of electricity; the sodium ion, for example, consists of an atom of
sodium with a certain positive charge, and the chlorine ion, or more
properly chloride ion, of an atom of chlorine with an equal negative
charge. It is these charges which differentiate the ions from the
free elements. It is free sodium, for example, which reacts with
water, not sodium ions; free chlorine which has a characteristic
colour and smell, not chloride ions. So when sodium chloride
dissolves in water it undergoes the following reversible change:

NaCl ^^ Na + Cl or NaCl ^=^ Na- + Cl',

these being alternative methods of representing electrolytic dissocia-
tion, or, more briefly, ionization. 1

Only, therefore, as a substance is ionized can it be reactive; only
then is it capable of electrolysis or chemical reactions in solution.

When hydrogen chloride is dissolved in water it is ionized thus:

HC1 ^ H-

and sodium hydroxide thus:

NaOH ^=^ Na- + OH'.

Water, however, which is not an electrolyte, is ionized thus:
H 2 O ^= H- + OH',

to the minutest extent; only, indeed, to the extent of 1 grm.-niole-
cule in 10 million litres. Since, therefore, H* and OH' ions cannot
coexist in the same solution, they will combine when solutions which
separately contain these ions are mixed. Now a solution which
contains H" ions contains an acid; indeed, according to the theory of
ionization, degree of acidity depends essentially upon the concentra-
tion of hydrogen ions; 2 whilst a solution containing OH 7 ions is a
solution of a base, that is of an alkali, the strength of which depends
similarly on the concentration of OH' ions. When, therefore, solu-
tions of an acid and an alkali are brought together, e.g. solutions of
HC1 and NaOH, a reaction such as the following takes place:
H- + Cl' + Na- + OH' = HOH + Na- + Cl'.

1 If the student has mastered Chapter VI he will know that, according to modern views,
sodium chloride becomes ionized as it is formed, so that solution in water but causes dis-
sociation of already existing ions.

2 See note in Appendix on Hydrion Concentration and pHL value.

214 CHEMICAL THEORY

The process of neutralization is the reaction:
H- + OH' = HOH,

and there is no other reaction. Neutralization is but the synthesis
of water molecules from their constituent ions; the sodium and
chlorine in the above case are no more united after neutralization
than before, for they remain as separate ions throughout the process.

This theory of the process of neutralization, now nearly a genera-
tion old, was revolutionary enough at the time of its inception, but
it has now become an integral part of physico-chemical theory;
it is a fundamental idea of modern chemistry which has stimulated
innumerable researches and greatly extended the boundaries of
scientific knowledge.

The first point in connection with electrolysis, cited on p. 211
the separation of the electrolyte into its constituent ions has now
been duly considered; and the conclusion is that this separation is
a spontaneous process accompanying the act of solution; it is neither
a part of the process of electrolysis, nor in any way dependent upon
the introduction of an electric current into the solution; but it is a
necessary precursor alike of electrolysis and every kind of chemical
activity of which the substance in solution is capable. The student
must therefore most carefully avoid the fallacy of supposing that
electrolytic dissociation is a part of electrolysis.

Nevertheless, since electrolytic dissociation involves the separa-
tion of the ions with electric charges upon them, it is seen how
closely chemical and electrical phenomena are connected.

With these ideas in mind the process of electrolysis may again
be reviewed.

Ions of Na" and 01', for example, stand free in every part of a
solution of NaCl, and are directed by the current to the cathode and
anode respectively. The electric charges of these ions are equal and
opposite, and their magnitude is connected with valency. Thus Na
and OH, being univalent, carry each a unit charge; Cu and S0 4 ,
however, both being bivalent, carry double charges. If the same
current passes successively through two electrolytic cells containing
solutions of NaOH and CuS0 4 respectively, the number of bivalent
Cu ions, carrying double charges, conveyed to the cathode will be half
the number of univalent Na ions which are similarly conveyed, or
in other words sodium and copper will be conveyed in equivalent
quantities by the current in its course.

THE PROPERTIES OF DILUTE SOLUTIONS 215

These conclusions are expressed in Faraday's laws of electrolysis,
discovered experimentally long before electrolytic dissociation was
thought of.

LAWS OF ELECTROLYSIS. I. The amount of chemical action in
an electrolytic cell is proportional to the current that passes.

II. The quantities of substances liberated at the electrodes when
the same current passes successively through different electrolytic
solutions are chemically equivalent.

From the second law it is apparent that chemical equivalents
may be found by means of electrolysis. For example, the same
current may be passed successively through solutions of sulphuric
acid and copper sulphate; the evolved hydrogen may be measured
and the deposited copper weighed. The weights of hydrogen and
copper will be found to be in the ratio of their equivalent weights.
The electro- chemical equivalent of copper, for example, is of course
simply its chemical equivalent determined by electrolysis.

It is possible now to fill a little more detail into the mental
picture of electrolysis of, say, sodium chloride. Na* and Cl' ions
react at the electrodes in equal numbers, but they need not have
travelled there at the same rate. Suppose a number of people
within an enclosure seek to leave it at two opposite exits. They
must pass through the exits at the same rate, let it be supposed, but
that does not influence the rate at which each individual may
approach the exit that he chooses. The quicker he reaches the exit
the longer must he wait there; that is all. In other words, unequal
approach to the exits will involve unequal crowding around them.
That is what happens in electrolysis. For example, in the electrolysis
of sodium chloride solution the chloride ions travel faster than the
sodium ions, with consequent increase in concentration of the
chloride ions in the neighbourhood of the anode.

A simple model, roughly illustrating electrolysis, may be made
in the following way. Two equal strips of plain cardboard, say
8 in. long and J in. wide, are taken, and a similar scale, say with
-in. divisions, is marked out on each. A plus sign is then made
in the centre of each division on one card, and a minus sign on
the other. The two cards are then held together by rubber bands,
with the scales face to face. This represents the state of a solution
previous to electrolysis. The process of electrolysis is shown
by moving the cards relatively to each other. An equal number
of plus and minus signs must necessarily be exposed at either end,

216

CHEMICAL THEORY

independently of the question whether the cards are moved at
equal or unequal rates, or whether one card even remains stationary.
It is the relative motion of the cards which determines the ex-
posure of the signs; similarly it is the relative motion of the ions
which determines the extent of electrolysis. The process of elec-
trolysis is illustrated similarly in the following diagram:

(2)

(4)

Fig. 44

In (1) the state before electrolysis is represented; (2) represents
electrolysis with ions migrating at equal rates; (3) shows the
cations migrating more slowly than the anions; (4) shows the
cations stationary whilst the anions move at the same rate as
before. The extent of electrolysis is influenced by the speed of
each ion, but the separated products are always equivalent. The
superior speed of the anions in (3) and (4) results in the solution
becoming more concentrated towards the anode.

What actually happens at the electrodes, however, requires
elucidation. When the electrolytic product is identical in com-
position with the ion there is no question; for example, in the
electrolysis of copper sulphate the cupric ion loses its charge at
the cathode, becomes an atom of copper, and is deposited as metal.
When, however, secondary products appear, as when, for example,
hydrogen and oxygen gases are obtained by the electrolysis of
dilute caustic soda solution, the matter is not so simple. What
then is the mechanism by which sodium ions at the cathode give
rise to hydrogen gas? The older explanation was that the ion
lost its charge and became an atom of sodium, which then reacted
with the water, evolving hydrogen, and thereby returned to its
former state. Why, however, should such a cycle of change be
assumed, why should sodium leave the solution only to re-enter it?

THE PROPERTIES OF DILUTE SOLUTIONS 217

There is certainly no evidence of the appearance of metallic sodium.

Here must be introduced the idea of solution pressure. When
zinc dissolves in dilute sulphuric acid, evolving hydrogen, it is
because it has a solution pressure greater than hydrogen, and by
reason of this pressure displaces hydrogen. This amounts to the
same thing as saying that zinc is more electro-positive than
hydrogen. Similarly, when sodium dissolves in water, displacing
hydrogen, it does so by its superior solution pressure.

Now, since hydrogen appears in the electrolysis of sodium
hydroxide solution it is displaced from water in preference to
sodium, which will never appear in the elementary state so long
as the less electro-positive hydrogen can take its place. And the
hydrogen that comes from water, whether by the action of metallic
sodium or during electrolysis, is the hydrogen formed by the
ionization of water thus:

H 2 ^ H' + OH',

although this ionization is so very slight; for molecular water,
according to the conception of the theory, would be inert.

The following, then, is what happens at the cathode during
the electrolysis of dilute sodium hydroxide solution. Sodium ions
arrive there, each with its positive charge. They do not leave
the solution, but instead displace the weaker hydrogen from the
ionized water. So these hydrogen ions lose their charges, which
were equal to those of the sodium ions that have displaced them,
and become molecules of gas, while the current flows on through
the cathode. More water is ionized in consequence of the dis-
turbance of equilibrium, its hydrogen being in turn displaced by
the arriving sodium ions, so that Na* and OH' may be present
at the cathode in equivalent amounts.

Meanwhile OH' ions arriving and giving up their charges at
the anode react with one another thus:

4 OH' = 2H 2 O + 2 .

This is because they cannot remain in the solution unbalanced
by cations. Consequently whilst hydrogen and oxygen appear
in equivalent quantities at the electrodes, the solution remains
unchanged in composition, the only effect being the displacement
of the alkali (NaOH) towards the cathode.

It may be noted that in the electrolysis of nearly anhydrous

218 CHEMICAL THEORY

fused sodium hydroxide, as in the original experiment of Davy,
metallic sodium appears instead of hydrogen at the cathode; this
is because there is no water whose hydrogen can be displaced by
the accumulating sodium ions.

The student should now be able to express the facts of electro-
lysis of other solutions that come under his notice in terms of the
ionization theory, r

Chemical Reactions in Solution.

The properties of acids, bases, and salts in dilute aqueous solution
are the properties of their ions, and reactions between them are
reactions between their ions.

This applies, for example, to the Colours of Salt Solutions.
Thus, the yellow colour of potassium chromate solution is the
colour of the chromate ion, the blue colour of copper sulphate
the colour of the cupric ion, and so forth. Certain ions are
essentially coloured, others essentially colourless.

The following basic and acidic ions are coloured:

BASIC IONS

Cations.
Cu," blue.
Co", pink.
Ni", bright green.
Cr'", purple.
Mn", pale pink.
Fe", green.

ACIDIC IONS

Anions.

CrO 4 ", bright yellow.
Cr 2 O 7 ", orange red.
MnO 4 ", deep green.
MnO 4 ', crimson.
[Fe(CN) fl ]"", yellow.
[Fe(CN) 6 ]'", yellow.

Since cupric ions, for example, are blue, and cobaltous ions
pink, dilute solutions of cupric and cobaltous salts are blue and
pink respectively, because their salts are ionized. Probably, how-
ever, these ions are hydrated.

Anhydrous eupric chloride is brown, and if dissolved in water
without ionization would probably give a yellow solution. A
concentrated aqueous solution of this salt is in fact green, but
becomes blue on dilution. The green colour is believed to be due
to admixture of yellow, non-ionized cupric chloride with blue cupric
ions. Since dilution promotes ionization it changes the yellow
colour of the non-ionized salt into the blue colour of cupric ions;
consequently the green colour of the mixture of CuCl 2 and Cu"
gradually gives place to the blue colour of the latter.

Again, a pink solution of cobaltous chloride of moderate con-

THE PROPERTIES OF DILUTE SOLUTIONS 219

centration is turned blue by the addition of concentrated hydro-
chloric acid, and also by heating.

In this case the blue colour is probably due not to the presence
of non-ionized Co01 2 , but to the formation of HCoCl 3 or H 2 CoCl 4 ,
the anions of which [CoCl 3 ]' or [CoClJ" are blue.

If ammonia is added to a solution of a cupric salt there is pro-
duced first a bluish-green precipitate, and then a deep-blue solution.
The colour of this solution is too intense for that of cupric ions;
from the solution, indeed, the compound CuSO 4 -4NH 3 H 2 O,
cuprammonium sulphate, will crystallize after the addition of
alcohol. The deep-blue colour has been shown to be due to the
cuprammonium cation [Cu(NH 3 )J".

Now, if potassium cyanide is added to -the deep-blue cupram-
monium solution the colour fades and the solution eventually
becomes colourless. This solution still contains copper, but plainly
the metal is now present neither as cupric nor cuprammonium
ions; for apart from its colour it fails to react with hydrogen
sulphide. From this colourless solution potassium cuprocyanide,
K 3 Cu(CN) 4 , can be isolated as a colourless salt; thus the copper
is present in solution as the colourless anion [Cu(CN) 4 ] /// .

Similarly, if potassium cyanide is added to cobaltous chloride
solution cobaltous cyanide is precipitated, and then dissolves in
excess of the cyanide to produce potassium cobaltocyanide,
K 4 Co(CN) 6 , which readily passes by oxidation into the cobalti-
cyanide, K 3 Co(CN) 6 . In this case the anions [Co(CN) 6 ]"" and
[Co(CN) 6 ]'" are formed in turn; these are pale yellow in colour,
and differ entirely in chemical reactions from cobaltous ions.
Thus the formation of cobalticyanide prevents the precipitation
of hydrated, black, cobaltic oxide by means of alkali and bromine
water, the corresponding oxide of nickel being precipitated by these
reagents in spite of the presence of cyanide, since it forms no stable
anions like cobalt.

The more familiar potassium ferro- and ferri-cyanide furnish
examples of anions containing iron, with reactions entirely different
from those of either ferrous or ferric salts.

These illustrations suffice to show the importance of the con-
ception of ions in analytical chemistry. The phenomena of pre-
cipitation and of hydrolysis which play so important a part in
analytical work will be dealt with later. There remains, however,
to be considered in an elementary way:

220 CHEMICAL THEOKY

The Theory of Indicators.

The indicators used in acidimetry and alkalimetry are natural
or synthetic colours which undergo a sharp colour change in a
liquid which changes from acid to alkaline or vice versa.

It has been stated above that an acid is a compound which
separates hydrogen ions in solution, and that acidic intensity
depends upon the concentration of these ions. Thus, different acids
differ in acidic strength because they are ionized to different extents
in equivalent aqueous solutions. Hydrochloric and nitric acids are
acids of maximum strength, being completely ionized in dilute
solution; sulphuric acid is about half, oxalic acid a fifth, as strong
as hydrochloric and nitric acids; whilst acetic acid has a strength
less than 1 per cent that of the strongest acids. A possible mis-
understanding may here be guarded against. Although oxalic acid,
for example, is only partially ionized in, say, decinormal solution,
it does not follow that it cannot be completely neutralized, i.e. con-
verted into sodium oxalate by titration with decinormal soda.
Experience shows that it can. The partial ionization is a state
of equilibrium reached when the acid is dissolved alone in water.
That state is disturbed when alkali is added and further ionization
takes place, so that the whole of the acid eventually provides
hydrogen ions to react with the hydroxide ions of the alkali and
complete neutralization.

Indicators are feeble acids or bases, the salts of which differ in
colour from the free acids or bases themselves. The common indica-
tors litmus, methyl orange, phenol-phthalein are acidic substances.

The acid of litmus is red, the alkali salt blue; the acid of methyl
orange is pink, the alkali salt yellow; the acid of phenol-phthalein
is colourless, the alkali salt crimson. Now the alkali salt of a
feeble, scarcely ionized acid undergoes much more ionization than
the acid itself, and so the above facts may be expressed thus: The
undissociated litmus acid is red, its anion is blue; the undissociated
methyl orange acid is pink, its anion yellow; the undissociated
phenol-phthalein acid is colourless; its anion crimson. It seems
extraordinary, however, that ionization should produce such re-
markable colour changes, and it is generally believed that these
changes depend upon intra-molecular rearrangement, i.e. the shuf-
fling of the atoms within the complicated molecules of the indi-
cators during neutralization.

THE PKOPERTIES OF DILUTE SOLUTIONS 221

Now it is well known that the three above-named indicators
differ in their behaviour towards certain acids and bases. The
facts are these:

Litmus is turned bluish-red by carbonic acid, but becomes blue
again on boiling the solution containing it, on account of the dissi-
pation of the carbonic acid. Litmus may be used as an indicator
for oxalic acid as well as mineral acids, and for ammonia.

Methyl orange is practically unaffected by carbonic and boric
acids, so that solutions of sodium carbonate and borax may be
titrated by mineral acid with the aid of methyl orange, as if they
were sodium hydroxide. Oxalic acid, however, produces a gradual
instead of a sharp colour change with methyl orange. This indi-
cator is therefore useless for titrations with oxalic acid.

Phenol-phthalein is sensitive to the weakest acids, even car-
bonic acid; it is indeed more sensitive than litmus. This indicator
gives no sharp end point, however, with ammonia, and is therefore
useless in titrating this alkali.

The ionic theory explains these facts satisfactorily as follows:
The acids of the indicators stand in this order of diminishing
strength:

Methyl orange, litmus, phenol-phthalein.

The acid of methyl orange is sufficiently weak to be displaced
sharply from its alkali salt by mineral acids, but it is sufficiently
strong to compete with oxalic acid, and therefore to be displaced
but gradually by the latter; moreover, it is sufficiently strong to
form a salt with the feebly basic ammonia which is ionized, so as
to show sharp colour change.

The acid of litmus is weak enough to be displaced to some
extent by carbonic acid, 1 and completely by oxalic acid; but it is
also strong enough to form an ionizable salt with ammonia.

Phenol-phthalein is so weak an acid that it reacts sharply even
with carbonic acid; but, owing to its feebleness, it cannot form a
non-hydrolyzable salt with ammonia. Therefore its colour change
with ammonia is gradual, because the coloured ions derived from
the ionized ammonium salt do not at once appear, owing to the
hydrolysis of this salt.

1 It is said that the pure colouring-matter of litmus gives a full red colour with water
saturated with CO 2 .

22-2 CHEMICAL THEOKY

SUMMARY

ELECTROLYTE. An electrolyte is a substance which conducts an
electric current whilst undergoing chemical decomposition which is
electrolysis.

An electrolyte, when dissolved in a suitable solvent, undergoes
spontaneous dissociation into ions. This is electrolytic dissociation
or ionization.

LAWS OF ELECTROLYSIS OF FARADAY.

i. The amount of chemical action in an electrolytic cell is pro-
portional to the current that passes.

ii. The quantities of substances liberated at the electrodes when
the same current passes successively through different electrolytic
solutions are chemically equivalent.

CHAPTER XII
THE COLLOIDAL STATE

When a finely -divided solid is mixed with water or other
solvent, either it may dissolve completely or some or all of it
may remain undissolved. These two conditions are easily distin-
guished. If the solid dissolves completely the resulting liquid,
whether coloured or not, is clear or transparent. If the solid does
not dissolve completely, the liquid when shaken will appear turbid
or opaque, and if the mixture is allowed to stand undisturbed
the solid in suspension will in time settle, leaving the supernatant
liquid clear.

The distinction between a substance in solution and one in
suspension appears fundamental; for a liquid containing suspended
matter may be filtered to be made clear, but it is needless to filter a
solution. Yet suspended solids differ in the fineness of their sub-
division, and in the ease with which they are removed by subsidence
or filtration. Sand, for example, will settle more quickly than pre-
cipitated chalk, and both of these can be filtered more easily than
precipitated barium sulphate or calcium oxalate. And the student
is familiar with substances even more difficult to filter than the last
named. The sulphur which separates when hydrogen sulphide gas
is passed through an oxidizing solution, or when acid is added to a
polysulphide, cannot be removed completely by means of ordinary
filter-paper, the pores of which are evidently too large to retain the
minute particles of which the precipitate consists.

Indeed, although the solution may become practically trans-
parent, yet the presence of suspended sulphur is revealed by a slight
opalescence. Precipitated silver chloride presents a similar pheno-
menon; whilst copper sulphide, imperfectly precipitated from cold
solution, may yield a brown filtrate, which, although transparent,
contains the sulphide in a very fine state of subdivision. It appears,
therefore, that a 'liquid may contain suspended matter so finely

223

224 CHEMICAL THEORY

divided as not to produce opacity or to be removable by subsidence
or ordinary filtration. It is pertinent to ask, therefore, how such
a suspension differs from a true solution. Meanwhile the subject
may be approached from a different point of view.

When an aqueous solution of sodium silicate, or soluble glass, is
acidified with dilute hydrochloric acid, silicic acid is liberated, and,
if the solution is in a concentrated state, will separate from it in
the form of " gelatinous silica ". If, however, the solution is suffi-
ciently dilute, there is no precipitate, the liquid remaining clear. It
might be supposed that the difference between these two conditions
depended simply on the amount of water present, there being enough
water to hold the silicic acid in solution in the one case, but not in
the other. If this were so, the gelatinous silica separated in the
former case would be in equilibrium with a saturated solution of the
same substance. This, however, is not so; the phenomenon here
exhibited is indeed quite different from an ordinary case of precipi-
tation, as will appear in the sequel.

This and kindred phenomena were first investigated by Graham
in 1 849 in connection with experiments on liquid diffusion. Graham
found that the rates of diffusion into pure water of different sub-
stances in aqueous solution were various, and that simple salts and
acids passed rapidly through an animal membrane or parchment
paper, whilst complex substances like gelatine or glue in aqueous
solution did not penetrate these membranes. These latter substances
Graham called colloids (/coXXa, glue), whilst acids and salts, being
crystallizable or related to crystallizable substances, he called crystal-
loids. So crystalloids and colloids can be separated from one
another by aqueous diffusion through a parchment or other suitable
membrane, 1 fixed on a frame like a drum and floating on water,
crystalloids passing through the membrane into the external water,
whilst colloids remain behind in the drum. The process is called
dialysis, because it involves separation of one substance from
another by passing it through a membrane, through which crystal-
loids in solution, but not colloids, can pass.

A diluted acidified solution of sodium silicate may be submitted
to dialysis. The sodium chloride formed in the reaction,

Na 2 SiO 3 + 2HC1 = H 2 SiO 3 +2NaCl,

and excess of hydrochloric acid pass through the membrane of the
x Not a " semi-permeable membrane ", which allows only the solvent to pass.

THE COLLOIDAL STATE 225

dialyser, leaving the silicic acid behind in pure aqueous solution.
Such a solution may be concentrated by evaporation to a strength
of about 14 per cent. After this it changes to a jelly, similar to
that obtained by acidifying a concentrated solution of soluble glass.
Not only does concentration cause dialysed silicic acid to coagulate,
but a trace of hydrochloric acid, or some simple salt, acting
catalytically, produces the same effect. It thus appears that there
are two forms of aqueous colloidal silicic acid; the clear form, which
seems to be a solution, and the gelatinous form, which evidently
is not. These two forms, in which colloids in general may occur,
are called respectively the kydrosol and hydrogel, or simply sol
and c/el.

Besides silicic acid and organic substances such as gums and
resins, glue and gelatine, various inorganic substances occur or can
be obtained in the colloidal state.

Graham prepared the sols of ferric, chromic, and aluminium
hydroxides by dialysis ; the sols of arsenious and antimonious
sulphides may be prepared by boiling arsenious oxide and tartar
emetic respectively with water, and adding hydrogen sulphide to
their solutions: the liquids become yellow and red respectively
because of the formation of the sulphides in the sol condition. A
drop of hydrochloric acid added to either solution precipitates the
yellow arsenious or orange-red antimonious sulphide, the sulphides
thus assuming the gel condition.

It will now be understood that the above-mentioned brown
liquid obtained in precipitating copper sulphide contained the sol of
this sulphide, whilst the opalescent liquid containing sulphur held
this element also in the state of colloidal sol.

The sols of certain metals are interesting, and often display
remarkable colours. Thus gold and silver may be separated from
their salts by hydrazine, formaldehyde, &c. Faraday produced
blue, violet, and rose-coloured liquids by reducing gold chloride by
means of an ethereal solution of phosphorus floating on the surface
of its solution; and Bredig obtained sols of gold, silver, platinum,
&c., by an electric discharge through water between poles of the
metal.

It is easy to understand that colloids have been regarded by
chemists with much interest from the time of their discovery to the
present day. They are of practical importance because they embrace
many common noii-crytallizable organic substances such as gum,

(D60) 16

226 CHEMICAL THEORY

resin, glue, starch paste, egg-albumin, casein, and gelatine; but
they are particularly interesting from the physico-chemical stand-
point because they present a fresh phase of the great subject of
molecular physics. Indeed, a transparent dialysed liquid, consisting
of silicic acid and water for example, which on the addition of a
suitable catalyst becomes a jelly that can be inverted without flow-
ing, presents to the scientific mind a subject for investigation full of
an interest that can scarcely be surpassed.

What is this transparent liquid? Is it a solution like a mixture
of sodium chloride and water? If it is, why does the silicic acid
remain in the dialyser whilst the sodium chloride passes through it?
It has already been suggested that the process of dialysis is a kind
of filtration, i.e. that the silicic acid molecules are too large to pass
through the pores of the parchment paper. But filtration is applied
to something in suspension. Is the silicic acid in this apparently
clear liquid really in suspension?

An answer to this question as regards colloids in general has
been gained by the use of the ultra - microscope invented by
Siedentopf and Zsigmondy in 1903. The principle underlying
the use of this instrument is that illustrated by the vision of the
" mote in the sunbeam ". It is well known that the moving dust
of the air, which cannot ordinarily be seen, is made visible in a
beam of sunlight entering a darkened room through a chink
in a shutter. At the same time the track of the beam itself is
clearly outlined; but if the air is free from dust the sunbeam dis-
appears.

This effect, studied by Professor Tyndall, is applicable also to
liquids, and will reveal the presence of suspended particles within
them in the same way that it shows aerial dust.

Moreover, the lesson to be learned is that very intense and
localized light, by increasing the intensity of reflection, greatly en-
hances our powers of vision. And if the dust of the air, otherwise
quite invisible, thus becomes apparent to the naked eye, particles
too small for microscopic vision under ordinary illumination may be
seen under illumination analogous to that of the sunbeam. This is
the principle of the ultra-microscope, in which a beam of sunlight,
or from the electric arc, passes through a slit horizontally, or is
focused into a liquid which is examined by the microscope verti-
cally. Any light which enters the microscope must then have been
reflected from the surface of particles suspended in the liquid.

THE COLLOIDAL STATE 227

Particles having a diameter only one -hundredth that of the
smallest particles visible under ordinary illumination can then be
seen as spots of light like planets in the darkness. And so colloidal
liquids have been seen to be suspensions, and the size of the
suspended particles has been estimated by counting the number of
them in a volume of the liquid containing a known weight of
material. Thus the particles of platinum, gold, and silver seen
in colloidal suspensions of these metals have been discovered
to have diameters ranging from 2X10"" 4 to 6x10""* mm. The
smallest particles detectable by this method, when illuminated
by bright sunlight, have a diameter of 4xlO~* mm., whilst the
individual molecules of substances like chloroform and alcohol
have diameters of O^xlO"" 6 to 0-8xlO~ 6 mm., and of hydrogen
0-1x10-* mm.

Thus the particles of colloidal metals in aqueous suspension
have diameters about a thousand times as great as those of mole-
cules which form mixtures with water regarded as true solutions;
whilst the smallest particles that can be rendered visible have
only about ten times the diameter of gaseous and other simple
molecules.

These metallic suspended particles are not, however, molecules,
but rather minute fragments of solid metal; since molecules of
solid metals, consisting of definite aggregates of atoms, can scarcely
be said to exist. Thus, they differ from silicic acid and complex
organic substances which are known to consist of very large mole-
cules. The molecule of egg-albumin is estimated to have a mole-
cular weight of 17,000, and the molecules of the enzymes emulsin
and invertin have molecular weights of about 45,000 and 54,000
respectively, with molecular diameters of about 6 X 10 ~ 6 mm.
Such molecules can be seen by the ultra-microscope, but there
seems no hope that the simpler inorganic molecules will ever be
revealed to the eye of man, although they lie but a little way
below his range of vision aided by this powerful instrument.

Finally, although colloids cannot ordinarily be filtered, the fact
that parchment paper retains them suggests that special methods
of filtration might effect their separation as parchment paper does.
Special filters have in fact been prepared by treating ordinary
filter -paper with collodion or gelatine, which have pores varying
in diameter between 930 X 10" mm. and 21 x 10~ 6 mm. By
means of these filters various colloids have been differentiated and

228 CHEMICAL THEORY

classified according to the sizes of their particles, with the following

results:

Suspensions of non-colloids.

Colloidal platinum.
Colloidal ferric hydroxide.
Colloidal arsenious sulphide
Colloidal gold.
1 per cent gelatin.
Colloidal silicic acid.
Litmus.
Dextrin.

Solutions of crystalloids.

Thus it is seen that colloids afford a gradation between what
are commonly regarded as suspensions and solutions; and so it is
evident that the idea of a colloid as a glue-like substance has been
extended so as to include all kinds of matter in a state between
that of molecules and that of gross particles which are ordinarily
perceptible. Therefore such natural products as clay and lime are
regarded as colloids, and many precipitates formed in analysis are
believed also to pass through the colloidal state before separating
as particles that can be filtered. Indeed it appears that there are
two ways in which a precipitate may be formed and pass into a
state of equilibrium with the liquid phase with which it is in
contact. These are either by direct crystallization, as in the case
of magnesium ammonium phosphate, as well as calcium carbonate
at elevated temperature; or the formation and coagulation of a
colloid, as in the case of many sulphide and hydroxide precipitates.
Silver chloride and calcium oxalate are also examples of this kind
of precipitate, and barium sulphate is probably a limiting case
between the two kinds.

The coagulation of colloids is a subject which has received much
attention. It has already been pointed out that a colloidal sub-
stance such as silicic acid can exist in two conditions, those of "sol"
and "gel"; and that an electrolyte has the power of coagulating a
sol, thus converting it into a gel. It is now recognized that the
particles of a sol are electrically charged; in some cases such as
metallic hydroxides positively, in others such as metallic sulphides,
silver chloride, and silicic acid, negatively. Consequently a col-
loidal sol is capable of a kind of electrolysis, the colloidal matter
travelling either to the cathode or the anode when a current is

THE COLLOIDAL STATE 229

passed through the liquid. This motion of colloidal particles in an
electric field is called cataphoresis.

The coagulation of a sol by means of an electrolyte, as, for
instance, that of arsenious sulphide sol by means of hydrochloric
acid, with formation of a flocculent precipitate, is probably an
electrical phenomenon, in which the particles lose their charges and
consequently coalesce. The reverse process may, however, take place
in too concentrated solution. For example, precipitated basic ferric
acetate is prone to pass into the sol condition, chiefly as hydroxide
owing to hydrolysis, and so form a slimy mass impossible to filter.
The formation of a sol from a gel, which is the opposite of coagu-
lation, is called peptization.

Connected with peptization is the subject of protective colloids.
Certain organic colloids have the power of hindering the coagula-
tion of inorganic colloids with which they are associated. The
persistence of the sol condition is the result of the mutual repul-
sions of similarly charged particles; and when from any cause
these repulsions cease to operate coalescence of particles may take
place with coagulation of the colloid. If, however, the individual
particles of the sol are protected by an envelope of any material
which hinders their electric discharge, the sol condition will persist
oven in presence of an electrolyte which would otherwise cause
coagulation. Thus, as was found by Faraday, gold sol, formed by
the reduction of auric chloride solution by phosphorus, is stabilized
by gelatin, and the precipitation of silver chloride is prevented
by the same colloid, the less stable silver chloride sol being pro-
tected against the electrolyte present by the more stable gelatin sol.

So far the study of colloids in this chapter has been restricted
to phenomena of the liquid state. In the original and narrower
sense in which the term was employed colloids necessarily contained
a liquid phase, but just as the term " solution ", which is generally
applied to liquid mixtures, may be extended to include gaseous
and solid mixtures, so the term colloid is now extended to include
certain phenomena of the gaseous and solid states.

The distinctive property, however, of the colloidal state, which
differentiates it from the state of solution, is heterogeneity; that is
to say, the colloid is dispersed through the medium, there is a
disperse phase and a dispersion medium. So far the disperse
phase has been regarded as solid, and the dispersion medium as
liquid. When, however, the conception is extended to gaseous and

230

CHEMICAL THEORY

solid systems, a comprehensive scheme results in which a variety
of interesting phenomena are included. Thus Wo. Ostwald has
proposed the following classification of colloids:

Disperse
Phase.

Dispersion
Medium.

Examples.

Solid

Carbon particles in iron. Gold in ruby glass.

Solid

Liquid

Colloidal solutions of metals, gelatin, starch, &c.

Solid

Gas

Smoke. Fine dust. Fumes.

Liquid

Solid

Certain minerals.

Liquid

Emulsions.

Liquid

Gas

Fog. Mist. Clouds.

Gas

Solid

Solidified froths, e.g. pumice.

Gas

Liquid

Froths and foams.

The only system excluded from this scheme is that of a gas
dispersed in a gas. This system, however, is homogeneous, like
a true solution, there being no distinction as regards molecular
dimensions between its different components; therefore it is not
colloidal.

It is evident from a consideration of the table that a very large
number of phenomena encountered in nature and employed in the
arts come under the category of colloids. Indeed, as Ostwald has
said: " It is simply a fact that colloids constitute the most universal
and the commonest of all things we know. We need only to look
at the sky, at the earth, or at ourselves to discover colloids or sub-
stances closely allied to them. . . . We have only recently come
to learn that every structure assumes special properties and a
special behaviour when its particles are so small that they can no
longer be recognized microscopically, while they are still too large
to be called molecules. Only now has the true significance of this
region of the colloid dimensions The World of Neglected Dimen-
sions become manifest to us."

SUMMARY

DIALYSIS is the separation of substances in solution by the use
of a membrane, through which crystalloids in solution will pass
but not colloids.

COLLOIDS, e.g. silicic acid, can exist in two states, the hydrosol
(or sol) state, and the hydrogel (or gel) state.

Sols are converted into gels by catalysis.

THE COLLOIDAL STATE % 231

By means of the ultra-microscope colloidal liquids have been
seen to be suspensions.

Colloids have been separated by the use of special filters, and
a gradation has been established between colloids in suspension
and crystalloids in solution.

Colloids consist of electrically charged particles which can be
separated by an electric current, the process being called cata-
phoresis.

A sol passes into a gel by coagulation, a gel into a sol by
peptization.

More stable colloids can act as protective colloids to less stable
colloids, thus preventing their coagulation by electrolytes.

A colloidal system includes a disperse phase and a dispersion
medium. The system may be gaseous, liquid, or solid.

PART III CHEMICAL COMPOUNDS
AND CHEMICAL CHANGES

CHAPTER XIII
TYPES OF CHEMICAL COMPOUNDS

The chemical elements are divided broadly into metals and non-
metals. The metals form only a limited number of compounds
among themselves, these being chiefly crystalline products that
separate from molten alloys. The non-metals, taken alone, how-
ever, form a considerable number of prominent compounds, amongst
which are important hydrides, chlorides, oxides, and oxyacids. Ex-
cepting carbon compounds, which for the most part form a category
by themselves, the largest number of compounds contain both metal
and non-rnetal, for amongst these are all the metallic oxides,
sulphides, halides, and oxy-salts.

There is no doubt that oxygen, and to a less degree hydrogen, is
a key -element; so that the study of oxides, and to a less extent
hydrides, yields a deep insight into the properties of the elements
and their compounds. From oxides there is an easy transition on
the one hand to hydroxides and oxy-salts, and on the other to
halides, sulphides, &c., and so the chief regions of inorganic chemis-
try may be traversed systematically.

Therefore, these various compounds will be studied in the follow-
ing sequence: hydrides, oxides and hydroxides, halides, sulphides,
oxy-salts.

Hydrides.

Except for cuprous hydride (CuH) n , which seems to be unique,
the hydrides of the elements fall into two classes: (a) non-volatile
hydrides of powerful metals, (fe) volatile, generally gaseous, hydrides
of non-metals and metalloids.

(a) Metallic Hydrides. The following non-volatile metallic
hydrides, in addition to (CuH), are known:

LiH, NaH, KH, RbH, CsH, CaH 2 , SrH 2 , BaH 2 .

The metals forming them are those of the alkalis and alkaline

234 CHEMICAL THEORY

earths, i.e. the most intense of all the metals, which themselves are
able to decompose water at atmospheric temperatiire.

These hydrides are crystalline solids formed by the combination
of the respective metals with hydrogen, and are decomposed by water
with the evolution of twice the volume of hydrogen which is evolved
when the metal alone reacts with water. Thus calcium hydride,
known technically as " hydrolith ", reacts with water:

CaH 2 + 2H 2 O = Ca(OH) 2 + 2 H 2 .

Such hydrides stand in great contrast to the volatile hydrides of
the non-metals next to be enumerated.

(b) Non-metallic Hydrides. The following non-metallic hydrides

are typical:

(B 2 H 6) <kc.) CH 4 NH 3 OH, FH

SiH 4 PH, BH 2 C1H

GeH 4 Asfi 3 SeH 2 BrH

(SnH 4 ) SbH 3 TeH 2 IH
(BiH 3 )

These compounds, and the relationships between them, nave, how-
ever, been fully studied in the chapter on the periodic system, to
which the student is referred.

Oxides and Hydroxides.

The following classes of oxides can be distinguished:
Neutral oxides, including suboxides.
Basic oxides.

Acidic oxides, including mixed anhydrides.
Saline oxides.
Peroxides, divided into poly- and super-oxides.

Neutral oxides include the following:

H 2 0, CO, N 2 0, NO,

as well as the sub-oxides Cu 4 O, Ag 4 O, Pb 2 O. It is doubtful, however,
if these compounds constitute a valid class of oxides. Cu 4 and
Ag 4 O may be Cu 2 O + Cu and Ag 2 O -f Ag respectively, and Pb 2 O
seems to be a basic oxide since subaalts such as PbCl exist.

Water is a truly neutral oxide, or, more strictly, it is equally basic
and acidic, since by the minute degree of ionization that takes place
within it hydrogen and hydroxide Ions are necessarily produced in
equivalent quantities. The view has recently been expressed that
water should be regarded as a base since it combines with hydrogen
chloride like ammonia:

OH 2 + HC1 -> OH 3 - + Cl'; NH 3 + HC1 -> NH 4 - + 01'.

Carbon monoxide is very slightly soluble in water, but the

TYPES OF CHEMICAL COMPOUNDS 235

solution is neutral Nevertheless this oxide is related to formic
acid, HCOOH; thus, adopting conventional formula:

CO + H 2 O H-C-OH;

that is to say, CO is produced from HCOOH by dehydration, but
HCOOH is not produced from CO by hydration. The explanation
appears to be that hydroxylation of CO thus:

does not take place; for even when another oxygen atom is present
in CO 2 , to fortify the first oxygen atom, the acid produced, viz.
carbonic acid, is unstable:

^= O=C-OH.

If CO is brought into contact with KOH at 100 C., however, the
case is different; CO is slowly absorbed to produce formate, thus:

C=OH-KOH H-C-OK.

This reaction involves rearrangement of the atoms concerned, i.e.
intra-molecular rearrangement, probably in the following way:

^
KOH C<QK ' HC OK;

but the reaction is possible owing to the basic power of the potash,
and the stability of the salt produced.

Nitrous oxide, N 2 0, whilst distinctly soluble in cold water, pro-
duces no acid. It is nevertheless derived from hyponitrous acid,
H 2 N 2 2 , by loss of water, thus:

N=N N=N

I I V + HA
OH HO O

Like carbon monoxide, nitrous oxide is an acidic anhydride from
one side only; it is produced from an acid by loss of water, but does
not combine with water to form that acid. A true acidic anhydride
must produce an acid by combining with water; nitrous oxide is
regarded as a neutral oxide because its aqueous solution is neutral.
Nitric oxide, NO, is also classed as a neutral oxide, because so
far as it dissolves in water it yields a neutral solution. There is,

236 CHEMICAL THEORY

however, an uncommon acid, nitrohydroxylamic acid, from which
this oxide is derived by loss of water.

Basic oxides are numerous, and differ much in basic power.
The oxides of the alkali metals slake very vigorously, producing
hydroxides which are caustic alkalis. The oxides of the alkaline
earth metals also slake, producing hydroxides which increase in
solubility from calcium to barium. A few other metallic oxides,
e.g. MgO and Ag 2 O, dissolve very slightly in water, giving faintly
alkaline solutions.

Metallic hydroxides may be produced by the reaction of metal
with water, by the combination of the corresponding oxide with
water, or by precipitation from a corresponding salt solution by an
alkali. Occasionally all three methods of preparation are possible;
e.g. calcium hydroxide may be produced by the reaction of calcium
with water, thus:

Ca + 2H 2 = Ca(OH) 2 + H 2 ;

by the slaking of lime thus:

CaO + H 2 = Ca(OH) 2 ,

or by precipitation from a salt solution thus:

CaCl 2 + 2 NaOH = Ca(OH) 2 + 2 NaCl.

Of these three methods of preparation the first and second are
limited to the most electropositive metals such as sodium and
calcium, the third can alone be employed for such hydroxides as
those of copper, zinc, iron, and aluminium, on account of their
insolubility and the inertness towards water of the corresponding
oxides.

Acids, Bases, and Salts.

A base is generally defined as a compound which neutralizes
an acid to form a salt, water being eliminated in the process of
neutralization.

Originally the formation of water was overlooked, bases and acids
being considered oxides which combined to form a salt, e.g.:

CaO + SO 3 = CaO* SO*

TYPES OF CHEMICAL COMPOUNDS 237

Then it was discovered that hydrogen chloride is an acid which
forms a salt with a base with elimination of water, thus:
CaO + 2 HC1 = CaCl 2 -f H 2 O,

and so two kinds of acids were recognized: oxyacids such as H 2 S0 4
derived from acidic oxides such as S0 3 , and the so-called hydracids,
which, containing no oxygen, are not derived from acidic oxides.
And since acidic oxides were found to combine with water to form
acidic hydroxides or oxyacids, they were called acidic anhydrides,
or simply anhydrides. Thus the nomenclature on the acidic side
became definite; an acid was recognized as a compound containing
hydrogen replaceable by a metal, and acidic oxides were no more
called acids. The dualism which divides an oxysalt into basic and
acidic oxides is sometimes retained for convenience, however, espe-
cially in tabulating analytical data. A carbonate such as dolomite,
for example, will be said to contain so much CaO, so much MgO,
and so much CO 2 .

It must be recognized, nevertheless, that so far as reactions
depending on specifically acidic properties are concerned, no dis-
tinction is to be drawn between oxyacids and hydracids. Each
kind of acid consists of hydrion combined by electrovalency with
anion, and whether the anion of an acid contains oxygen or not is
of secondary importance.

It has been usual to represent sulphuric acid as if it were sul-
phury 1 hydroxide, SO a (OH) 2 , because it is related to sulphuryl
chloride, SO,C1 2 , from which it can be derived by hydrolysis. Such
a method of formulation, however, is not to be commended, since it
shows two of the oxygen atoms in the sulphate radicle differently
placed from the other two; and, moreover, no acid is a hydroxide
in the same sense as a base is a hydroxide. The modern method of
formulating sulphuric acid has been shown on p. 126.

The essential characteristic of an acid, then, is that it contains
loosely bound hydrogen ions which it easily parts with; and the
simplest conception of a base is that it is complementary to an acid,
and receives the hydrogen ions which the acid so easily gives away.
Thus an acid is a donor, and a base an acceptor of hydrogen ions. 1

There is, however, more than one way in which a base can
accept hydrogen ions. The way most generally recognized is that
of providing hydroxide ions to form water, as in the reaction:
Na' + OH' 4- H' + 01' = H 2 + Na" + 01';

1 Lowry, Chemistry and Indmlry^ 1923, 46.

238 CHEMICAL THEORY

and when a base thus provides hydroxide ions in solution, it is an
alkali as well as a base.

An anhydrous oxide can, however, behave as a base in the sense
of accepting hydrogen ions; for example, zinc oxide will dissolve in
an acid producing a salt, and thus neutralize the acid by accepting
its hydrogen ions and converting them into water:

Zn-O" + 2 H- + SO 4 " = H 2 O + Zn" + SO 4 ".

In both these cases water is produced by the neutralization of an
acid; ammonia, NH 3 , however, probably acts directly as a hydrogen
ion acceptor thus:

NH 3 + H- + 01' = NH 4 - + Cl'.

There is no need for the assumption which is generally made
that the ammonia first necessarily reacts with water to produce
ammonium hydroxide, NH 4 OH, before it can neutralize an acid.
It is true that NH 3 does react with water to form an alkaline
solution containing ammonium and hydroxide ions; but it is more
likely when added to dilute hydrochloric acid solution to react with
the hydrogen ions of the acid than with similar ions derived from
the water.

Acidic Oxides.

The oxides of non-metals are generally acidic oxides, and com-
bine more or less readily with water to form oxyacids. Thus
N 2 O 5 , P 2 6 , S0 3 , have a great affinity for water, and can with
difficulty be separated from it; in this respect they resemble the
oxides of the alkali metals, which combine with water with great
vigour to form hydroxides. Other non-metallic oxides, such as
C0 2 , S0 2 , B 2 3 , As 4 O 6 , do not so vigorously combine with water,
and are more easily separated from it; Si0 2 , probably because it
consists of polymerized molecules, does not directly combine with
water. The strongest acidic oxides, like the strongest basic oxides,
are those which combine most readily with water; and the feeblest
are those which, like silica, have very little attraction for water.

If the basic hydroxides are arranged roughly in order of
diminishing basic strength, and the acidic hydroxides similarly in
order of diminishing acidic strength, the interesting observation is
made that the two series overlap in the centre thus:
NaOH Ca(OH) 2 Mg(OH) 2 | A1(OH) 3

Diminishing basic strength.

HA1O(OH) 2

H 3 PO 4 H 2 S0 4 HNO 3

Diminishing acidic strength.

TYPES OF CHEMICAL COMPOUNDS

239

Thus A1(OH) 3 is both a feeble base and a feeble monobasic
acid, and is said to be amphoteric. 1 This is shown by the fact
that this hydroxide dissolves in acid to form an aluminium salt,
and in alkali to form an aluminate. That it is both a base and an
acid is shown by the following series of reactions which can easily
be carried out:

i. AlCl 3 + 3NaOH = A1(OH) 3 + 3 NaCl

(salt) (base) (base) (salt)

ii. HA10(OH) 2 + NaOH = NaAlO(OH) 2 + H 2 O

(acid) (base) (salt) (water)

iii. NaA10(OH) 2 + HC1 = HA1O(OH) 2 + NaCl

(salt) (acid) (acid) (salt)

iv. A1(OH) 8 + 3 HC1 = A1C1 3 + 3 H 2 O

(base) (acid) (salt) (water).

Thus A1(OH) 3 completes the cycle:

base; acid; acid; base.

The mineral known as spinel, MgOAl 2 O 3 or MgAl 2 O 4 , contains
alumina as aluminate; and analogous to this is chrome ironstone,
which is ferrous chromite, FeO Cr 2 O 3 . Other amphoteric hydroxides
are: Sn(OH) 2 , Sn(OH) 4 , Sb(OH) 3 , and perhaps Zn(OH) 2 ; though
the existence of Zn(ONa) 2 is improbable, and of Zn(OH)(ONa)
doubtful.

It may here be recorded that certain of the less electro-positive
metals form both basic and acidic oxides, the lower oxides being
basic, the higher acidic. The best-known examples of such metals
are chromium and manganese, to which may be added uranium on
account of an interesting point in connection with the trioxide of
this metal. The facts are set forth in the following table.

Cr.

Mn.

CrOj wholly basic.
CrjOto basic, feebly acidic,

e.g. in FeCr 2 O 4 .

CrOs, wholly acidic.
(CrOjCl 2 is not a salt, but
an acidic chloride).

MnO, wholly basic.
MnOp possibly basic,
feebly acidic.
MnO, wholly acidic.
J/?i 2 7 , acidic.
(MnOjhSOt suggests
basic property.

U0 2 , wholly basic.
U0 3 , acidic; basic with
regard to one oxygen
atom in the uranyl
salts,

e.g. UO,(NO^

These examples illustrate the acid-producing tendency of added
oxygen; they justify the name oxygen (acid -producing). The

1 1.e. facing both ways. See note in Appendix on Amphoteric Hydroxides.

240 CHEMICAL THEORY

compound (Mn0 3 ) 2 SO 4 , permanganyl sulphate, formed by dissolving
potassium permanganate in concentrated sulphuric acid, shows that
Mn 2 7 possesses a vestige of basic power; the uranyl salts, e.g. the
nitrate, U0 2 (N0 3 ) 2 , are interesting for a similar reason. Uranium
is the highest member of the sixth group of the periodic system;
therefore its oxides should be more basic than those of any other
member of the group. This is so, and basic power here extends
even to the trioxide, although it does not in the case of the less-
basic chromium trioxide. It is only one of the three oxygen atoms,
however, which can be replaced by acidic radicles so that basic
salts, the uranyl salts, are formed. These are nevertheless true
salts; the nitrate, for example, dissolves freely in water.

Mixed Anhydrides. There are a few non-metallic oxides which
combine with water or bases to produce two acids or salts. The
chief of these are nitrogen and chlorine peroxides.

Nitrogen peroxide, N0 2 or N.,O 4 , readily dissolves in cold water,
producing an equimolecular mixture of nitrous and nitric acids,
thus:

2N0 2 + H 2 = HN0 2 + HN0 3 ;

it is therefore regarded as the mixed anhydride of these two acids.

Chlorine peroxide or dioxide, C10 2 , does not form any acid with
water, although it yields a yellow solution from which the hydrate
C10 2 8H 2 O can be crystallized. With alkali, however, a mixture
of chlorite and chlorate is formed, thus:

2ClO 2 + 2NaOH = NaClO 2 + NaClO 3 + H 2 O.

In this limited sense, therefore, C1O 2 is a mixed anhydride.

Saline Oxides. It has been remarked above that the higher
oxides of certain metals are acidic whilst the lower are basic.
For example, Cr0 3 is acidic whilst Cr 2 s is basic. If potassium
chromate is added to a chromic salt solution, a dingy, yellow
precipitate is obtained which has the empirical composition CrO 2 .

This is basic chromic chromate, Cr 2 3 -CrO 8 or (CrO) 2 Cr0 4 ,
formed thus:

2 CrCl 3 + K 2 Cr0 4 + 2 H 2 O = Cr 2 O 3 -CrO 3 + 2 KC1 + 4 HC1;

it is therefore a saline oxide, composed of a basic and an acidic
oxide of the same metal.

The following are other examples of saline oxides:

Fe 3 O 4 , Mn 3 O 4 , Pb 2 O 3 , Pb 3 O 4 .

TYPES OF CHEMICAL COMPOUNDS 241

Fe B 4 , ferroso-ferric oxide, is a compound of FeO and Fe 2 O 8 . If
a mixture of ferrous and ferric salts is precipitated by alkali,
black hydrated FesO 4 is produced. Mn^O^ may be regarded as
MnOMn 2 O 8 , or perhaps 2MnOMnO 2 . The former constitution
is suggested by the behaviour of this oxide with concentrated
sulphuric acid, in which it dissolves, forming a mixture of man-
ganous and manganic sulphates, thus:

Mn 3 O 4 + 4 H 2 SO 4 = MnSO 4 + Mn 2 (SO 4 ) 3 + 4 H 2 O.

So Mn 2 O 3 , acidic relatively to MnO, is really also a basic oxide.

P6 2 3 and P6 S 4 are PbO-PbO 2 and 2PbO-PbO 2 respectively, i.e.
they are lead meta- and ortho-plumbate, related to the hypothetical
acids:

= Pb<g| and i8>PK8n
thus:

O = Pb<g>Pb and Pb<g>Pb<g>Pb.

It might be expected that tin, in the same group of the periodic
system as lead, would form similar saline oxides. That it does not
is probably due to the fact that SnO is not basic enough to combine
with Sn0 2 , for SnO and SnO 2 are undoubtedly more acidic than
PbO and PbO 2 respectively.

Peroxides,

A peroxide is an oxide which readily yields some of its
oxygen either as a gas or by behaving as an oxidizing agent.
Thus Pb0 2 , besides being an acidic oxide, is a peroxide which
behaves as follows when heated alone (i) or with hydrochloric

acid (ii): ^

i. 2PbO 2 = 2PbO + 2 .

ii. PbO a + 4 HC1 = PbCl 2 + 2 H 2 O + C1 2 .

This description, however, is not precise enough; for BaO 2 be-
haves like PbO 2 in regard to these two reactions, but is also capable
of a third reaction, with dilute acid, viz.:

BaO 2 + 2 HC1 = BaCl 2 + H 2 O 2 .

Now the difference between BaO 2 and Pb0 2 thus revealed is
fundamental, for lead belongs to the fourth periodic group, and
is quadrivalent in Pb0 2 , whilst barium, belonging to the second

(D60) 17

242 CHEMICAL THEORY

group, is only bivalent. The constitutions of these two oxides are
therefore ordinarily represented thus:

O=Pb=0;

and so BaO 2 yields H 2 O 2 with dilute acid thus:

/O H-O

Ba<M + 2 HC1 = BaCl 2 + I ,

a reaction of which PbO a is plainly incapable.

Peroxides are divided into two classes, of which the above are
examples. These are sometimes known as poly- and superoxides
respectively (Mendel^eff):

POLYOXIDES.

Pb0 2 .
MnO.
CIO.
N0 2

'2-

SUPEROXIDES.

BaO 2 .

NaA-
Ti<X

The term peroxide is applied to a particular oxide, irrespective
of any other property it may possess. Thus, amongst the poly-
oxides are C10 2 and N0 2 , which are also mixed anhydrides; and
amongst the superoxides are those which are basic: Ba0 2 and
Na 2 2 ; and those which are acidic: S 2 O 7 and Ti0 3 .

The acidic, like the basic superoxides, contain a chain of 2 oxygen
atoms in lieu of an increased valency of the nuclear atom. Thus the
constitutions of the two latter oxides are generally represented thus:

O O

Halides.

The binary compounds of the halogens fluorine, chlorine,
bromine, and iodine differ m*uch in character, according to the
diversity of the elements which form them. Attention may be
confined to the chlorides, and those of silicon and sodium may be
chosen first of all as typical of non-metallic and metallic chlorides
respectively.

Silicon tetrachloride, SiCl 4 , the chloride of a non-metal, may be

TYPES OF CHEMICAL COMPOUNDS 243

prepared by the union of its elements, or more usually by passing
chlorine over a heated mixture of silica and carbon, thus:

Si0 2 + 2 C + 2 C1 2 = SiCl 4 + 2 CO.

It is formed as a vapour, and may be condensed as a colourless
liquid which is soluble in such a solvent as benzene. It cannot be
prepared by the action of aqueous hydrochloric acid on silica, for it
is instantly decomposed, i.e. hydrolyzed, by water, thus:

SiCl 4 + 4H 2 O Si(OH) 4 + 4HCL

It is not, however, decomposed by concentrated sulphuric acid, for
silicon forms no sulphate.

Sodium chloride, Nad, presents an extreme contrast to silicon
chloride in mode of preparation and properties. It is obtained by
neutralizing base by acid in aqueous solution, thus:

NaOH + HCl NaCl + H 2 O,

and is consequently not decomposed by water. It is a crystalline
solid, insoluble in a solvent like benzene, which dissolves silicon
chloride; it is not volatile, except at high temperature, but it is
decomposed by concentrated sulphuric acid, because sodium forms
a sulphate, and HC1 is more volatile than H 2 SO 4 .

Now consider aluminium chloride, A1C1 3 . Alumina, A1 2 O 3 , is
amphoteric, and the chloride also is intermediate in character.
Thus it is prepared anhydrous as a sublimate by passing either
chlorine or hydrogen chloride over the heated metal, and it may
also be obtained in aqueous solution by dissolving the hydroxide
in excess of hydrochloric acid, and then crystallized as A1C1 3 6H 2 O.
If, however, the anhydrous or hydrated chloride is heated with
water, hydrolysis takes place and hydrated alumina separates.

So the following reaction is reversible:

A1C1 3 + 3H 2 O ^= A1(OH) 3 + 3HC1.

Contrast with this:

SiCl 4 + 4H 2 Si(OH) 4 + 4HCl,

and

NaCl + H 2 O NaOH + HCl.

Thus it is seen that, as with oxides, there are chlorides in which
basic and acidic characters predominate respectively, and chlorides
of intermediate or amphoteric character.

244 CHEMICAL THEORY

Sulphides.

In the course of qualitative analysis the student becomes
acquainted with various types of metallic and metalloidal sul-
phides. Non-metallic sulphides are not very important; carbon
disulphide will serve as an example of these.

Since oxygen and sulphur are in the same group of the periodic
system, a comparatively close analogy between oxides and sulphides
may be expected. Thus, for example, members of the following
pairs of compounds may be expected to show somewhat close
relationships to each other:

C0 2 ,CS 2 ; As 2 3 ,As 2 S 3 ; K 2 O, K 2 S;
and they do.

Carbon dioxide and carbon disulphide differ in physical pro-
perties no more than oxygen and sulphur differ; and they are
distinctly analogous to each other in chemical properties. Thus,
as carbon dioxide combines with basic oxides to form carbonates,
so carbon disulphide combines with basic sulphides to form thio-
or sulpho-carbonates:

CO 2 + 2 NaOH = Na 2 CO 3 + H 2 O.
CS 2 + 2 NaSH = Na 2 CS 3 + H 2 S.

These oxy- and thio-salts are respectively decomposed by acids,

thus:

Na 2 CO 3 + 2 HC1 = 2 NaCl + H 2 CO 3 .
Na 2 CS 3 + 2 HC1 = 2 NaCl + H 2 CS 3 .

Thiocarbonic acid, H 2 CS 3 , separates in the latter case as a liquid,
and is thus more stable than carbonic acid, H 2 C0 3 , which exists
only in dilute solution. These acids, however, readily decompose
into H 2 O and CO 2 and H 2 S and CS. 2 respectively.

Practice in qualitative analysis teaches the student the relation
between As 2 O 3 and As 2 S 3 . Thus As 2 O 8 or more correctly As 4 6
is an acidic oxide, dissolving in alkali hydroxide to form arsenite,
and similarly As 2 S 3 is an acidic sulphide dissolving in alkali hydro-
sulphide to form thioarsenite;

As 2 O 3 + 2 NaOH = 2 NaAsO 2 + H 2 O.

Frequently, however, As 2 S 3 is dissolved in NaOH, and then the
reaction is:

= 3 NaAsS 2 + NaAsO 2 + 2 H 2 O.

TYPES OF CHEMICAL COMPOUNDS 245

When, however, the solution formed in this way is acidified, As 2 S 8
is reprecipitated on account of its insolubility:

3 NaAsS 2 + NaAsO 2 + 4 HC1 = 2 As 2 S 3 + 4 NaCl + 2 H 2 0.

Basic sulphides resemble basic oxides, thus:

Na 2 + H 2 = 2NaOH
= 2NaSH;

hydrosulphide being analogous to hydroxide. When an alkali
sulphide is dissolved in water, as might be expected, it reacts with
water in this way:

Na 2 S + H 2 O = NaSH + NaOH,

and is strongly alkaline on account of the hydroxide formed.

Calcium sulphide, formed in the dry way as in the black-ash
process, is scarcely soluble in water; nevertheless it is slowly hydro-
lyzed by water, thus:

2 CaS + 2 H 2 O = Ca(SH) 2 +

Consequently it is not formed in presence of water; and calcium is
not precipitated from solution as sulphide.

A characteristic of many metallic sulphides is their exceeding
insolubility in water; and in consequence they are unacted upon by
this substance, and the corresponding metals are quantitatively pre-
cipitated from aqueous solutions of their salts by hydrogen sulphide,
as in the well-known analytical reactions. The sulphides of some
metals, however, whose oxides are feebly basic or amphoteric, cannot
be precipitated in presence of water, so that the hydroxides appear
in place of the sulphides.

Thus when ammonium hydrosulphide is added to an aluminium
salt solution, A1(OH) 3 is precipitated, owing presumably to hydro-
lysis of the hydrosulphide:

A1C1 3 + 3 NH 4 SH = A1(SH) 3 + 3 NH 4 C1.
A1(SH) 3 + 3 H 2 O = A1(OH) 3 + 3 H 2 S.

A similar reaction occurs with chromium.

Oxy-salts.

The formation of oxy-salts, such as sulphates, is a criterion of a
metal. A chloride may or may not be a salt, a sulphate necessarily

246 CHEMICAL THEORY

is. Consider, for example, the following series of chlorides and
sulphates:

Chlorides: PC1 6 SiCl 4 A1C1 3 MgCl 2 NaCl
Sulphates: A1 2 (SO 4 ) 3 MgSO 4 Na 2 SO 4 .

PC1 6 and SiCl 4 are certainly not salts, and phosphorus and
silicon form no sulphates; A1C1 8 is a chloride of intermediate char-
acter possessing some saline qualities, and A1 2 (S0 4 ) 3 is an imperfect
salt somewhat hydrolyzed by water; Mg01 2 is also hydrolyzed when
heated with water, MgS0 4 scarcely so; NaCl and Na 2 S0 4 are true
salts yielding neutral solutions in which there is no hydrolysis.
That the chlorides A1C1 3 and MgCl 2 appear more hydrolyzable than
the corresponding sulphates is probably due to the volatility of
hydrogen chloride, which escapes with the steam when solutions of
these salts are evaporated.

It appears that an element must possess a certain minimum
metallic strength in order to form a sulphate; to form an acid
sulphate a metal must be of the strongest character. Thus it is
only the sulphates of the alkali metals which combine with sulphuric
acid to form solid acid sulphates, e.g. NaHS0 4 and KHS0 4 ; though
a few other metals appear to form such sulphates in solution. The
study of carbonates is very instructive. The elements fall into five
categories as regards power to form carbonates:

i. No carbonates.

ii. Basic carbonates only,
iii. Normal as well as basic carbonates.
iv. Normal carbonates only.

v. Acid as well as normal carbonates.

Non-metals form no carbonates, and power to form basic carbon-
ates emerges amongst the metalloids. Thus, of the fifth-group
elements, nitrogen, phosphorus, arsenic, antimony, bismuth, the
last alone forms a carbonate, which however is basic, and so reveals
itself as more metallic than any other of these five elements.

The elements of the fourth group, carbon, silicon, germanium,
tin, and lead, show similar relationships. Lead is the only member
of these five that forms a carbonate, and this metal appears to be
more metallic than bismuth, because in presence of excess of carbonic
acid to prevent hydrolysis it yields the normal carbonate PbCO 8 .

A number of other metals more readily form basic than normal
carbonates; amongst these are copper, mercury, zinc, magnesium.

TYPES OF CHEMICAL COMPOUNDS 347

Undoubtedly the failure to form a normal carbonate is due not only
to the weakness of a metal but to the feebleness of carbonic acid.
Thus magnesium, which is prone to form a basic instead of a normal
carbonate by precipitation, shows no tendency to form a basic sul-
phate, because sulphuric acid is sufficiently powerful to enable its
salts to resist hydrolysis.

Incidentally it may be supposed that a normal carbonate is first
formed in the act of precipitation, and that the basic carbonate which
actually appears is due to subsequent hydrolysis. For example, in
the case of lead, the basic carbonate may be supposed formed from
the normal carbonate, thus:

3PbCO 3 +2H 2 O 2PbCO 3 -Pb(OH) 2 + H 2 CO 3 ,

Sometimes this process of hydrolysis may actually be observed.
For example, when sodium hydrogen carbonate is added to mercurous
nitrate solution the precipitate first formed is almost white, and
consists of the normal carbonate, the hydrolysis of which is pre-
vented by the bicarbonate present. On dilution and warming, how-
ever, the precipitate darkens, and basic carbonate, and finally oxide,
is formed, because of the decomposition of the alkali bicarbonate in
solution, and the consequent hydrolysis of the precipitated mercurous
carbonate.

The series of the metals must be traversed far in the upward
direction before those metals are reached which form normal but
not basic carbonates. Such are the metals of the alkaline earths
and the alkalis. Calcium carbonate, for example, is normal when
precipitated, and never becomes basic. This salt is soluble to a
minute extent in pure water, to which it imparts a faintly alkaline
reaction. It is true that hydrolysis of the dissolved salt takes place,
but this results in the formation of hydroxide and bicarbonate, thus:
2CaC0 3 + 2H 2 O ^ Ca(HCO 8 ) 2 + CaCOH),,

OH' ions showing alkalinity being due to the ionization of Ca(OH) 2 .
The same phenomenon appears more markedly in the case of alkali
carbonates, which show a strongly alkaline reaction. Thus, sodium
carbonate reacts with water in this way:

Na 2 C0 3

and so its strongly alkaline reaction is explained; for, according to
the ionization theory, a salt producing basic and acidic ions only, e.g.
2Na* and C0 3 ", would be neutral in reaction.

248 CHEMICAL THEORY

It has been seen above that only the most powerful metals form
hydrogen or acid sulphates; the same is true in regard to hydrogen
or acid carbonates. Thus it is only the alkali metals that form solid
hydrogen carbonates, and these increase in stability from sodium to
caesium in the series:

NaHC0 8 , KHCOg, KbHCO 3 , CsHC0 8 .

The alkaline earth metals, with ferrous iron and magnesium, form
unstable hydrogen-carbonates in solution however, e.g. Ca(HC0 3 ) 2 ,
and these hydrogen-carbonates, existing only in solution, differ from
the solid hydrogen-carbonates of the alkali metals by being more
soluble than the corresponding normal carbonates. Whilst, there-
fore, CaC0 3 is dissolved by passing carbon dioxide through its sus-
pension in water, NaHC0 3 is precipitated when the same gas is
passed through a cold saturated solution of Na 2 C0 3 .

The superior solubility of calcium carbonate in water containing
carbon dioxide in solution over its solubility in pure water is of
profound importance in nature; for it is the cause, not only of the
temporary hardness of water, but of the disintegration of calcareous
rocks, as well as of their original formation through the agency of
marine organisms, which form their shells from calcium carbonate
held in aqueous solution by carbonic acid. These facts are repre-
sented by the following reversible reaction:

CaCO 3 + H 2 O + CO 2 ^= Ca(HCO 3 ) 2 .
Hydrated Salts.

Water of crystallization is of common occurrence in crystallized
salts, and since its presence has a great influence on physical pro-
perties, the student must on no account ignore it in formulat-
ing a salt. The influence of temperature and atmospheric con-
ditions on hydrated salts will be dealt with in another place; it
may here be remarked, however, that when such salts are coloured
the corresponding anhydrous compounds are invariably of a different
colour. Thus, for example:

CuSO 4 -5H 2 O is blue; CuSO 4 is white.

FeSO 4 -7H 2 O , green; FeSO 4 white.

NiSO 4 -7H 2
CuCl 2 -2H 2 O
FeCl 3 -6H 2 O
CoCl 2 -6H 2 O
CoBr 2 -6H 2 O
CoI 2 -6H 2 O

deep green; NiSO 4 yellow,

bluish green; CuCl 2 brown,

yellow; FeCl 3 iron-black,

crimson; CoCl 2 blue,

dark red; CoBr 2 green,

dark red; CoI 2 violet

TYPES OF CHEMICAL COMPOUNDS 249

The proportion of water varies much in different hydrated salts;
ammonium oxalate, for example, has 1 molecule of water to 1 mole-
cule of salt, ordinary sodium phosphate has 12, and the alums have
24; whilst between these extremes there are salts containing 2, 5, 6,
V, and 10, and less frequently 3, 4, and 8 molecules of water. Occa-
sionally, too, the same salt will crystallize with varying proportions
of water according to the temperature of its formation. Thus, for
example, manganous sulphate, MnSO 4 , forms crystallo-hydrates with
1, 4, 5, and 7 molecules of water at different temperatures.

The way in which water is combined chemically in crystallo-
hydrates constitutes a problem the discussion of which is beyond
the scope of the present work; nevertheless it will be well to
tabulate here the commonest hydrated salts according to the mole-
cular proportions of water they contain:

H 2 O .... 2CaS0 4 -H 2 O.

H 2 O Na 2 CO 3 -H 2 Oj (NH 4 ) 2 C 2 O 4 -H 2 O.

2H 2 O BaCl 2 -2H 2 O; CuCl 2 -2H 2 O; CaS0 4 -2H 2 O.

3H 2 O K 4 Fe(CN) 6 -3H 2 O.

4H 2 O NaNH 4 HP0 4 -4H 2 O.

5H 2 O Na 2 S 2 O 3 -5H 2 O; CuSO 4 -5H 2 O; Bi(NO 3 ) 3 -5H 2 O.

6H 2 O CaCl 2 -6H 2 O; MgCl 2 -6H 2 O; CoCl 2 -6H 2 O; FeCl s -

6 H 2 O; CrCl 3 -6 H 2 0; FeSO 4 -(NH 4 ) 2 SO 4 -6 H 2 O,
and similar double sulphates.

7H 2 O MgSO 4 -7H 2 O; ZnSO 4 -7H 2 0; FeSO 4 -7H 2 O;

NiSO 4 -7H 2 O; CoSO 4 -7H 2 O.

8H 2 O Ba(OH) 2 -8H 2 O; BaO 2 -8H 2 O.

10H 2 O Na 2 C0 3 -10H 2 O; Na 2 SO 4 -10H 2 O; Na 2 B 4 O r -10H 2 O.

12H 2 O Na 2 HPO 4 -12H 2 O; Na 2 HAsO 4 -12H 2 O.

18H 2 O A1 2 (SO 4 ) 3 -18H 2 O.

24 H 2 O K 2 SO 4 Al^SO^ - 24 H 2 O, and other alums.

Double and Complex Salts.

DOUBLE SALTS are those which have a definite chemical indi-
viduality in the solid state, but break up more or less completely
in aqueous solution into their constituent single salts. Crystallized
potassium alum, K 2 SO 4 A1 2 (SO 4 ) 8 24 H 2 O, for example, is un-
doubtedly a chemical compound, and not a mixture of its two
constituent salts; but when dissolved in water it gives the separate
reactions of aluminium and potassium sulphates, so that its solution
contains a mixture of these two salts, the process of solution having
been accompanied evidently by disintegration of the double salt.

The formula for alum is sometimes halved, thus:

KA1(SO 4 ) 2 '12H 2 0.

250 CHEMICAL THEORY

Now it is always wise to accept the simplest available formula
in default of evidence to the contrary; but it may be objected that
the above formula suggests a complex rather than a double salt,
since it does not show complete molecules of the two constituent
sulphates. This objection would perhaps have little weight were it
not for a peculiar change which solid chromic alum undergoes when
heated to 90 0. The violet crystals then turn green, with loss of
water, changing into a salt which contains no free sulphate, since its
solution gives no precipitate with barium chloride. The change is
thus formulated:

K 2 S0 4 .Cr 2 (SO 4 ) 3 K 2 [Cr 2 (SO 4 ) 4 ] or 2 K[Cr(SO 4 ) 2 ].

So a double salt becomes a complex salt; potassium chromic
sulphate becomes potassium chromisulphate, the potassium salt of
chromisulphuric acid, HC^SO^g, a compound which is actually
formed when chromic sulphate is warmed with sulphuric acid.

That K[Cr(SO 4 ) 2 ] is so different from K 2 SO 4 -Cr(SO 4 V24 H 2 O
is a good reason for not writing the formula for any alum in a way
to suggest relationship to the former of these compounds.

Nevertheless there is evidence that saturated solutions of double
salts contain complex ions, and probably solutions of the various
alums form no exception.

Double salts are very numerous. Besides the double sulphates
and isomorphous selenates there are double chlorides, bromides, and
iodides, and less frequently double carbonates and nitrates.

DOUBLE SULPHATES. The alums, and salts of which ferrous
ammonium sulphate is a well-known example, may be mentioned.

Alums are isomorphous salts of the type

M 2 -SO 4 -X 2 "-(S0 4 ) 3 .24H 2 0,
where M = Na, K, NH 4 , Kb, Cs, Tl,
and X = Al, Fe, Cr, Mn, Ga, Ti, Rh.

They are formed by mixing the constituent salts in aqueous
solution, in proportions which may vary within wide limits, and
crystallizing. The alums are less soluble than their constituent
salts, and this is particularly the case with those of the extremely
electropositive metals rubidium and caesium.

Double sulphates of the ferrous ammonium sulphate type are
the salt FeSO 4 -(NH 4 ) 2 SO 4 -6H 2 O and others in which Mg", Zn",
Cu", Mn," Co", Ni" may take the place of Fe", and other alkali
metals that of NH 4 \ The relation between these double sulphates
and the heptahydrated sulphates, e.g. FeS0 4 -7H 2 O, is interesting.

TYPES OF CHEMICAL COMPOUNDS -251

It was found by Graham that one of the seven molecules of
water in this salt required a higher temperature for its expulsion
than the other six. This seventh molecule Graham called con-
stitutional water, because it appeared to enter into the constitution
of the salt more intimately than the other six molecules. It is
probably associated with the sulphate radicle, and is displaced by
ammonium or other alkali sulphate in the formation of the double
salt. The relationship may be thus shown:

FeSO 4 -H 2 O-6H 2 O : FeS0 4 -(NH 4 ) 2 S0 4 '6H 2 O.

It should be ^remarked that the ammonium sulphate in the
double salt exerts a protective influence over the ferrous sulphate,
for ferrous ammonium sulphate is less oxidizable by the air than
ferrous sulphate, and for this reason is preferred for the purpose
of volumetric analysis.

DOUBLE CHLORIDES. The mineral carnallite is KC1 MgCl 2 6 H 2 O,
to which there corresponds the ammonium salt NH 4 ClMgCl. 2 -6H 2 0.
The solubility of magnesium and manganous hydroxides in am-
monium chloride solution, with the corresponding fact that the
hydroxides of these metals are not precipitated by ammonia in
presence of ammonium chloride, is sometimes attributed to the
formation in solution of complex ions, such as (MgCl 3 )', derived
from NH 4 Cl-MgCl 2 . The salts themselves, however, are usually
regarded as double rather than complex salts. The double chloride
NaCl-AlCl 8 is a volatile compound, the formation of which was a
part of an early process for the preparation of metallic aluminium.

Examples of double salts of another type are sodium potassium
tartrate (Rochelle salt), NaKC 4 H 4 O 6 4H 2 O, microcosmic salt,
NaNH 4 HP0 4 4H 2 O, and magnesium ammonium phosphate,
MgNH 4 P0 4 -6H 2 0.

These are formulated differently from the alums and other
double salts, as containing two or more metallic radicles within
the same molecule. Since, however, these salts show no complex
ions in dilute solution, and their molecular magnitudes are unknown,
it may be that they should be put in the same category as other
double salts.

COMPLEX SALTS are those which, derived originally from single
salts, are so stable as to maintain their individuality in solution,
one of the metals appearing as a basic ion, whilst the other has
become part of a complex acidic ion, so that its metallic nature is

252 CHEMICAL THEORY

masked. Potassium ferrocyanide, K 4 Fe(CN) 6 , is a familiar example
of a complex salt. It appears to be composed of 4KCN + Fe(CN) 2 ,
and is indeed formed by adding potassium cyanide to ferrous sul-
phate solution until the precipitated cyanide has been redissolved,
and then boiling the solution. Thus a remarkable change takes
place; the iron ceases to behave as a basic radicle and becomes
part of an acidic complex, so that it gives no ferrous reactions in
solution. No ferrous salt is present, only a potassium salt potas-
sium ferrocyanide which ionizes in solution thus:

K 4 Fe(CN) tf 4 K' + [Fe(CN)J'''.

So profound is this change, and so stable the complex salt, that
from its concentrated solution sulphuric acid separates hydrof erro*
cyanic acid, H 4 Fe(CN) 6 , as a white solid.

Alum and potassium ferrocyanide, as representatives of double
and complex salts respectively, present extremes, but there are
gradations between them. The behaviour of nickel and cobalt
salts with potassium cyanide furnishes a case in point. The fol-
lowing reactions take place:

NiS0 4 + 2KCN = Ni(CN) 2 + K 2 SO 4 ; Ni(CN) 2 + 2KCN = K 2 Ni(CN) 4 .
CoSO 4 + 2KCN = Co(CN) 8 + K,SO 4 ; Co(CN) a + 4KCN = K 4 Fe(CN) 6 .

Both K 2 Ni(CN) 4 and K 4 Co(CN) 6 are complex rather than double
salts, for they do not contain nickelous and cobaltous ions; more-
over, K 4 Co(CN) 6 is plainly analogous to K 4 Fe(CN) 6 . From each of
these solutions, however, the simple cyanide Ni(CN) 2 or Co(CN) 2
is reprecipitated by dilute acid. These are examples of complex
salts, therefore, which are less stable than ferrocyanide. When a
solution of potassium cobaltocyanide is boiled in presence of air it
undergoes oxidation to cobalticyanide thus:

2 K 4 Co(CN) 6 + H 2 O + O = 2 K 3 Co(CN) 6 + 2 KOH,

and this latter salt is much more stable than cobaltocyanide, in this
respect resembling ferro- or ferricyanide. The fact that nickel
forms no such stable complex salt, nickelic salts being unknown,
underlies the well-known separation of these two metals.

The student meets with other examples of complex acids and
salts in the course of chemical analysis. Hydrofluosilicic acid,
H 2 SiF 6 , is evidently composed of 2HF-fSiF 4 , but it contains the
complex ion [SiF 6 ]". Potassium platinichloride, or chloroplatinate,

TYPES OF CHEMICAL COMPOUNDS 253

K 2 PtCl 6 , and the corresponding acid H 2 PtCl 6 , formed when platinum
is dissolved in aqua regia, are of the same type, and so is the
corresponding stannichloride, K 2 SnCl 6 . Potassium cobaltinitrite,
K 3 Co(NO 2 ) 6 , formed as a yellow crystalline precipitate when potas-
sium nitrite is added to a cobaltous solution acidified with acetic
acid, is of the same type as K 3 Co(CN) 6 and K 3 Fe(CN) 6 .

Ammonium phospho-molybdate is a complex salt of a different
kind, in which 12 molecules of MoO 3 are combined with (NH 4 ) 3 P0 4 .
It is formed in presence of nitric acid, and when dissolved by
ammonia suffers hydrolysis into simple phosphate and molybdate.
Potassium antimonyl tartrate, or tartar emetic, [KSbOC 4 H 4 6 ] 2 H 2 O,
is a complex rather than a double salt, for it dissolves in water
without hydrolysis, which antimonious salts will not do. It is
therefore best regarded as the potassium salt of antimonyl-tartaric
acid, [K(SbOC 4 H 4 6 )] 2 H 2 0.

SUMMARY
Types of Chemical Compounds

HYDRIDES. Metallic and non-metallic.
OXIDES AND HYDROXIDES.

Neutral oxides, including suboxides.

Basic oxides.

Acidic oxides, including mixed anhydrides.

Saline oxides.

Peroxides divided into poly- and superoxides.

HALIDES. Metallic and non-metallic.

SULPHIDES. Metallic, metalloidal, and non-metallic.

OXYSALTS. Sulphates, carbonates, &c.

HYDRATED SALTS.

DOUBLE AND COMPLEX SALTS.

CHAPTER XIV
CHEMICAL CHANGE IN GENERAL

A classical illustration of chemical change, at once simple and
valuable, is furnished by the work of Priestley and Lavoisier on
mercuric oxide. Priestley heated mercuric oxide by concentrating
the sun's rays upon it with a lens, so as to decompose it into
mercury and oxygen. The reaction is represented thus:

2HgO 2Hg + O 2 .

This mercuric oxide could previously be obtained, as was shown
by Geber, by gently heating mercury for a long time in the air,
when atmospheric oxygen united with the metal, thus:
2Hg + O 2 _* 2HgO.

Lavoisier combined these two operations by first heating mercury
at a moderate temperature in a confined space, and noting the volume
of air absorbed, and then collecting the mercuric oxide formed and
heating it more strongly; this resulted in the evolution of a volume
of oxygen equal to that of the air which was previously absorbed.
So the possibility of reversing a chemical reaction was established,
a fact now represented thus:

2Hg + O 2 ^ 2HgO.

This simple illustration has been chosen becaiise it gives rise to
various questionings, the consideration of which leads far into the
subject of chemical change in general.

Thus it is a surprising thing that the compound of mercury and
oxygen should be a red, crystalline powder, so different from its
constituent elements, and the question at once occurs whether the
whole of chemistry is full of surprises like that. The elementary
student is rather led to suppose that it is. At least there are many
such surprises which lend to chemical science a fantastic charm for
the youthful mind. For example, the vapour of sulphur is led over

264

CHEMICAL CHANGE IN GENERAL 255

red-hot charcoal, and, instead of yellow crystals and black lumps,
there appears a colourless liquid with a quite extraordinary smell;
or ammonia and hydrogen chloride gases are brought together, and
instead of a neutral gas resulting from the combination of an acid
and an alkaline gas, there is dense white smoke which settles down
as solid sal ammoniac.

" It is the unexpected that happens" might apparently be said of
chemical change. Indeed, the difference between physical mixture
and chemical combination often appears to be this the properties
of a mixture are what might be expected, they are the mean of
those of the constituent parts of the mixture, whilst the properties
of a chemical compound often could not be expected, for they are
unrelated to those of the constituent elements.

Yet, if the impression gained from the facts above considered
were true, and if the above epigram were a generalization of
chemistry then there would be no chemical science; chemistry
would be but a catalogue of curious material phenomena.

The scientific thinker is thus met with the fundamental question
of the relation between the properties of compounds and those of
their constituent elements; he is led, indeed, to the threshold of a
field of inquiry as broad as chemistry itself.

This particular inquiry may, however, be carried a little further.
The greatest differences in properties are seen between elements
and their simplest compounds. When a compound is converted
into another compound by the addition or substitution of other
elements, the physical differences brought about are not so great.
Consider, for example, the series of paraffin hydrocarbons
C w H 2n+2 (see p. 164). An increment of CH 2 causes no surprising
change in the properties of a hydrocarbon; on the contrary it
causes an almost constant alteration of boiling-point and other
physical properties. Or, consider the influence of the substi-
tution of chlorine for hydrogen in the CH 3 group of acetic acid,
CH 8 .COOH. The chloracetic acids CH 2 C1-COOH, CHC1 2 -COOH,
CC1 3 COOH, stand in order of increasing strength; thus the electro-
negative element chlorine has had a specific influence in increasing
the strength of the acid.

The colour, also, of a complex chemical compound is definitely
related to its constitution, and is modified by the substitution of
one element or group for another within the molecule. The art
of producing synthetic dyes depends among other things on a

256 CHEMICAL THEORY

knowledge of the influence of certain substituents on the colour
of the compound formed. The same is true regarding the thera-
peutic properties of synthetic drugs.

But, more generally, if, according to the periodic law, the
properties of the elements and their compounds are periodic
functions of the atomic numbers, then this law should at least
relate the properties of a particular compound to those of a similar
compound of an analogous element. That it does this is shown
by a systematic study of oxides, chlorides, and other simple com-
pounds. Our present knowledge of the law, however, fails to
account completely for the properties of a particular compound.
It is known, for example, that the solid iodides of imperfect metals
are brightly coloured, although the constituent ions are colourless;
such iodides are: PbI 2 , Hg, 2 I 2 , HgI 2 , SnI 4 , SbI 3 ; but why PbI 2 is
yellow, for instance, and SnI 4 scarlet, is not known, though probably
it is referable to the electronic constitution of the iodine atom, as
well as to the intimate mode of union of iodine with the metallic
atom in the solid state, which prohibits ionization. The colours of
ions such as green manganate and crimson permanganate may be
connected with the different modes of distribution of electrons in
the several external layers of the central atom.

These are examples of the questionings to which a consideration
of the superficial properties of a simple chemical compound gives
rise. The human mind desires an explanation of the unexpected.
Why is mercuric oxide red ? Why is permanganate solution crimson ?
These questions cannot at present be answered fully, though their
answer is bound up with the problems of atomic constitution.

Reversible Reactions.

A second question suggested by the reaction between mercury
and oxygen is that of the reversibility of a chemical change. Are
all reactions reversible? If not, why not?

Speaking generally, the possibility of reversing a chemical
reaction depends on the realization of suitable conditions. Some
chemical changes brought about by heat are so profound that
their reversal in the narrower sense is not possible. Sugar, for
example, is destroyed if heated strongly; it is commonly said to
be burnt, and the final products of the burning of sugar in air arc
carbon dioxide and water. Can such a change be reversed?

Growing plants can reconvert the carbon of carbon dioxide

CHEMICAL CHANGE IN GENERAL

257

and water into sugar, and the chemist can laboriously synthesize
a kind of sugar; but that is not a true reversal of the chemical
reaction of decomposition, because the synthetic changes do not
follow the same route as the changes of decomposition.

Now, reverting to the reaction between mercury and oxygen,
it would appear from Lavoisier's experiment that there is a certain
minimum temperature at which visible combination between the
elements takes place, and a somewhat higher temperature at which
visible decomposition of the compound formed sets in. These
temperatures cannot be stated because they are conditioned, but it
may be judged that below or above a limited range of temperature
oxygen and mercury do not combine or remain in combination
respectively.

Such a statement, however, is not very satisfactory, because,
whilst it recognizes the increased activity of mercury and oxygen
molecules, due to rise of temperature, which first promotes com-
bination between the two elements, and subsequently causes
disruption of the compound formed, it takes no account of the
physical state or concentration of the combining elements; or
otherwise, since oxygen is a gas, that it may escape from the
mercury altogether when evolved and so render a reversal of the
reaction impossible.

It is worth while to attempt to gain a clear mental picture
of this reversible reaction, since it illustrates fundamental principles
which underlie chemical reactions in general.

Chemical .Equilibrium.

Suppose the flask A in the figure con-
tains mercury and mercuric oxide in con-
tact with oxygen gas, the pressure of the
latter being indicated by the manometer B,
consisting of a U-tube containing mercury;
and suppose that the flask is heated in a
chamber, the outline of which is shown by
the dotted line, to a temperature t Q C., within
the limits between which a reaction be-
tween mercury and oxygen is known to

take place. Then the manometer will show in which direction the
reaction is proceeding. If combination is taking place, the mercury
will rise in the limb nearer the flask, owing to diminution of gaa

Fig. 45

(1)00)

268

CHEMICAL THEORY

pressure; if decomposition, the mercury in the nearer limb will
be depressed because of increase of pressure. But in either case
equilibrium will eventually result, and the mercury in the mano-
meter will become and remain stationary, registering a certain gas
pressure; this is not because nothing is taking place, but because
the two opposite reactions are proceeding at equal rates, and a
state of dynamic equilibrium has been attained, which is suitably
represented by the equation:

2Hg + O 2 ^ 2HgO.

If the temperature is altered the pressure will likewise alter. If,
for instance, the temperature is raised, further decomposition will
begin, but the accumulating oxygen will soon bring it to a stand-
still and a higher constant pressure will be registered corresponding
to the higher temperature.

This reversible chemical reaction, depending upon temperature,
is an example of thermal dissociation', and the pressure at which
equilibrium is reached at a given temperature is called the dis-
sociation pressure for that temperature.

The following dissociation pressures have been measured for
the reaction under discussion. 1

Temperature.
360 C.
380
400
420
440
460
480

Millimetres Hg Pressure.

90
141
231
387
642
1017
1581

If the temperatures are represented graphically by abscissae, and
the corresponding pressures by ordinates, the relation between
them is given by a curve which takes the form shown in fig. 46.

The curve may be thus interpreted: at the lower temperatures
there is little increase of pressure as the temperature rises, tendency
to dissociation is small, and a small pressure suffices to prevent it,
whilst a pressure greater than that necessary to prevent dissociation
promotes absorption of the gas with formation of more of the
oxide. At the higher temperatures, as shown by the steepness
of the curve, there is a great increase of pressure with small rise
of temperature, or, otherwise expressed, great pressure of oxygen

i Taylor and Hulett, /. Pkytikal Ckem., 1913, 27, 565.

CHEMICAL CHANGE IN GENERAL

259

is necessary to prevent dissociation, and combination apart from
such great pressure is impossible. Although this thermal dis-
sociation is a chemical change it closely resembles the physical
change involved in the evaporation of water, and the vapour-
pressure curve of water takes a similar form.

16UU
1600
1400
1300
1200
1100
o 1000
8 900
| 800
s 700
g 600

fi 500
400
300
200
100 (

360 380 400 420 440 460 480

Temperature.
Fig. 46. Dissociation Pressure Curve of Mercuric Oxide

If the student has followed the exposition of this simple chemical
reaction thus far he will be prepared for the conclusion that the
reaction depends on three determining factors. First, there is the
specific power of combination between mercury and oxygen. Com-
pared with other elements, mercury has little intrinsic attraction for
oxygen; if, for example, copper were substituted for mercury, the
result would be very different. The quality of mercury thus mani-

260 CHEMICAL THEORY

f ested may be called its chemical affinity, a convenient but originally
rather vague term.

Second, there is the effect of temperature. The power of mercury
to combine with oxygen remains latent until by rise of temperature
the molecules of metal and gas are made sufficiently active. Chemi-
cal change is always confined within certain temperature limits.
At very low temperatures matter is inert and incapable of chemical
change ; at very high temperatures matter is too active, its atoms
are too restless to enter into chemical union.

Third, there is the effect of pressure, that is of the concentration
of oxygen molecules around the mercury. If the surface of the
mercury in the flask is thought of as in a state of flux, some mole-
cules of the oxide being continuously produced whilst others are
simultaneously undergoing decomposition, it will be understood that
the chances of combination are increased and of decomposition
simultaneously diminished by the crowding of oxygen molecules
just above the surface. Thus increase of pressure promotes com-
bination, and diminution of pressure decomposition, altogether apart
from the question of temperature. In other words, the extent and
apparent direction of a chemical reaction depend on the active
'masses of the reacting substances.

These effects illustrate a general principle known as Le Chatelier's
Theorem, which may be expressed thus: If a system in equilibrium
is subjected to a constraint by which equilibrium is disturbed, a
reaction tends to take place by which the effect of the constraint
is destroyed. So in this case increase of oxygen pressure would be
a restraint which would be counteracted by combination of oxygen
with mercury until the pressure was reduced to its former value.

Thus, apart from temperature, chemical action is determined by
chemical affinity and active mass. By active mass is meant con-
centration when a substance is a gas or in solution, and in the case
of a solid the minimum amount necessary to secure equilibrium.
Thus the active mass of a solid is constant under the same physical
conditions; a very little solid is necessary in contact with a solution
to secure saturation; and more is superfluous.

It must be understood that the proportions in which elements
are in contact influence the extent of combination, but not the pro-
portions in which this occurs.

Thermochemistry.

There is one other aspect of this simple chemical reaction between

CHEMICAL CHANGE IN GENEBAL 261

mercury and oxygen which must be considered. The effect of
temperature on the chemical reaction has been noticed; but heat
is applied in quantity, and some of this heat is transformed into
other kinds of energy during the progress of the reaction.

In the decomposition of the oxide, for example, oxygen leaves the
solid and becomes a gas; not only is heat used up in the actual dis-
ruptive process, but also doubtless in the gasification of the oxygen,
as well as in the vaporization of some of the mercury, though the
latter heat is given up again if the mercury is condensed. Thus
heat is absorbed in the change

2HgO

and it has been estimated that when 2 gram-molecules (i.e. 433-2
grin.) of mercuric oxide are decomposed, producing 2 gram-atoms of
liquid mercury (401-2 grm.) and 1 gram-molecule of oxygen gas
(32 grm.) 44,000 calories are absorbed; and similarly, when these
quantities of the separate elements unite to form mercuric oxide, the
same quantity of heat is liberated.
These facts may be expressed thus:

2 HgO = 2 Hg + O 2 - 44,000 cals.
and 2 Hg + O 2 = 2 HgO -f 44,000 cals.

So mercury and oxygen in the system [2Hg + O 2 ] contain between
them an excess of energy over that in the system [2HgO] amounting
in terms of heat energy to 44,000 calories. It must, however, be
clearly understood that nothing is here said or known concerning
the total energy in either of these systems. The rise or fall of the
tide may be measured against a rock in a sea the depth of which is
unplumbed. Similarly, the rise or fall of energy content in a
material system may be measured although the total energy in the
system cannot be estimated. The energy of a system which is
transferable by chemical change is called the free energy, the energy
which is not transferable is the bound or latent energy, these two
make together the total energy in the system.

Now, every distribution of matter which we call a chemical
change is accompanied by a corresponding distribution of energy.
The branch of chemistry concerned with energy changes manifested
by heat phenomena is called thermo-chemistry, the kind of equation
which sets forth heat change accompanying changes in matter is called
a thermo-chemical equation, and such an equation is based upon the

262 CHEMICAL THEORY

principle of the conservation of energy, just as an ordinary chemical
equation, which is a mass equation, is based upon the principle of
the conservation of mass.

It will be well to consider now some representative thermal
equations chosen to illustrate the range of the subject. And first
it may be remarked that thermal equations need not necessarily be
thermo-chemical equations. They may express thermal phenomena
accompanying physical rather than chemical change.

For example, the transformations between ice, water, and steam,
may be represented thus:

H 2 O (water) = H 2 O (ice) + 1440 cal.
H 2 O (steam) = H 2 O (water) + 9670 cal.

These equations express, in perhaps unusual guise, the latent heats
of water and steam.

Or the heat of transformation of one allotropic form of an ele-
ment into another may be represented thermo-chemically thus:

S (monoclinic) = S (rhombic) + 64 cal.
C (graphite) = C (diamond) + 500 cal.

Thus when 32 grm. of monoclinic sulphur are converted into the
same weight of rhombic sulphur, 64 calories are evolved; and simi-
larly in the conversion of 12 grm. of graphite into diamond 500
calories would be evolved. These figures are not, however, deter-
mined directly, but are derived from the differences between the
heats of combustion of different forms of the element. For example,
the heats of combustion of graphite and diamond are:
Graphite, 94,810 cal.; diamond, 94,310 cal.;

so the difference, 500 calories, must represent the heat of transfor-
mation of graphite into diamond.

It may next be observed that in the case of a reversible change,
whether physical or chemical, if the evolution of a certain quantity
of heat accompanies the change in one direction, the absorption of
the same quantity of heat takes place when the change is reversed.

This may be represented simply by transposing the terms of the
thermal equation; thus in the case of ice and water

H 2 O (water) = H 2 O (ice) + 1440 cal.;
but H 2 O(ice) = H 2 O (water) 1440 cal

Thus when 18 grm. of water are frozen, 1440 calories are removed

CHEMICAL CHANGE IN GENERAL 263

from the water, and to melt the ice a similar quantity of heat must
again be supplied.

Changes in which heat is evolved are exothermic changes; those
in which heat is absorbed are endothermic changes.

These terms are of value in considering chemical changes in-
volving much heat energy. Thus, reverting to the original example,
the formation of mercuric oxide from its elements is an exothermic
change; the decomposition of the compound into its elements an
endothermic change; the two changes, as regards energy, as well
as matter, are equal and opposite. One, the heat of formation, is
+ 44,000 cal.; the other, the heat of dissociation, is 44,000 cal.

Other kinds of thermo-chemical data of value are heats of
solution of substances in water and other solvents; heats of evapo-
ration of solvents containing specific dissolved substances; heats of
neutralization of acids and alkalis. It may be pointed out that
when chemical operations are being conducted on a large scale,
thermal effects, often negligible on the small scale, assume much
importance. For example, how to divert or usefully to dispose of
the heat evolved in a chemical action, or how most economically
to conduct the evaporation of a given solution, having regard to
its latent heat of evaporation, often become problems of pressing
importance to the chemical engineer.

Here may be mentioned the law of Hess, which, briefly, states
that the heat of formation of a compound is independent of its
mode of formation. This law follows naturally from the principle
of the conservation of energy, but it is capable of experimental
proof.

Lastly, tES<jonni6ction between heat evolved and chemical
affinity may berorijBlJy considered.

Exothermicyconroounds are stable; endothermic compounds often
unstable. Exothermic compounds may be compared with a boulder
at the base of a mountain which has reached its position by rolling
down the mountain's side, expending its potential energy the while,
but havwag^ reached the lowest ground is likely to remain there.
Endotfitohflfc compounds may be compared with a boulder raised
and p<w above the valley, and liable, after a slight disturbance,
to roll <JB(pm and expend energy in its fall.

High-explosives, and detonators such as fulminate of mercury,
are good examples of endothermic substances, as well as ozone,
.acetylene, and carbon disulphide.

264 CHEMICAL THEORY

Ozone is formed by the absorption of energy from the electric
sparks which generate it from oxygen; acetylene derives its in-
herent energy from the electric furnace in which calcium carbide
is produced; carbon disulphide is similarly formed from its elements
at high temperature.

These compounds are evidently highly artificial, and unlikely
to be produced by nature in her quieter moods, though electric
or volcanic action might produce them. Solar energy is, however,
continuously being absorbed by growing plants, and the compounds
produced by them, such as starch and sugar, may, in a wider sense,
be regarded as endothermic. Similarly, wood, coal, and petroleum
are in this sense endothermic, though it is unusual to extend the
term to include these things.

In general, endothermic compounds are likely to be produced at
high temperature, because the heat to be absorbed in their forma-
tion is thus available; whilst exothermic compounds are likely to
be formed at lower temperature, because the heat evolved in their
formation can more easily be dispersed.

The connection between the heat effect of a chemical change
and the chemical affinities of the reacting substances may now be
considered briefly.

It is generally supposed that if the reaction between two
chemical elements is vigorous and accompanied by the evolution
of much heat, the chemical affinity between these elements is great;
and that if the reaction is sluggish and accompanied by small
heat evolution, the affinity between the elements concerned is small.
This is illustrated by the heats of formation of the halogen hydracids,
which are:

HF 38,600 calories.

HC1 22,000

HBr 8,400

HI - 6,000

Thus the heats of formation fall from HF to HI, and that of
the latter is actually negative; HI is an endothermic compound.
The affinities of the halogens for hydrogen undoubtedly fall from
fluorine to iodine, but that of iodine cannot be regarded as actually
negative, as the thermal value might suggest. So that whilst the
heat effect of chemical union is some indication of the strength of
attraction between the elements, it cannot be regarded as a quan-
titative measure of chemical affinity.

CHEMICAL CHANGE IN GENERAL 265

There are other considerations bearing on thermo- chemistry,
and the connection between heat effect and chemical affinity, which
the student may postpone to a more advanced course.

Bate and Limits of Chemical Change.

The rate of a chemical change is much influenced by the condi-
tions under which it takes place. Thus the physical state of a
reacting substance or substances determines whether or how a
chemical change occurs. So it is often the case that substances
which interact in solution will remain in contact with one another
in the solid state without chemical change. The familiar example
of a mixture of sodium bicarbonate and tartaric acid, which effer-
vesces when water is poured upon it, may be quoted.

The modifying influence of temperature upon chemical change
is noteworthy. Consider, for example, the formation of water from
its elements or its dissociation into them.

If hydrogen and oxygen gases are mixed together in the propor-
tions necessary to form water, no change is perceived at atmospheric
temperature, but if the mixed gases are gradually heated, combina-
tion becomes perceptible at about 450 C., whilst at 700 C. there is
a rapid and complete combination, accompanied by explosion. It is
possible, however, to raise the mixed gases to such a temperature
as not to promote but actually to prevent their combination. Thus
above 2000 C. steam is decomposed into its elements, and hydrogen
and oxygen will not combine. So there are temperature limits
within which the combination of hydrogen and oxygen can take
place, but outside which no combination between these gases
occurs. And what is true of this particular change is true of all
chemical changes. Chemical reactions, vigorous under ordinary
conditions, become sluggish with reduction of temperature, and
cease altogether at very low temperatures. The velocity of a
reaction is frequently halved by a fall of 5 C., and if this rate
of diminution proceeds regularly, the velocity will become infini-
tesimal, the reaction coming practically to a standstill before the
temperature of liquid air is reached. On the other hand, chemical
union between the elements is impossible in the hotter stars; only
in the coldest stars, for example, do hydrocarbons exist.

The effect of heat in promoting change is too common to need
specific illustration. The use of a Bunsen gas flame to further
chemical action soon becomes instinctive to the student.

266 CHEMICAL THEORY

Catalysis,

Further study of the interaction of hydrogen and oxygen
reveals the fact that the degree of combination between these gases
is much influenced by the presence or absence of water vapour.
Thus, when the pure gases have been dried perfectly by long
exposure to phosphoric oxide, they do not combine at a red heat,
and silver can be melted in a mixture of them at a temperature
of 960 C.

Minute quantities of other foreign substances also affect the rate
of combination, and this is notably the case with certain metals,
especially if they are finely divided. Thus, when the mixed gases
are passed through a heated capillary tube containing a thread of
platinized asbestos, gradual combination takes place, whilst the
thread becomes red-hot, and would eventually cause the gases kept
in contact with it to explode. The platinum, however, remains
chemically unaltered at the end of the process.

These are examples of catalytic action, or catalysis, and the
promoter of the chemical change is called a catalyst. Catalysis
plays an important part in modern chemistry, and especially in
chemical industry. A brief study of the subject may therefore be
undertaken here.

The student will already be familiar with several cases of
catalysis. He knows that manganese dioxide causes potassium
chlorate to yield its oxygen at a lower temperature than that at
which the pure salt is decomposed; also that "nitrous fumes"
promote the formation of sulphuric acid from sulphur dioxide,
atmospheric oxygen, and water by a cycle of reactions which may
be represented thus:

2H 2 S0 4

So nitric oxide carries oxygen from the air to sulphurous
acid, being alternately oxidized and reduced. Platinized asbestos,
however, causes sulphur dioxide and oxygen to combine thus:
2S0 2 +O 2 = 2S0 3 , without itself undergoing perceptible change.

Cuprous chloride, employed in the Deacon process in the
oxidation of hydrogen chloride to chlorine and water, behaves
similarly to nitrous fumes in the chamber sulphuric acid process,
as an oxygen carrier, thus:

CHEMICAL CHANGE IN GENERAL 267

, * 2CuCl 2 =

| Cu 2 Cl 2 +O = Cu 2 OCl 2 ,

Cu 2 OCl 2 + 2 HC1 = 2 Cu01 2 + H 2 O ;

the net result being seen by adding these reactions thus:
2HC1 + = H 2 + C1 2 .

These well-known examples of catalysis illustrate the variety
of the mechanism of the reactions which may occur. Thus in the
case of the older sulphuric acid process, as well as in the Deacon
process, there is no doubt that a series of reactions takes place,
involving the catalysts, whilst finely divided platinum does not
appear to be altered chemically when it acts as a catalyst. Man-
ganese dioxide, however, probably undergoes a cycle of reactions
when heated with potassium chlorate, being oxidized to a higher
oxide which is immediately decomposed again. This view is sup-
ported by the fact that when crystallized manganese dioxide is
employed to assist the evolution of oxygen from potassium chlorate,
the dioxide recovered from the process is in a finely divided state.

The following are well-known examples of catalytic agents
and catalytic action,

i. Influence of Water.

It has already been seen that a trace of water-vapour greatly
influences the combination of hydrogen and oxygen to form water.

In the entire absence of water the following reactions will not

take place:

2CO + O 2 = 2CO 2 .

2NO + O 2 = 2NO 2 .

H, + C1* = 2HC1.

CaO + SO 3 = CaSO 4 .

CaO + 2NH 4 Cl = CaCl 2 + 2NH, + H 2 O.

Also, mercuric chloride and ammonium chloride do not dissociate
when quite dry, and sodium and potassium can be distilled in
perfectly dry oxygen without combustion. There is thus a variety
of reactions which depend for their progress on the influence of
water as a catalyst.

ii. Influence of Hydrogen and other Ions.

An ester, such as ethyl acetate, may be formed or hydrolyzed
according to the following reversible reaction:

2 ^ C 2 H 6 .C 2 H 3 2

268 CHEMICAL THEORY

The progress of the reaction in either direction, however, is
exceedingly slow in the absence of a catalyst; but hydrogen
chloride, by reason of hydrogen ions, under suitable conditions
accelerates both the formation and the hydrolysis of the ester.

The hydrolysis of cane-sugar thus:

C 12 H 2 Ai + H 2 = C 6 H 12 + C 6 H 12 6 ,

(dextrose) (levulose)

is also promoted by hydrogen ions, introduced as dilute acid, and
the extent of hydrolysis, in this case as well as in that of an ester,
is directly proportional to the concentration of the hydrogen ions;
consequently, these reactions can be used to compare the strengths
of acids. As regards the added acid, analysis proves that its
amount remains unchanged.

Another interesting example of the catalytic influence of
hydrogen and other simple ions is the coagulation of a colloid.
If hydrogen sulphide gas is passed into a dilute aqueous solution
of tartar-emetic potassium antimonyl tartrate there is no pre-
cipitate, but a deep orange, faintly-opalescent liquid results. This
liquid contains antimonious sulphide in the colloidal state. If then
a few drops of dilute hydrochloric acid or ammonium chloride are
added, an orange precipitate is formed, the colloidal state of quasi -
solution having been disturbed by the ions of the acid or salt added.

iii. Influence of Finely Divided Metals.

The catalytic influence of platinized asbestos has already been
noticed. The contact process for the manufacture of sulphuric
acid depends on the synthesis of sulphur trioxide in presence of
platinized asbestos.

The realization of this reaction on a commercial scale was
found to depend on keeping the platinum free from dust, and
especially from arsenious oxide, derived from the iron pyrites used
as the source of sulphur. The arsenic has been said to "poison"
the catalyst; it may be regarded as a negative catalyst, counter-
acting the influence of the finely divided metal. Scrupulous
cleanliness is always necessary for catalytic action at solid surfaces.
Platinum black, precipitated from platinic solutions by reducing
agents, is also very active in promoting gaseous synthesis. Finely
divided osmium and uranium have been found valuable in promoting
the synthetic formation of ammonia; platinum, as well as iron
containing copper or bismuth, is employed in bringing about the

CHEMICAL CHANGE IN GENERAL 269

oxidation of ammonia to nitric acid; finely divided nickel at 250 C.
causes hydrogen to be added to unsaturated hydrocarbons and their
derivatives, a reaction employed for the hardening of fats.

Another example of the catalytic action of a finely divided
rnetal is that of silver on hydrogen peroxide.

If aqueous hydrogen peroxide is added to silver oxide the
oxide is reduced to metal with evolution of oxygen, and then if
more hydrogen peroxide is poured upon the metal the evolution
of gas continues. Alternate oxidation and reduction of the silver
may here be postulated thus 1 :

2 Ag + H 2 2 = Ag 2 + H 2 O.
Ag 2 O + HA = 2 Ag + H a O + 2 .

Colloidal metals are particularly active owing to their state
of minute subdivision. Bredig has obtained platinum, silver, &c.,
in colloidal suspension in water by striking an electric arc between
poles of the metal beneath the surface of the water. In this way
the metal becomes diffused through the water in a colloidal state,
and is very active catalytically.

iv. Infliience of Oxides and Salts.

Various oxides and salts may act as catalysts; and the use of
oxides of nitrogen in sulphuric acid manufacture, of manganese
dioxide in the preparation of oxygen, as well as of cuprous chloride
in the Deacon chloride process, have already been noticed.

The use of cobalt oxide in the preparation of oxygen from
bleaching powder is a well-known lecture experiment; and the
role of copper oxide in the combustion and ultimate analysis of
organic compounds may be regarded as catalytic because of alter-
nate reduction and oxidation by the compound and atmospheric
oxygen respectively. Ferric oxide is employed in the manufacture
of hydrogen by the reaction:

C0 + H 2 C0 2 + H 2 ;

in the removal of sulphur compounds from coal gas, the "spent
oxide" being revivified by the combustion of its sulphur in sul-
phuric acid manufacture; and also in the promotion of the reaction:
2H 2 S + 2 2H 2 O + 2S,

1 Hydrogen peroxide, moreover, is very sensitive to catalytic action; precipitated man-
ganese dioxide will decompose it, as well as finely divided platinum, and even powdered
glass. Indeed, sharp points and edges promote its decomposition. If a concentrated solution
of hydrogen peroxide is poured into a glass vessel the surface of which has been scratched,
the evolution of gas bubbles along the line of the scratch may be observed.

270 CHEMICAL THEORY

by which sulphur has been recovered from alkali-makers' waste
by the Glaus process.

Vanadic oxide is used as a catalyst to promote oxidation in
some processes of organic chemistry; cerium oxide in the Welsbach
gas mantle is believed to act as a catalyst in promoting combustion,
whilst magnesia and other oxides are employed as catalysts to
promote the surface combustion of furnace gases.

As regards salts, manganous sulphate is known to act as a
catalyst in promoting the oxidation of oxalic acid by permanganate
in volumetric analysis, and ferrous sulphate likewise hastens the
liberation of iodine from potassium iodide by hydrogen peroxide,
whilst silver nitrate is well known to promote the oxidation of
small quantities of manganese in solution to permanganate by
persulphate.

v. Influence of Enzymes.

Enzymes are unorganized ferments produced by living organisms,
and they are very powerful catalysts.

For example, diastase, produced in germinating seeds, converts
starch into sugar, thus:

2C 6 H 10 6 + H 2 = C 12 H 22 O n ,

and so renders stored up food available for the growing plant;
zymase, in yeast, also converts sugars of the type C 6 H 12 O into
alcohol and carbon dioxide, thus:

C C H 12 O = 2C 2 H 6 O + 2CO 2 .

Similarly, there are unorganized ferments in the human
alimentary canal, which aid the digestion of food. All these
actions are catalytic.

Theory of Catalysis.

The foregoing examples of catalysis are sufficient to show the
wide influence of this phenomenon; indeed, according to Ostwald,
" there is probably no kind of chemical reaction which cannot be
influenced catalytically, and there is no substance, element or com-
pound, which cannot act as a catalyser ".

When various examples of 'catalysis are reviewed with the
purpose of forming a general theory of their action, it becomes
evident that some catalysts, such as the oxides of nitrogen, are
definitely chemical in their action, forming part of a cycle of
reactions, whilst others, such as finely divided platinum, are of

CHEMICAL CHANGE IN GENERAL 271

such a nature that it is difficult to postulate for them chemical
participation in the reactions which they promote.

Nevertheless platinum, like manganese dioxide, is found to
undergo change when it is employed as a catalyst. Thus photo-
micrographic study of platinum gauze before and after use shows
that the metal becomes coated with minute craters which are
tinged with grey or black platinum. Some chemists, indeed,
believe that an oxide of platinum is formed and reduced during
the catalytic activity of the metal.

An alternative view of the action of platinum is that the finely
divided metal presents a very large surface on which the reacting
gases are condensed and brought into an intimate contact which
promotes their chemical combination. Even in this case, however,
the action must probably be regarded as chemical, for surface
action is now known to be brought about on the external molecular
layer of a solid by forces which are chemical rather than physical
in nature.

According to Lowry, the modern view of the action of catalysts
may be summarized thus: "A catalyst is a reagent which is repro-
duced as one of the products of the reaction."

A catalyst is generally present in relatively small amount, and
its amount may be minute. Thus it serves over and over again, and
may remain active for an indefinite length of time. The velocity of
the reaction, however, is often proportional to the amount of catalyst
present.

Further, a catalyst is believed by some chemists not to initiate,
but only to hasten, or sometimes to retard, a chemical change. It
is true that in the absence of the catalyst the change often appears
not to occur, but this may be due, as in the synthesis of water, to
the reaction being exceedingly slow.

A catalyst has been compared with the oil which facilitates the
running of a machine; and just as the friction between the parts
of a machine may be too great to allow movement to begin until
oil is added, so there may be chemical reactions which, under
certain conditions, cannot begin until the necessary catalyst is
present. This appears to be the case regarding the influence of
water on certain reactions.

Moreover, a catalyst does not effect the final state of equilibrium
of the reacting substances; it only alters the time in which this
equilibrium is reached. This must be the case, since it contributes

272 CHEMICAL THEORY

no energy to the reaction, but only provides a stimulus; otherwise
a kind of perpetual motion would have been discovered. Lastly,
where intermediate reactions are known to occur, they must, when
taken together, necessarily proceed faster than the original reaction
in order to hasten it; there must be " a saving of time in the longer
way round ".

SUMMARY

THERMAL DISSOCIATION. Thermal dissociation is a reversible
chemical change caused by heat.

Exo- AND ENDO-THERMIC COMPOUNDS. An exothermic com-
pound is a compound produced from its elements with the evolution
of heat.

An endothermic compound is a compound produced from its
elements with absorption of heat.

The heat of formation of a compound is the heat evolved in its
formation from its elements.

CATALYST. A catalyst is a substance which alters the velocity
of a reaction, and is reproduced as one of its products.

CHAPTER XV
CHEMICAL CHANGES CLASSIFIED

It is customary to classify chemical reactions in the following
manner:

I. Combination of elements or compounds, e.g.:

2 = 2HgO.
SO 3 + H 2 O = H 2 S0 4 .

II. Decomposition of compounds into simpler compounds or ele-

ments, e.g.:

CaCO 3 = CaO + CO 2 .
Ag 2 O

III. Re-arrangement of the atoms within a molecule, e.g. the
formation of urea from ammonium cyanate, thus:

NH 4 CNO = CO(NH 2 ) 2 .

IV. Condensation of two or more molecules into one molecule
(polymerization), e.g.:

V. Single displacement, e.g.:

CuSO 4 + Zn = ZnSO 4 + Cu.

VI. Double decomposition, e.g.:

BaCl 2 + H 2 SO 4 = BaSO 4 + 2 HC1.

Such a classification, however, is not exhaustive; it is rather
formal, and not very illuminating, because it overlooks cause and
effect in chemical change; and reactions which are similar in form
may differ essentially in nature.

Compare, for example, the above reaction of double decomposi-
tion with the following:

2 NaCl + H 2 SO 4 = Na 2 SO 4 + 2 HC1.

(DCO) 273 19

274 CHEMICAL THEORY

The two reactions are similar in form, but quite different in cause
and effect.

More chemistry may be learned by studying reactions according
to the conditions under which they take place than according to a
few set types accepted a priori.

Thus, the study of changes effected by the action of heat on
solid substances, either singly or mixed, or by bringing together
solids and liquids such as non-metals and metals and their com-
pounds on the one hand, and water, acids, and alkalis on the other,
or of changes occurring in aqueous solution, will be found to cover
a very wide field of chemistry, and yield much insight into the
nature of chemical change. Or, again, the phenomena of oxidation
and reduction well repay classification and thoughtful study.

Chemical Action of Heat on Compounds.

It has already been seen that the change in a substance produced
by heat may be reversible or irreversible, and that in the former
case the change is one of dissociation, in the latter one of decomposi-
tion. A change which is really dissociation may, however, be
described as decomposition if one or more products are gaseous and
escape. Thus, limestone may be said to be decomposed by heat
because the carbon dioxide is allowed to escape; similarly the action
of heat on lead nitrate will generally be regarded as one of de-
composition unless the reversibility of the reaction is recognized.

Thermal Dissociation.

The general principles of thermal dissociation have been studied
in the previous chapter by reference to the case of mercuric oxide;
the subject will now be further illustrated by a few chosen
examples.

Nitrogen Peroxide. Nitrogen peroxide exists at ordinary
atmospheric temperature as a brown gas which becomes deeper in
colour when heated, but paler when cooled, yielding at 10 C. a
yellow liquid which at 10 *C. forms colourless crystals. The
alteration in colour of the gas is due to the following reversible

change:

N 2 O 4 ^ 2NO 2 ,

N0 2 being deep brown, while N 2 O 4 is pale. Thus one molecule of
N 2 4 is dissociated by heat into two similar molecules of N0 2 . The

CHEMICAL CHANGES CLASSIFIED 275

progress of dissociation is indicated by an increase of pressure of
the gas at constant volume, or by a reciprocal decrease in density
at constant pressure after a correction has been made in either case
for change of temperature. Thus the density of nitrogen peroxide
gas at 4-2 C. is 2-588 (air = 1) to which a molecular weight of
74-8 corresponds, whilst its density at 97-5 C. is 1-783, with a
corresponding molecular weight of 51-5. These calculated mole-
cular weights show that 62-6 and 11-9 per cent respectively of
the mixed gases is N 2 O 4 .

This is the simplest example of dissociation that can be given,
for the whole system is gaseous, and there is only one kind of dis-
sociation product. The formation of N 2 4 from two molecules of
NO 2 may be regarded as a case of polymerization.

Phosphorus Pentachloride. The vapour of phosphorus penta-
chloride readily dissociates thus:

PC1 5 ^ PC1 3 + C1 2 ,

the dissociation at various temperatures being indicated by the fol-
lowing densities, reduced to normal temperature and pressure:

Temperature C 182 200 250 300 336

Density 73-3 70-0 57-6 52-4 52-5

The normal density of PC1 6 is 104-2; thus dissociation, considerable
at 182 C. is complete at 300 C.

Now, suppose that instead of vaporizing into empty space,
PC1 5 is made to vaporize into a space containing much vapour of
PC1 3 . Every C1 2 molecule derived from PC1 5 by incipient dissocia-
tion would then encounter many molecules of PC1 3 , and its chance
of remaining uncombined with one of these surrounding molecules
would be very small. Thus PC1 3 molecules would hinder the dis-
sociation of PC1 5 . This has been found to be the case. Wiirtz
vaporized PC1 5 into PC1 3 vapour and found the PC1 6 vapour density
at 160 to 175 C. to be nearly 104-2, that of the undissociated
compound.

This is a notable case of the influence of mass on chemical
change. Vaporization of PC1 5 into chlorine gas would produce a
similar result.

Ammonium Chloride.

This salt, when vaporized, dissociates into NH 3 and HC1, thus:
NH 4 C1 ^ NH. + HC1.

276 CHEMICAL THEORY

The dissociation may be proved by the separation of ammonia from
hydrogen chloride by the more rapid diffusion of the lighter am-
monia through a porous pipe stem.

Mercurous Chloride. The vapour density of mercurous chloride
ordinarily corresponds to the molecular weight of HgCl, but it
may be shown that the vapour is a mixture of Hg and HgCl 2 ,
since the mercury will diffuse through a porous tube and condense
in globules outside. Consequently the formula for mercurous
chloride is Hg 2 01 2 , and the vapour dissociates thus:

Hg 2 Cl a ^ Hg + HgCl 2 .

It has been shown, however, by Baker that when water is very
carefully excluded the vaporized salt does not dissociate, but shows
a density corresponding to the formula Hg 2 Cl 2 .

Calcium Carbonate.

This compound dissociates thus:

CaCO 3 ^ CaO + CO 2 ,

and the dissociation pressures of this salt at various temperatures
are these:

t C. Pressures.

547 27 mm. Hg.

610 46

740 255

810 678

865 1333

Dissociation ceases at any given temperature when the pressure
reaches the corresponding value; consequently the accumulation of
the gaseous product of dissociation hinders dissociation. This is a
principle of mass action which is true in general. Its application
in the case under consideration may be seen in this way. If pre-
cipitated chalk is to be converted into quicklime by ignition in
a crucible, whilst it is desirable to have as high a temperature
as possible, the lid must not be allowed to cover the crucible too
tightly or carbon dioxide, will accumulate within it and hinder
dissociation. This example of dissociation differs from the fore-
going, because one of the dissociation products is permanently
solid.

CHEMICAL CHANGES CLASSIFIED 277

Barium Peroxide.

Barium peroxide furnishes an important example of dissociation.
When heated to a dull red heat it yields oxygen, thus:
2Ba0 2 =?=: BaO + O 2 ,

and the resulting BaO will reabsorb oxygen at a lower temperature
than that at which the peroxide dissociates. If, however, the
material is kept at a constant temperature of 700 C., the direction
of the reaction may be reversed by altering the pressure. In con-
tact with air at 2 atmospheres pressure BaO at 700 C. absorbs
oxygen forming BaO 2 ; but when the air is pumped away the re-
action is reversed, and oxygen is evolved from the BaO 2 . This
is the principle of Erin's process for obtaining oxygen from the
air, and as the direction of the reaction depends upon the concen-
tration of the reacting substances, it is a good example of the action
of mass.

Cry stallo- hydrates. The numerous salts and other -substances
containing water of crystallization afford material for illustrating
the phenomena of dissociation. Water of crystallization is often
but loosely retained by a salt, and it may sometimes be lost even
at atmospheric temperature.

The student will now appreciate the fact that the condition of
a salt with reference to water of crystallization depends not only
on temperature, but also on the concentration of water vapour in
the vicinity of the salt. There is a certain range of temperature
and external pressure of water vapour within which a given salt
or its hydrate can exist permanently. Outside this range the com-
pound will give rise to a product containing a different proportion
of combined water. Lowering of temperature and increase of ex-
ternal water- vapour pressure will promote the formation of a higher
hydrate; raising of temperature and diminution of external vapour
pressure will cause the salt to lose water, so as to form a lower
hydrate or become anhydrous.

Thus some hydrated salts are permanent in air because of
favourable temperature and water -vapour pressure, while others
tend to lose water or absorb it from the air.

Crystals of blue vitriol, CuSO 4 -5H 2 O, for example, neither lose
nor gain water under ordinary atmospheric conditions, but crystals
of washing-soda lose water according to the scheme:

Na 2 CO s -10H 2 O ^ Na 2 CO 8 -H 2 O + 9H 2 O,

27d CHEMICAL THEORY

because the vapour pressure of the system consisting of the deca-
and mono-hydrate is greater than the pressure of water vapour in
the air at ordinary temperature. So the salt effloresces, becoming
opaque on the surface, owing to the formation of the powdery
monohydrate, efflorescence being the phenomenon in which hydrated
crystals lose water vapour to the air.

Hydrated salts having small vapour pressures are frequently
very soluble in water, and their saturated solutions may have
vapour pressures less than the vapour pressure of water vapour
in the air. Atmospheric water vapour will be absorbed by a
salt of this kind until a liquid solution is produced of such a
strength as to have a water-vapour pressure equal to that in the
superincumbent atmosphere. Then the salt is said to 'deliquesce.
Crystallized calcium chloride, CaCl 2 6H 2 O, is a good example of
a deliquescent salt.

Thermal Decomposition.

It has been seen in the previous chapter that whilst some
reactions caused by heat are reversible, being cases of dissociation,
others, such as the decomposition of complex organic substances,
are undoubtedly irreversible. The question now arises as to the
extent and boundaries of reversible and irreversible thermal change,
and whether many changes which at present appear irreversible
are really so. The investigation of this question leads to some
interesting results, but not to a definite conclusion in every case.

The thermal decomposition of oxides, hydroxides, and oxysalts
furnishes a sufficiently wide field for investigation.

Thermal Decomposition of Oxides.

The student has already learned that the thermal decomposition
of mercuric oxide is reversible; the action is one of dissociation,

thus:

2HgO ^ 2Hg + O 2 .

Silver and auric oxides are similarly decomposed by heat:

2Ag 2 O = 4Ag + O 2 ,
2Au 2 O 3 =

but these reactions are not reversible; gold and silver are not sus-
ceptible of atmospheric oxidation; they are too electro-negative, too
inert chemically, for that. Silver, however, is oxidized by ozone.

CHEMICAL CHANGES CLASSIFIED 279

Cupric oxide is an interesting case; when heated to a white
heat it loses half its oxygen, the following reaction being reversible:
4CuO ^: 2Cu 2 O + O 2 .

No oxides of metals more electro-positive than mercury lose all
their oxygen when heated,

The loss of some of their oxygen by higher oxides under the
influence of heat furnishes an instructive series of phenomena.
Consider, for example, the oxides N 2 O 5 , P 2 5 , As 2 5 , Sb 2 O 5 , Bi 2 O 5 ,
which are decomposed by heat, thus:

2N 2 O 6 4NO 2 + O 2 .

[P 2 O 6 stable.

l2P 2 O 6 P 4 O fl + O 2 .

2As 2 O 6 As 4 O 6 + 2O 2 .

2Sb 2 O 6 2Sb 2 O 4 + O 2 .

Bi 2 O 6 Bi 2 O 3

None of these reactions appears to be reversible. The behaviour
of N 2 O 5 is unique, like that of the element nitrogen itself. P 2 6 is
easily produced by the atmospheric oxidation of the element or the
lower oxide; but < arsenic is less oxidizable than phosphorus, and
the higher oxide must be formed by the chemical oxidation of the
lower oxide, which alone is produced when the element burns. The
oxide Bi 2 O 5 is a peroxide, formed only in the wet way; it has feebly
acidic properties, but easily loses > oxygen. Thus the oxides P 2 6 ,
As 2 5 , Sb 2 O 6 , Bi 2 O 5) with the possible exception of Sb 2 O 6 , stand in
the order of decreasing stability.

The dioxides of the fourth group CO 2 , SiO 2 , GeO 2 , SnO 2 , PbO 2>
are similarly related as regards stability; Pb0 2 alone is decomposed
by heat. Or consider the superoxides Na 2 2 , Ba0 2 , SrO 2 , CaO 2 ,
Mg0 2 , &c. Sodium peroxide does not lose oxygen when heated
short of very high temperatures; Ba0 2 ,and the oxides which follow
it, are decomposed with increasing readiness as electro-positiveness
diminishes. Only the more electro -positive metals form super-
oxides, i.e. derivativestof hydrogen peroxide, at all.

The reversible reaction,

6PbO + O 2 =s=: 2Pb 3 O 4 ,

is instructive, and it is noteworthy that PbO 2 is not formed by the
atmospheric oxidation of PbO; Pb 8 O 4 is a compound of 2PbO
and PbO 2 , that is, it is a salt, and the two oxides being mutually

280 CHEMICAL THEOEY

satisfied, the PbO in Pb 3 O 4 is not free to combine with more
oxygen.

If, however, another base is present with which Pb0 2 can
combine, PbO may be completely oxidized to PbO 2 by atmospheric
oxygen. Thus the following reaction is reversible:

2PbO + O 2 + 2Na 2 CO 3 ^=r 2Na 2 PbO 3 + 2CO 2 .

This is also a good example of mass action, since preponderance of
oxygen or carbon dioxide determines the direction of the reaction.
The case presented by the following reaction is similar:

4CrO 3 2Cr 2 O s + 3O 2 .

The reverse reaction takes place only in the presence of alkali,
when atmospheric oxygen may be absorbed at high temperature to
form chromate, thus:

4 Na 2 CrO 4 + 4 CO 2 .

Amongst non-metallic oxides the case of sulphur trioxide is
important. It is well known that the reaction

2SO 2 + O 2 > 2S0 3

takes place to any considerable extent only in presence of a
catalyst, whilst the reverse reaction,

2SO 3

occurs when the vapour of the trioxide is passed through a tube
heated to 1000 C. Nevertheless, when the dioxide is heated to
1200 C., or is submitted to the action of electric sparks, it undergoes
the following reversible change:

3SO 2 ^=f: 2SO 3 + S.

Thermal Decomposition of Hydroxides

Hydroxides may be basic or acidic, and their stability depends
on the basic or acidic intensity of the corresponding oxides.
Consider, for example, the basic hydroxides:

NaOH, Ba(OH) 2 , Ca(OH) 2 , MgCOH)* A1(OH) 3 , Cu(OH) 2 , AgOH, NH 4 OH.

NaOH and Ba(OH) 2 , producible from the corresponding oxides
by violent reactions with water, cannot be dehydrated by heat
alone; Ca(OH) 2 , produced also from the oxide by slaking, is easily
dehydrated by heating above 150 C.; Mg(OH) 2 is formed by very

CHEMICAL CHANGES CLASSIFIED 281

slow combination of MgO with water, but more readily by precipita-
tion, and is stable at 100 C.; A1(OH) 3 , formed by precipitation in the
cold, loses water, forming, A1 2 O(OH) 4 or A1OOH, when heated at
100 0. or dried over sulphuric acid; Cu(OH) 2 is formed as a blue
precipitate in the cold, which turns dark brown when warmed
with water, forming Cu 4 O 3 (OH) 2 ; AgOH exists only in very dilute
aqueous solution, for the precipitate formed by adding alkali to a
silver salt is Ag 2 O, which dissolves very slightly in water, producing
a faintly alkaline solution containing AgOH.

Ammonium hydroxide. NH 4 OH, is unique. A hydrate,
NH 3 -H 2 O exists in crystals, melting at 78 C., and also in solu-
tion. According to modern views 1 this must be a covalent compound

H
H:N:H:O:H;

but whether it ionizes thus: NH 4 OH ^^= NH 4 * + OH', causing
alkalinity, or whether alkalinity is due to the more direct reaction
NH 3 + H- + OH ^^ NH 4 * + OH', it is perhaps impossible to
say. The presence both of NH 3 and of hydroxide ions is proved
by the use of ammonia solution to form an anmiine, such as
CuSO 4 -4NH 3 -H 2 O, as well as to precipitate a base such as Fe(OH) 3 .
When an aqueous solution of ammonia is boiled, however, all the
ammonia escapes as gas.

The following acidic hydroxides 2 (oxy-acids) may be considered:
S0 2 (OH) 2 , PO(OH) 3 , SO(OH) 2 , CO(OH) 2 .

Sulphuric acid, SO 2 (OH) 2) is fairly stable towards heat and
may be distilled, for SO 3 is a powerfully acidic oxide, and as such holds
its combined water tenaciously. At 440 C., however, dissociation
according to the scheme: H 2 SO 4 + S0 3 + H 2 O is complete.

Phosphoric acid, PO(OH) 3 , loses water when heated, thus:
2PO(OH) 3 PA(OH) 4 PA(OH) 2

Ortho- Pyro- Meta-phosphoric acid;

but it is noteworthy that the meta-aeid is the final dehydration
product, from which water cannot be removed. The meta- and the
pyro-acid both revert to the ortho-acid in contact with water,
and there is evidence that the meta-acid passes through the pyro
form on the way back to the ortho-acid.

Sulphurous acid, SO(OH) 2 or SO 2 : H(OH), shares with carbonic

1 Vide Lowry, Cltemittry and Industry, 1928, 1233.

2 It may be questioned, however, if oxy-acids should be regarded as hydroxides; but see
p. 287.

282 CHEMICAL THEORY

and nitrous acids, which also have gaseous anhydrides, the property
of existing only in aqueous solution; several crystallo-hydrates of
this acid exist however. The following reaction is reversible:

SO 2 + H 2 O ^ SO(OH) 2 ,

though, like ammonia, some sulphur dioxide exists in aqueous
solution without being hydroxylated. It may be that sulphurous
acid in aqueous solution passes into the unsymmetrical form, thus:

SO(OH) 2

though modern theory seems to make this view superfluous.

Carbonic acid, CO(OH) 2 , is formed when carbon dioxide dis-
solves in water, thus:

CO 2 + H 2 O ^=- CO(OH) 2 ;

the carbonic acid so formed being ionized to a considerable extent,

thus:

H 2 CO 3 ^f H- + HCO 3 '.

When a solution of carbonic acid is boiled, both these reactions are
reversed, and all the carbon dioxide escapes from the liquid.

Thermal Decomposition of Oxy-Salts.

The usual mode of thermal decomposition of oxy-salts is into
basic and acidic oxides or their decomposition products; e.g.:

CaC0 3 CaO + CO 2 .

2Cu(NO 3 ) 2 -

Whether such a reaction occurs depends upon the relative basic
and acidic strengths of the respective oxides, and upon the volatility
of the acidic oxide.

Thus, when both basic and acidic oxides are powerful, as, for
example, in the case of sodium sulphate, thermal decomposition
does not take place, and even when the acidic oxide is feeble an
alkali oxide is sufficiently powerful to retain it, as in the case
of sodium carbonate, which is not .decomposed at 1000 C. The
sulphates of the feebler basic oxides lose sulphur trioxide when
heated; for example, copper sulphate forms a basic salt at a dull
red heat, and ferrous sulphate loses all its SO 3 , leaving a residue of
ferric oxide.

The behaviour of nitrates and chlorates when heated is instruc-

CHEMICAL CHANGES CLASSIFIED 283

tive. Those of the alkali rnetals do not lose their nuclear nitrogen
or chlorine; thus potassium nitrate forms nitrite, and chlorate
chloride. Nitrates and chlorates of feebler metals, however, leave
a residue of oxide; e.g. the lead salts decompose thus:

2Pb(N0 3 ) 2 2PbO + 4NO 2 + 2 .
2Pb(C10 3 ) 2 2PbO + 2Cl 2 + 5O 2 .

An interesting question arises here as to reversibility. The
reaction:

2KC10 3 2KC1 + 30 2

appears not to be reversible. Why is this? Hypochlorite easily
undergoes self -oxidation and reduction thus:

3KOC1 -* 2KC1 + KC10 3 ;

yet chloride seems incapable of oxidation to chlorate. The explana-
tion is to be found in the fact that the chloride ion is more stable
than the ion of any oxyacid of chlorine, and that the ions of the
oxy acids stand in the following order of increasing stability: hypo-
chlorite, chlorate, perchlorate. These facts are illustrated by the
formulae attributed to these ions according to the electronic theory
of valency (q.v.).

Besides relative basic and acidic strengths of oxides in an oxy-
salt, the volatility of the acidic oxide determines the stability of
the salt. Silica, for example, is not volatile like carbon dioxide,
and so calcium silicate, unlike calcium carbonate, is not decomposed
by heat. For the same reason silica at high temperature displaces
carbon dioxide from chalk, thus:
CaC0 3

a reaction common in metallurgy.

Phosphates, moreover, are not decomposed by heat; for example,
bone ash, Ca 3 (PO 4 ) 2 , is unchanged at a white heat; for, although
phosphoric oxide itself is volatile, it cannot be separated from a
combined basic oxide any more than it can be separated from
water combined with it in metaphosphoric acid, H 2 P 2 O6-

In general, the mode of decomposition of an oxy-compound
when heated reveals the relative stabilities at high temperatures
of the possible decomposition products.

Thus the isomorphous salts KC1O 4 and KMn0 4 behave very
differently when heated, because in spite of the superficial resem-
blance between them and the similarity of their constitution, the
elements concerned, namely chlorine and manganese, are widely

284 CHEMICAL THEOKY

different in chemical nature. The following reactions represent
the manner of decomposition of these salts:

KC1O 4 = KC1 + 2O 2 .

2 KMnO 4 = K 2 MnO 4 + Mn0 2 + O 2 .

The chief ammonium oxy-salts are decomposed by heat thus:

/(NH 4 ) 2 SO 4 = NH 4 HSO 4 + NH 3 .

\NH 4 HS0 4 = H 2 + S0 3 + NH 3 .

NH 4 NO 3 = 2H 2 O + N 2 0.

NH 4 NO 2 = H 2 + N*

(NH 4 ) 2 HP0 4 = HPO 3 + H 2 + 2 NH 3 .

/(NH 4 ) 2 C0 3 = NH 4 HC0 3 + NH 3 .

\NH 4 HC0 3 = H 2 + C0 2 + NH 3 .

Chemical Interaction of Water with Elements and Compounds.

The interaction of elements and pure water is limited almost
completely to the behaviour of a few electro-positive metals.

The metals of the alkalis and alkaline earths decompose water
at atmospheric temperature, displacing hydrogen with formation
of the hydroxide of the metal. The vigour of the reaction with
water increases with the electro-positiveness of the metal from
lithium to caesium, and from calcium to barium. The reaction is
attributable firstly to the ionization of water, although this is so
small, and secondly, to the solution pressure of the metal which
is superior to that of hydrogen.

The student has met with this latter idea before, under the
subject of electrolysis (q.v.). Thus metallic sodium, striving to
assume the ionic state when brought into contact with water,
displaces the hydrogen ions of the latter, causing them to lose their
electric charges and escape as gas.

Besides the metals of the alkalis and alkaline earths, amalga-
mated aluminium reacts with water at atmospheric temperature
evolving hydrogen, and powdered aluminium as well as magnesium
decomposes steam. The rusting of iron is believed to be due, first
of all, to the action on water of the metal containing impurities
which set up slight electro-potential differences throughout the
mass; and it is well known that steam is decomposed by red-hot
iron, whilst hydrogen can reduce heated iron oxide to metal.
Copper, however, has no action on water or steam, for it is less
electro-positive than hydrogen and does not displace this element
from water or dilute acids.

CHEMICAL CHANGES CLASSIFIED 285

So the following reactions are instructive:
H 2 O -> MgO + H 2 .

ii. 3Fe + 4H 2 O :^ Fe 3 O 4
iii. Cu + H 2 O - CuO + H 2 .

Reactions (i) and (iii) proceed in single and opposite directions;
reaction (ii) is reversible because of the intermediate character of
iron.

Of the non-metals the halogens alone react with water.
Fluorine decomposes water in the dark, and chlorine in the day-
light, both with evolution of oxygen. It is probable that the
following is the first reaction between chlorine and water:

C1 2 + H 2 O ^= HC1 + HOC1;

for chalk and chlorine water yield calcium chloride and hypo-
chlorous acid, the latter not reacting with chalk.

Hydrolysis.

The chemical decomposition of compounds by water is hydro-
lysis. This, again, is to be attributed to the hydrogen and
hydroxide ions existing in water.

Ideal salts, consisting of powerfully basic and acidic ions, are
not hydrolyzed by water; their aqueous solutions are ionized
but remain neutral.

Normal salts are not necessarily neutral. Sodium chloride and
trisodium phosphate, Na s PO 4 , are both normal salts, but whilst
NaCl is also a neutral salt in the sense of giving a neutral solution
with water, this is not the case with Na 3 PO 4 , the solution of which
is alkaline on account of hydrolysis with production of free alkali,
i.e. OH' ions in solution, thus:

Na 3 PO 4 + HOH ^ Na 2 HP0 4 + NaOH.
NaOH ^=e Na- + OH'.

Even the hydrogen salt Na 2 HP0 4 yields an alkaline solution
owing to hydrolysis, thus:

Na 2 HPO 4 + HOH ^ NaH 2 PO 4 + NaOH;

for although the salt NaH 2 PO 4 is itself acid, it does not produce
such a concentration of H* ions as the equivalent of NaOH pro-
duces of OH' ions.

286 CHEMICAL THEORY

The reactions of the three sodium phosphates are:

Na 3 PO 4 , strongly alkaline;
Na 2 HPO 4 , alkaline;
NaH 2 PO 4 , acid;

the point of neutrality thus lying between Na 2 HP0 4 and NaH 2 P0 4 .

An alternative way of explaining the alkalinity of Na s P0 4>
applicable to the effect of mixing equivalents of NaOH and H S P0 4 ,
is that NaOH being stronger as a base than H 3 PO 4 is as an acid,
the alkali provides a larger proportion of OH' ions than the acid
of H' ions, and so an excess of OH' ions remains over after all
the H" ions from the acid have been neutralized.

These considerations have an important bearing on the volu-
metric estimation of acids and alkalis. Only acids and alkalis
which produce salts not appreciably hydrolyzed in aqueous solution
can be titrated in the ordinary way; obviously phosphoric acid
cannot be directly titrated with an equivalent standard alkali.

As regards salt hydrolysis in general, whilst the salts of strong
bases with weak acids are alkaline in reaction, those of weak bases
with strong acids are acid. An example of the latter is furnished
by ferric chloride which is hydrolyzed in aqueous solution with
formation of a basic salt and free acid, somewhat as follows:

FeCl 3 + HOH ^f FeOHCl 2 + H' + Cl'.

The formation of the basic salt is shown by the darkening of the
solution. The same thing is especially noticeable when ferric alum
is dissolved in water. The pale-violet crystals yield a brown
solution which becomes colourless when a little sulphuric acid is
added to convert the basic salt into the normal salt. If a solution
of ferric alum is poured into much boiling water a precipitate of
ferric hydroxide separates owing to complete hydrolysis. Ferric
acetate is similarly hydrolyzed by boiling water with the precipi-
tation of a basic acetate.

That salt hydrolysis is a frequent phenomenon is shown by
the following facts:

Soluble salts of the following acidic radicles have alkaline
reactions in solution:

Borate, carbonate, chromate, cyanide, hypochlorite, nitrite,
phosphate, silicate, sulphide, sulphite.

Soluble salts of the following metallic radicles possess an acid
reaction:

CHEMICAL CHANGES CLASSIFIED 287

Mercurous, mercuric, cupric, aluminium, chromic, ferrous, ferric,
stannous, stannic, antimonious, bismuthous.

Salts of weak bases with weak acids are generally insoluble
in water, and on that account less amenable to hydrolysis.

A fat, may be hydrolyzed by superheated steam with the pro-
duction of glycerol and fatty acid, thus:

_ C 3 H 5 X 3 + 3 HOH = C 3 H 6 (OH) 3 + 3HX.
(C 3 H 6 ) = glyceryl.

X = C 15 H 31 COO.(palmitate),orC l7 H 36 COO.(stearate).

The catalytic effect of H" ions in promoting the hydrolysis of
an ester or of cane sugar has already been noticed (p. 267).

Sometimes the basic salt produced by hydrolysis is insoluble
and is precipitated. Thus, bismuthous and antimonious chlorides
are not only hydrolyzed by water but yield precipitates of the
basic chlorides:

BiCl 3 + H 2 O ^=5 BiOCl + 2HCl.
SbCl 3 + H 2 ^= SbOCl + 2HCL

On this account hydrolysis is accentuated, for the reversal of
the reaction is greatly hindered by the separation of the hydrolytic
product in the solid state.

Incidentally it may be observed that whilst BiOCl is stable
towards water, SbOCl, on account of the feebler basic properties
of antimonious oxide, loses all its chlorine when boiled with water.

So far the hydrolysis of salts has been considered; but the
process is not confined to salts.

Consider the series of chlorides:

NaCl, MgCl 2 , A1C1 3 , SiCl 4 , PC1 5 .

The series begins with a powerfully metallic chloride, and ends
with a non-metallic chloride. The chloride of a powerful metal,
being a true salt, is not hydrolyzed by water, whilst the chloride
of a non-metal is at once and completely hydrolyzed. Between
these two extremes there is a gradation of hydrolysis.

MgCl 2 gives a neutral solution with water, but when this
solution is evaporated nearly to dryness hydrogen chloride gas
escapes with the steam, and the resulting pasty mass, containing
the basic salt Mg(OH)Cl, reacts alkaline. By ignition the oxide
MgO is eventually formed.

AlCli dissolves in water, reacting vigorously with it if anhydrous,

288 CHEMICAL THEORY

and produces an acid solution by incipient hydrolysis; when this
solution is evaporated to dryness and ignited, the oxide A1 2 O 3
remains.

SiCl^ is instantly decomposed by water with separation of
gelatinous silica H 2 Si0 3 :

H 2 SiO s

To a certain extent the reaction is reversible, for H 2 Si0 8 is more
soluble in hydrochloric acid than in water.

PCls is similarly decomposed, liquid POC1 3 being first formed,
and then H 8 P0 4 thus:

PC1 6 + H 2 O POC1 3 + 2HC1.
POC1 3 + 3H 2 O PO(OH) 3 + 3HC1.

This is an extreme case of hydrolysis, for the reaction is not
reversible.

Chemical Interaction of Acids and Metals.

It has been seen that there is an interaction between water and
some metals, and that this is believed to be due to the slight ioniza-
tion of water without which this substance would be inert. Now
acids owe their essential nature to the presence of hydrogen ions,
the strength of an acid in aqueous solution beiug measured by the
concentration of these ions. Consequently acids, containing much
higher concentration of hydrogen ions, behave much more vigorously
toward metals than water.

Consider the following metals arranged in order of decreasing
electro-positi veness :

Cs, Rb, K, Na, Ba, Sr, Ca, Mg, Al, Mn, Zn, Cd, Tl, Fe, Co, Ni, Sn, Pb,

Sb, Bi, As, Cu, Hg, Ag, Pd, Pt, Au.

Metals preceding hydrogen can displace it as a gas from dilute
hydrochloric or sulphuric acid; those that follow H do not generate
this gas in contact with dilute acid. Or, using an idea the student
is already familiar with: metals preceding hydrogen have a greater,
metals following it a less, solution pressure than this element. It
has already been seen that metals as far as calcium, and under
some circumstances aluminium, displace hydrogen from cold water,
and that if steam is used and the metal is heated reactivity
extends as far as iron.

CHEMICAL CHANGES CLASSIFIED 289

The extent to which a metal displaces hydrogen from a dilute
acid depends upon:

i. The nature and purity of the metal.

ii. The nature and state of dilution of the acid.

i. The nature and purity of the metal. If dilute hydrochloric
acid is poured upon fragments of magnesium, zinc, iron, and tin,
in separate test-tubes, the reaction will be very vigorous with
magnesium, less so with zinc, still less so with iron, and very slight
with tin; indeed, it is necessary to heat tin with moderately con-
centrated hydrochloric acid to dissolve the metal with evolution
of hydrogen. All this is quite in accord with the order of electro-
potential of the metals.

The purity of a metal has a marked effect on its interaction with
acid. If a stick of pure zinc is placed in dilute hydrochloric acid,
very slow action takes place; but if a length of fine copper wire is
coiled round the zinc, action becomes vigorous and the hydrogen is
seen to be coming off from the surface of the copper, although this
metal is found unchanged when the zinc has been dissolved. These
different effects are due to the fact that in the first case the displaced
hydrogen forms a protective film on the surface of the zinc, thus
polarizing it, whilst in the second case electrical action is set up,
hydrogen flows with the electric current from the zinc to the copper,
and is evolved from the surface of the latter metal, whilst the zinc
is left continuously exposed to the action of the acid. A similar
effect is produced by adding a few drops of copper sulphate or
platinic chloride to the acid in which the zinc is immersed. Copper
or platinum is deposited on the zinc, and causes the production of
electric circuits with the consequent promotion of chemical action.
If the zinc is impure, and contains a small proportion of a less
electro-positive metal, solution is hastened by this impurity without
the help of an added metal. Such action may be regarded as catalytic.

ii. The nature and state of dilution of the acid. The strength
of an acid, i.e. its degree of ionization, determines the rate at which
a metal displaces hydrogen from it. Thus, if parallel experiments
are done with zinc in (a) dilute hydrochloric acid, (6) dilute acetic
acid of equivalent strength, the rate of evolution of hydrogen from
the acetic acid will be exceedingly slow, whilst that from the hydro-
chloric acid may be vigorous. The explanation is simple. Hydro-
chloric acid is a strong acid, almost completely ionized in dilute

(DCO) 20

290 CHEMICAL THEORY

solution, whilst acetic acid is a weak acid which is but slightly
ionized, and therefore approaches water in its behaviour towards
metals.

The state of dilution of a particular acid also determines the rate
of evolution of hydrogen. If an acid is much diluted the concentra-
tion of hydrogen ions is correspondingly reduced, together with the
vigour of the reaction. If, however, too little water is present there
may not be room enough in the solvent for extensive ionization,
and consequently displacement of hydrogen by a metal will not be
effected.

This is not the case with concentrated hydrochloric acid, which
contains only about 32 per cent of hydrogen chloride; but it is the
case with concentrated sulphuric acid, which consists of about 98 per
cent of the absolute acid and 2 per cent of water. Thus zinc does
not displace hydrogen from concentrated sulphuric acid; a reaction
commences when this acid is heated with the metal; the gas evolved,
however, is not hydrogen but sulphur dioxide, thus:
Zn + 2 H 2 SO 4 = ZnSO 4 + 2 H 2 O + SO 2 .

Since a similar reaction takes place with copper, it is not necessary
to assume that free hydrogen has anything to do with it.

Interaction of Nitric Acid and Metals.

Since nitric acid contains much oxygen, and is easily reducible
the question arises whether hydrogen can escape from the acid when
displaced from it by a metal, or whether the hydrogen will neces-
sarily reduce the acid instead. If the acid is reduced, since there
are various reduction products of nitric acid, the possible reactions
of the acid with metals are various.

Hydrogen is among the gases evolved from dilute nitric acid by
magnesium, the most electro-positive of the metals which do not
react with cold water; but with this exception hydrogen is not
obtained.

The possible reduction products of nitric acid may be shown by
separating water from the acid, and breaking up the resulting
anhydride thus:

2HN0 3 = H 2 + N 2 6 ,

= H 2 O + 2N0 2 + 0,
= H 2 O + N 2 3 + 2O,
= H 2 O + 2NO + 3O,

CHEMICAL CHANGES CLASSIFIED 291

Besides the five reduction products, N0 2 , N 2 O 8 , NO, N 2 O, N 2 ,
hydroxylamine, NH 2 OH, and ammonia, NH 3 , are sometimes pro-
duced.

The question now arises whether displaced hydrogen or the
metal itself reduces the acid. Hydrogen cannot reduce nitric acid
in the case of a metal which does not displace this element from an
acid. It can scarcely be assumed that nascent hydrogen is respon-
sible for the reduction of nitric acid by copper, since this metal does
not displace hydrogen from any acid. As a matter of fact, metals
may be divided into two categories as regards their behaviour to-
wards nitric acid. The metals zinc, cadmium, iron, tin, and others
more electro-positive than hydrogen may reduce nitric acid as far
as ammonia; the metals bismuth, copper, mercury, silver, which are
less electro-positive than hydrogen, do not reduce the acid beyond
the stage of nitric oxide.

Presumably the reduction to nitric oxide does not necessitate the
intervention of hydrogen; reduction beyond this stage may. Dis-
placed hydrogen must evidently play a part in the formation of
NH 2 OH and NH 3 ; it may be indirectly the cause of the evolution
of nitrogen, which may be derived from ammonium nitrite thus:

NH 4 N0 2 = N 2 + 2H 2 0,

whilst N 2 O may come from hyponitrous acid thus:
HON : NOH = N 2 O + H 2 O.

With regard to the reduction -of nitric acid by metals of the
copper series there is evidence that the pure, diluted acid, free from
nitrous acid, does not react with these metals, but that nitrous acid
is necessary to start the reaction, which proceeds thus:

i. Cu + 4HNO 2 = Cu(NO 2 ) 2 + 2H 2 O + 2NO;

ii. Cu(NO 2 ) 2 + 2HNO 3 = Cu(NO 3 ) 2 + 2 HNO 2 ;
iii. HNO 3 + 2 NO + H 2 O = 3 HNO 2 ;

and adding

iv. Cu + 3 HNO 3 = Cu(N0 3 ) 2 + HN0 2 + H 2 O.

So, by the aid of a little HN0 2 more is produced, which, instead of
accumulating, decomposes thus:

v. 3 HNO 2 = HNO 3 + H 2 O + 2 NO.

Thus nitric oxide gas is evolved, and the nitrate of the metal is
formed in solution.

292 CHEMICAL THEORY

This theory of the activity of nitrous acid does not invalidate
the usual equation:

3 Cu + 8 HNO 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO,

which may be obtained by multiplying equation (iv) by three, and
adding the product to equation (v).

The concentration of the nitric acid employed affects the nature
of the reduction products. This may be seen by pouring very dilute
nitric acid on zinc in a test-tube, and then gradually increasing the
strength of the acid. At first no gas is evolved, then a colourless
gas, and afterwards a brown gas.

With the most dilute acid ammonia is produced, and combines
with the acid, forming ammonium nitrate; then, with increasing
strength of acid, nitrogen and nitrous oxide appear, and afterwards
nitric oxide, and possibly still higher oxides of nitrogen with the
strongest acid. In general, the more dilute the acid the more
perfectly it is reduced.

The following equations represent the formation of nitrous acid
and ammonia respectively when zinc and nitric acid interact,

4 Zn + 10 HNO 3 = 4 Zn(NO 3 ) 2 + 5 H 2 O + N 2 O.

10HNO 3

That the quantities on the left side of the equation are the
same in both cases is interesting. To argue that the same products
should therefore result would be fallacious, for chemical equations
do not represent the concentrations in which reacting substances
are brought together.

Effect of Solubility and Volatility on Chemical Change.

There are two laws associated with the name of Berthollet
which state (i) that if two substances by reacting in solution can
yield a product less soluble than themselves, that product will be
formed; and (ii) that if two substances when heated together can
yield a product more volatile than themselves, that product will
result. The facts thus expressed find ready explanation according
to the principles of mass action.

The direction of the reaction,

AX + BY ^ AY + BX,

will depend on the effective concentration of each component; if
one component, from any cause, is placed partly or wholly outside

CHEMICAL CHANGES CLASSIFIED 293

the sphere of action, its power of directing the course of the reaction
will be so far diminished. Consider, for example, the following
two reactions placed at the beginning of this chapter:

BaSO 4 + 2HCL

Superficially these reactions appear to be alike; actually they
differ very much. The first is supposed to take place in dilute
solution, BaSO 4 being precipitated and HC1 remaining in solution;
the second takes place when solid NaCl and concentrated H 2 SO 4
are heated together, solid Na 2 SO 4 resulting, while HC1 escapes.

If the conditions are reversed, the results are not so effective.
Thus if solid BaCl 2 is heated with concentrated sulphuric acid the
reaction will be only partial, because the BaCl 2 will be protected
from the acid by a crust of BaS0 4 ; and if NaCl and H 2 S0 4 are
mixed in dilute aqueous solution no HC1 gas escapes.

The first reaction ordinarily proceeds to finality because of the
insolubility of BaSO 4 . When BaCl 2 and H 2 S0 4 interact in aqueous
solution, BaS0 4 and HC1 are first formed in solution, and if both
remained in this state they would be effective in establishing an
equilibrium with the original substance, and no change would be
apparent.

But BaS0 4 is at once precipitated because it is so very slightly
soluble in water, and only the very minute quantity of this salt
remaining in solution can have any effect in reversing the change.
This effect is quite negligible, and, practically speaking, the reaction
proceeds to completion from left to right.

The second reaction proceeds to finality because of the volatility
of hydrogen chloride. When NaCl is heated with concentrated
H 2 S0 4 , torrents of HC1 gas are evolved, and the reversal of the
reaction by means of this gaseous product which has escaped is
out of the question.

The principles here illustrated are very far-reaching, and the
student should make sure he grasps them,

Consider another example: The preparation of nitre by inter-
action of sodium nitrate and potassium chloride in aqueous solution,

thus:

NaNO 3 + KCl ==-

All the four possible salts are soluble in water, but sodium chloride
is the least soluble of the four. When, therefore, sodium nitrate

294 CHEMICAL THEORY

and potassium chloride are mixed in equivalent proportions in
solution and the latter is evaporated, sodium chloride is the first
salt to crystallize; and thus potassium nitrate remains in solution,
and may be crystallized after the sodium chloride has been re-
moved.

A salt such as sodium chloride may be eliminated more effectually
if it is formed in contact with a non-aqueous solvent in which
it is not soluble. An interesting example of this is furnished by
the preparation of free hydroxylamine, NH 2 OH, from the hydro-
chloride by causing the following reaction to take place in methyl
alcoholic solution:

= NaCl + HOCH 3 + NH 2 OH.

Sodium methoxide, NaOCH 3 , reacts effectively with hydrogen
chloride because of the insolubility of sodium chloride in methyl
alcohol, CH 3 OH, and the methyl alcohol produced at the same time
becomes part of the solvent, so that after the precipitated sodium
chloride has been removed nothing remains in solution but free
hydroxylamine.

The reaction,

which is of wide, terrestrial significance, serves further to illustrate
the principle of volatility. If silica is heated with sodium carbonate,
the reaction proceeds from left to right, because carbon dioxide is
volatile and silica is not; but if carbon dioxide acts on sodium silicate
in aqueous > solution, the reaction is reversed, because carbonic acid
is soluble in water, whilst hydrated silica is not so soluble, and
largely separates from solution. Thus, in the early ages of the
earth's history much silica was probably combined with bases in
the earth's crust, and much free carbon dioxide existed in the
atmosphere: but "weathering" of siliceous rocks has now been
going on for ages in presence of water and at moderate tempera-
ture, with the fixation of carbon dioxide and consequent separation
of silica.

The phenomena of precipitation, % so important in analysis, may
now be seen in a clearer light. Some precipitates are amorphous,
some crystalline. Suddenly - formed precipitates are frequently
amorphous, because they have not had time to crystallize, and
occasionally they become crystalline when kept in contact with
the liquid from which they have been separated.

CHEMICAL CHANGES CLASSIFIED 295

Thus magnesium ammonium phosphate, MgNH 4 PO 4 6H 2 O, when
it is precipitated from concentrated solutions, appears amorphous,
though when it separates gradually from dilute solutions it is
distinctly crystalline.

Similarly barium sulphate, precipitated from cold, dilute solu-
tions, is so finely divided as to run through the pores of ordinary
filter-paper, but when the precipitate is heated with water it slowly
becomes granular, that is, micro-crystalline, and can then be easily
filtered. The same condition is more quickly assumed if the pre-
cipitation is carried out by mixing boiling solutions.

Calcium carbonate is another example of an amorphous precipi-
tate which may become crystalline; for when heated in the liquid
from which it has been precipitated it gradually assumes the crys-
talline form.

Sulphide precipitates are, however, colloidal, 1 and do not crystal-
lize. They are exceedingly insoluble in water, though some of them
may assume "the state of colloidal suspension. A curious case is
presented by the sulphides of nickel and cobalt. These precipitates
do not dissolve readily in dilute hydrochloric acid, though this acid
prevents their formation. The explanation of this paradox is that
on precipitation these sulphides polymerize, i.e. simple molecules
combine to form more complex and therefore less soluble molecules.
Thus dilute acid prevents the formation of the simple and more
soluble molecules of the sulphides, but is unable to dissolve readily
the more complex molecules formed by precipitation from alkaline
solution.

The student should now understand that crystallization and
precipitation are linked phenomena. Indeed, crystallization from
solution is slow precipitation; the slower the precipitation the
larger and more perfect are the crystals; the more rapid the preci-
pitation the smaller and less perfect are the crystals, until precipi-
tation becomes too rapid, and the separated solid is amorphous.

He will also understand that chemical affinity has little if any-
thing to do with precipitation. For example, the reaction,

NaCl + AgN0 3 = AgCl +

takes place in aqueous solution, not because silver has a greater
affinity for chlorine than sodium, for that is not the case, or because
sodium has a greater affinity for nitrate than for chloride, but

1 Vide Chapter XII, on the colloidal state.

296 CHEMICAL THEORY

because silver chloride happens to be practically insoluble in water,
and a reaction which might be reversible fails to be so because
the effective mass of the silver chloride is so small. Thus silver
nitrate is a reagent for all chlorides, or, more properly, silver
ions are a reagent for chloride ions from whatever source they
are derived.

Oxidation and Reduction.

Berzelius called oxygen the pole of chemistry; and it is true
that no other element occupies so significant a position among the
rest. It has already been seen that the kinds of oxides an element
forms go far to reveal the chemical nature of the element.

Again, Lavoisier invented the name oxygen, thinking that this
element was a constituent of all acids; and although this idea was
false the name was not badly chosen after all, for the addition of
oxygen to a substance generally enhances its acidic character.

Consider, for example, the oxides of manganese:

MnO, Mn s O 4 > Mn 2 O 3 , MnOa, MnO 3 , Mn 2 O 7 .

The first oxide, MnO, is wholly basic, but with successive
additions of oxygen the oxides lose their basic power, becoming
more and more acidic, until Mn 2 7 , a strongly acidic anhydride,
is reached.

Oxidation is the addition of oxygen to an element or compound',
reduction is its removal.

The removal of hydrogen is also regarded as oxidation, and
sometimes also the addition of chlorine or an equivalent electro-
negative atom or radicle.

Thus, alcohol is oxidized to aldehyde by the removal of
hydrogen:

C 2 H 6 O + O = C 2 H 4 O + H 2 O;

and the conversion of a ferrous into a ferric compound is oxidation,
however effected; e.g.

2 FeS0 4 + H 2 S0 4 + O = Fe 2 (S0 4 ) 3 + H 2 O,
or 2FeCl 2 + Cl 2 = 2FeCl 3 . '

Oxidation generally involves an increase in the valency of the
nuclear element of the compound . oxidized; for example iron is
bivalent in ferrous compounds and trivalent in ferric. This is not

CHEMICAL CHANGES CLASSIFIED

297

always the case, however, for carbon is quadrivalent in both alcohol

and aldehyde:

H H

CHy-i-OH, CH 3 -C=0,

H
and barium is bivalent in both BaO and Ba0 2 :

Ba=O,

Similarly, the addition of hydrogen to a molecule may be
regarded as reduction, as for example in the formation of leuco-
compounds from organic dyes through addition of hydrogen by
means of sulphurous acid, thus:

H 2 SO 3 + QH 2 = H 2 S0 4 + 2 H.

Oxidation and reduction are reciprocal processes; the oxidation
of one substance often involves the reduction of another; as for
instance the oxidation of ferrous sulphate by nitric acid:

6 FeSO 4 + 3 H 2 SO 4 + 2 HNO 3 = 3 Fe 2 (S0 4 ) 3 + 4 H 2 O + 2 NO.

This reciprocal relationship necessarily exists when the reducing
and oxidizing substances are both compounds; indeed the process
is essentially the transference of oxygen from one compound to
another. v /

Common Oxidizing and Reducing Agents.

The following are the principal oxidizing and reducing agenta
met with in inorganic chemistry.

OXIDIZING AGENTS

Free oxygen and ozone.

Hydrogen peroxide and other per-
oxides.

Keducible basic oxides, e.g. Ag 2 0.

The halogens, oxyacids and their salts,
e.g. nitric and chloric acids and
their salts.

Higher oxides and oxyacids of metals,
e.g. chromic and permanganic acids
and their salts.

REDUCING AGENTS

Hydrogen, gaseous and " nascent ".
Unstable hydrides, e.g. H 2 S, HI.
Oxidizable elements, e.g. carbon, the

alkali metals, and magnesium.
Lower oxides, e.g. CO.
Lower oxyacids and their salts, e.g.

sulphites, nitrites.
Lower salts, e.g. ferrous, stannous.
Cyanides, formates.

298 CHEMICAL THEORY

Conditions of Action.

Oxidation and reduction may take place either in presence or
absence of water. These two conditions are well illustrated by
the reactions of qualitative analysis carried out in the dry way
and in solution respectively.

A familiar example of oxidation in the dry way is the for-
mation of a chromate by oxidizing fusion, thus:

Cr 2 O 3 + 2 Na 2 CO 3 + 3 NaNO 3 = 2 Na 2 CrO 4 + 3 NaNO a + 2 CO 2 ,
or Cr 2 O 3 + 3Na a O 2 = 2Na 2 CrO 4 + Na 2 O;

and of reduction in the dry way the liberation of a metal from
its oxide by heating it with charcoal in a reducing flame, e.g.:
SnO 2 + 2C = Sn + 2CO.

The dry reactions of qualitative analysis, which are frequently
neglected, are at any rate of value because they illustrate in
miniature important manufacturing processes.

The oxidation of chromic oxide, for example, in presence of
alkali, is employed to obtain chromium compounds from the
natural source of chromium, chrome ironstone or ferrous chromite:
FeOCr 2 O 3 . It is interesting to observe, moreover, that although
Cr 2 O 3 is not oxidized when heated alone in the air, and that on
the contrary Cr0 3 loses oxygen under these conditions, nevertheless,
in presence of a base with which CrO 3 can combine chromate is
formed, thus:

4 FeO -Cr 2 O 3 + 7 O 2 + 8 CaCO 3 = 8 CaCrO 4 + 2 Fe 2 O 3 + 8 CO 2 .

From this it may be inferred that the following is a reversible

reaction:

2 Cr 2 O 3 + 3 O 2 ^= 4 CrO 3 ,

the presence of* a base with which it can combine promoting the
formation of the acidic oxide CrO 3 , whilst the absence of a base,
being the absence of the condition of stability of this oxide, permits
its decomposition by heat.

Further, it is easy to understand why Cr0 3 oxidizes hydro-
chloric acid according to the reaction

2 CrO 3 -f 12 HC1 = 2 CrCl 3 + 6 H 2 O + 3 C1 2 ,

for the acid is oxidizable, and at the same time promotes the for-
mation of the salt of the lower and basic oxide Cr 2 O 3 .

Similar considerations apply to the oxides of manganese. The

CHEMICAL CHANGES CLASSIFIED 299

oxide Mn0 8 loses oxygen when heated alone, but manganates
corresponding to it are formed when any other manganese oxide
is heated in air with alkali; e.g.:

2 MnO 2 -f O 2 H- 2 Na 2 CO 3 = 2 Na 2 MnO 4 + 2 CO 2 .

Thus a deep-green mass of sodium manganate is produced when
a trace of a manganese compound is heated strongly with fusion
mixture. Without alkali, however, no manganate appears; so
when manganese is tested for by means of a borax bead, the bead
is not green, but amethyst in the oxidizing and colourless in the
reducing flame, these colours being due to manganic and manganous
borates, derived from Mn 2 3 and MnO respectively.

The processes of metallurgy, i.e. the winning of metals from
their ores, afford numerous examples of reduction in the dry way;
and these are often imitated on a small scale by blowpipe and other
laboratory reactions.

The following reactions are typical of reduction in the dry way:

K 2 CO 3 -t-2C = 2K + 3CO.

Fe 2 3 + 3C = 2Fe + 3CO.

4KOH + 3Fe =* 4K -f Fe 3 O 4 -f 2H 2 .

Sb 2 S 5 + 3 Fe = 2 Sb + 3 FeS.

Na 2 SiF + 4 Na = Si + 6 NaF.

CoO + H 2 = Co + H 2 O.

2 AgCl -f 2 Hg = 2 Ag + Hg 2 CI 2 .

Cr 2 O 3 + 2 Al = 2 Cr -f A1 2 O 3 .

2PbS + 3O 2 = 2PbO-f2SO 2 /i

2PbO + PbS = 3Pb + SO 2 . I

PbSO 4 + PbS = 2 Pb + 2 S0 2 . J

The last three reactions are remarkable. They concern the
metallurgy of lead by what is paradoxically called the "air-
reduction process". A similar reaction occurs in the metallurgy
of copper, between cuprous oxide and sulphide, thus:

2 Cu 2 O + Cu 2 S = 6 Cu + SO 2 .

Oxidation and reduction in solution are frequent laboratory
operations.

The reactions of nitric acid have already been studied.

Sodium peroxide, or hydrogen peroxide in presence of alkali,
is frequently employed, e.g. to oxidize precipitated chromic
hydroxide to chromate:

2 Cr(OH) 3 + 3 Na 2 O 2 = 2 Na 2 CrO 4 + 2 NaOH + 2 H 2 O.

300 CHEMICAL THEORY

Similarly sulphide is oxidized to sulphate:

Na 2 S + 4 Na 2 a + 4 H 2 O = Na^SO^ + 8 NaOH.

The oxidation of sodium sulphide, or rather hydrogen sulphide,
by nitric acid, produces, however, not sulphuric acid in the first
instance, but free sulphur, thus:

2HN0 3 =

although by the action of hot, concentrated nitric acid sulphur is
gradually converted into sulphuric acid. This difference between
the manner of oxidation of a sulphide in acid and alkaline solution
is a further illustration of the influence of conditions on the course
of a reaction.

Chloric acid, derived from a mixture of potassium chlorate and
hydrochloric acid, is sometimes employed in qualitative analysis to
oxidize finely-divided sulphur to sulphuric acid When a chlorate
is heated with hydrochloric acid a mixture of chlorine and chlorine
dioxide, Davy's " euchlorine ", is evolved, the reaction being most
simply represented, thus:

2 KC1O 3 + 4 HC1 = 2 KC1 + 2 C1O 2 -f C1 2 + 2 H 2 O.

It does not follow, however, that this proportion between
chlorine dioxide and chlorine is maintained, for the former may
in turn oxidize hydrochloric acid, thus:

2C1O 2 + 8HC1 = 4H 2 O + 5C1 2 ,

so that the proportion of chlorine is increased; but that is im-
material, for the mixture of gases evolved represents quantitatively
the oxygen content of the chlorate, so that an equivalent amount
of iodine would always be liberated from hydriodic acid, thus:

2KC1O 3 +12HI = 2KCl + 6H 2 O-f 6I 2 .

A very sensitive reaction is that between iodate and iodide in
presence of dilute acid, thus:

KI0 8 + 5 KI + 3 H 2 SO 4 = 3 K 2 SO 4 + 3 H 2 + 3 1 2 ;

indeed the acidity of a solution may be estimated by the iodine
liberated when it is mixed with excess of iodate and iodide, which
do not interact in neutral solution.

Among the most important oxidizing reactions in solution are

CHEMICAL CHANGES CLASSIFIED 301

those of permanganate and dichromate, so commonly employed in
volumetric analysis.

The reactions of permanganate afford a most interesting example
of the principle already noticed from time to time that differing
conditions determine which of several possible reactions shall take
place.

The following scheme represents in terms of oxides the reduction
of permanganate in stages, with corresponding colours of the reduc-
tion products:

Mn 2 O 7 * 2MnO 3 + O -* 2MnO 2

Crimson Green. Brown. Pale Pink

Permanganate. Manganate. Hydrated manganese dioxide, Manganoua

or manganous acid. salt.

These stages may be readily observed by the use of sulphite
as reducing agent. Thus, if a little neutral or alkaline sulphite
solution is added to a dilute solution of permanganate, made some-
what alkaline, the colour becomes deep green, owing to the forma-
tion of manganate. If more sulphite is added, and especially if the
solution is not too alkaline, and is warmed, the green solution gives
place to a brown and turbid liquid containing hydrated manganese
dioxide. Finally, if sulphite is added to an acidified permanganate
solution, the colour is instantly discharged, manganous salt being
produced. Thus an alkaline solution with little reducing agent pro-
motes the formation of manganate; a nearly neutral solution with
more reducing agent causes the manganese dioxide stage to be
reached, whilst the presence of acid is the best condition for
complete reduction to manganous salt.

All this is what might be expected; MnO 3 is an acidic oxide, most
likely to be permanent as a salt in presence of alkali; MnO is a
basic oxide, and its salts are likely to be formed in presence of
acid; whilst MnO 2 is neither strongly basic nor acidic, and is there-
fore likely to result when the solution is nearly neutral.

It must not, however, be assumed that reduction cannot proceed
beyond the stage of MnO 2 in absence of acid, for this is not true.

It is possible to create conditions of stability for compounds
corresponding to MnO in presence of alkali. Thus, if a few drops
of permanganate solution are added to a solution of alkali sulphide
containing free alkali, the brown precipitate of hydrated MnO 2 first
observed quickly becomes paler as it mixes with the solution, being
converted by excess of reducing agent into the less-soluble MnS.

302 CHEMICAL THEORY

Or, if the same permanganate solution is added to excess of alkali
sulphite containing much ammonium chloride, the brown precipitate,
when heated with the liquid, completely dissolves, yielding a colour-
less solution, because, as the student should know, manganous solu-
tion can remain unprecipitated by alkali in presence of much
ammonium chloride.

Occasionally a compound intermediate in oxygen content be-
tween two other compounds may behave either as an oxidizing or
a reducing agent, according to circumstances. Instances of such
compounds are nitrous acid and the aldehydes. This double func-
tion of nitrous acid may be expressed in terms of oxides, thus:

-o 4-20

2 NO N 2 3 N 2 6 ;

that is to say, nitrous acid may be oxidized to nitric acid or reduced
to nitric oxide.

If a solution of nitrite is carefully added to an acidified solution
of permanganate the latter is decolourized, whilst the nitrite is
oxidized to nitrate without evolution of gas, thus:

2 KMnO 4 + 5 HNO 2 + 3 H,SO 4 = K 2 SO 4 + 2 MnSO 4 + 3 H 2 O + 5 HNO 3 .

If, however, nitrite solution is mixed with hydriodic acid or
acidified potassium iodide, the following reaction occurs:

2HNO 2 + 2HI = 2H 2 O + 2NO + I 2 ;

the hydriodic acid being oxidized with liberation of iodine, whilst
nitric oxide gas escapes as the reduction product of nitrous acid.
Nitrous acid can also oxidize ammonia, nitrogen resulting, both as
the reduction product of the former and the oxidation product of
the latter; so ammonium nitrite decomposes, thus:

NH 4 NO 2 = N 2 + 2H 2 O.

Aldehydes are intermediate as regards oxygen content between
alcohols and carboxylic acids, e.g.:

CH 3 -CH 2 OH CH 3 COH CH 3 -COOH

Ethyl alcohol. Acetaldehyde. Acetic acid.

So an aldehyde may be reduced to an alcohol, thus behaving as an
oxidizing agent, or oxidized to an acid, thus behaving as a reducing
agent.

Nascent hydrogen reduces an aldehyde thus:
CH 3 -COH + 2H = CH 2 -CH 2 OH;

CHEMICAL CHANGES CLASSIFIED 303

but silver oxide oxidizes an aldehyde thus:

CH 3 -COH + Ag 2 O = CH 3 COOH + 2Ag.

Hydrogen peroxide may perform both oxidizing and reducing
functions, though in both cases it is reduced to water, so that when
it reduces a compound, oxygen gas is evolved. The following
equations will make this plain:

H 2 O 2 + X = XO + H 2 O, oxidizing function.
H 2 O 2 + XO = X + H 2 O + O 2 , reducing function.

The behaviour of hydrogen peroxide towards manganese com-
pounds furnishes an admirable example of the influence of acidity
or alkalinity on the course of a reaction. If H 2 O 2 is added to
manganese hydroxide in presence of alkali, oxidation occurs, thus:

Mn(OH) 2 + H 2 O 2 = MnO 2 + 2 H 2 O;

but if the liquid containing the precipitated Mn0 2 is now acidified,
reduction to manganous salt is brought about by H 2 O 2 , thus:

MnO 2 + H 2 SO 4 + H 2 O 2 = MnSO 4 + 2 H 2 O + O 2 .

Permanganate is similarly reduced in acid solution, thus:

2 KMnO 4 + 3 H 2 SO 4 + 5 H 2 O 2 = K 2 SO 4 + 2 MnSO 4 + 8 H 2 O + 5 O 2 ;

reduction of a compound by H 2 O 2 does not, however, necessarily
take place in presence of acid, for chromic acid is oxidized in acid
solution to the deep blue perchromic acid. But in general, since
the lower oxides of a metal are basic, they are likely to be produced
by reduction in presence of acids, whilst the higher oxides, if acidic,
are more likely to result by oxidation in presence of alkali.

The above reaction between permanganate and hydrogen per-
oxide isva case of the mutual reduction of two oxidizing agents,
each of which contributes one atom of oxygen towards a molecule
of this gas. Similar examples are furnished by the reactions be-
tween H 2 O 2 and silver oxide and ozone respectively:

H 2 O a + Ag 2 O = 2 Ag + H 2 O + O 2 .
H 2 O 2 + O 8 = H 2 O + O 2 + Oj,

A reaction which may be regarded as the converse of the above
is the partition of molecular oxygen between water and an oxidizable
substance, as in the following reaction, which is believed to be part

304 CHEMICAL THEORY

of what takes place in the process of solution of gold by potassium
cyanide in presence of air:

2 Au + H 2 O + O 2 = Au 2 O + H 2 O 2 .

This reaction is an example of auto-oxidation, i.e. spontaneous
oxidation at atmospheric temperature. Another example is the
following interaction of zinc, water, and air in presence of very
dilute sulphuric acid:

Zu + 2H 2 O + O 2 = Zn(OH) 2 + H 2 O 2 .
Zn(OH) 2 + H 2 SO 4 = ZnS0 4 + 2H 2 O.

There are a few compounds, intermediate in oxygen content,
which, under suitable conditions, undergo self-oxidation and reduc-
tion, producing compounds poorer or richer in oxygen respectively
than themselves. Thus a hypochlorite solution, when boiled,
passes into chloride and chlorate, thus:

3KOC1 = 2KC1 + KC1O 3 ;

whilst chlorate, when heated suitably, yields chloride and per-
chlorate, thus:

4KC10 3 = KC1 + 3KC10 4 .

Phosphite and hypophosphite pass by heating into phosphate
and phosphine, thus:

4 Na 2 HPO, = 2 lSra 3 PO 4 + Na a HPO 4 + PH 3 .
2NaH 2 PO 2 =

and sulphite and thiosulphate behave thus when heated:

4Na 2 S0 3 =
4 Na 2 S 2 O 3 = 3 Na 2 S0 4

In all these cases the compounds pass into others more stable
under the given conditions.

SUMMARY

HYDROLYSIS. Hydrolysis is the chemical decomposition of a
compound by water.

OXIDATION AND REDUCTION. Oxidation is the addition of
oxygen to an element or compound; reduction is its removal.

CHAPTER XVI
EQUATION-BUILDING

The student of chemistry becomes familiar early with chemical
equations. It may be that before he mastered the rudiments of
the atomic and molecular theories he was taught to employ the
equation as a brief and pointed way of stating what happens in
chemistry; and that without fully understanding their significance
he has committed to memory a number of equations representing
the reactions that occur in the varied preparations of his elementary
course. Consequently he has been in danger of magnifying the
equation unduly and regarding it as a kind of talisman by means
of which natural processes may be foretold or brought to pass.
He may have supposed that by manipulating the formulae of
certain reacting substances, and of other substances that might
result from their interaction, a chemical change may be successfully
expressed without any practical experience of that change. This,
in fact, is how beginners often behave in answering questions.
Their chief concern is to make the equation balance, in faithfulness
to the principle of the indestructibility of matter, supposing that
the exigencies of chemical science are completely satisfied if nothing
is lost by the way.

The student must learn that chemistry is not a branch of
mathematics, that chemical equations are not to be solved like
algebraic equations. To employ a chemical equation for a reaction
until that reaction is understood qualitatively and quantitatively,
so far as the distribution of matter it expresses is concerned, is
unscientific and vain.

What is the use of a chemical equation? it may be asked.

Such a question is best answered by considering what chemistry
would be without equations. The science might still exist, but it
would be an exceedingly clumsy science, and would probably be
in a much more rudimentary state than it is at present. Nature

(D00) 306 21

306 CHEMICAL THEORY

has not ordained the chemical equation; it is a human invention;
and man might still be mixing or heating things together and
watching for results if he had never invented a means of expressing
briefly the results of his discoveries. What language is to thought,
so roughly the equation is to chemistry; and just as the educated
man chooses language to express his thoughts, and avoids the
danger of allowing his language to outstrip his thoughts, so the
chemist uses equations to express discovered chemical facts, and
avoids allowing his equations to outstrip his facts.

Chemical equations, then, are chemical notation, that is, a means
of noting chemical facts in a convenient form. Facts first; equation
afterwards. Let the student remember this, and he will at least
be in the way of learning chemistry properly.

Nevertheless, to express chemical reactions by equations, not
learned by rote, but developed intelligently from well-understood
principles, is an art which should be learned by every student of
chemistry. It is the present purpose to show how even the most
complex equations met with in inorganic chemistry may be built
up when the underlying principles of the reactions they express
are understood.

Consider the simplest possible reaction; the preparation of
hydrogen by the interaction of zinc and dilute sulphuric acid.

If the student knows the symbol Zn and the formula H 2 SO 4 ,
and knows also that zinc displaces the hydrogen from the sulphuric
acid, he may still go wrong with his equation.

For how shall he decide whether:

2Zn + H 2 SO 4 = Zn 2 SO 4 + 2H,
or Zn + H 2 SO 4 = ZnSO 4 + 2H,
or Zn + 2 H 2 SO 4 = Zn(SO 4 ) 2 + 4 H,

or some other relation between the zinc and the displaced hydrogen
is the right one? He may be told what the right relation is, and
therefore what the proper equation should be, but he should also
be told why.

Now, whilst the equivalent weight of zinc is 32-65, the atomic
weight of this metal, as shown by the method of specific heats and
other methods is 65-3. Therefore, an atom of zinc displaces 2
atoms of hydrogen, and the equation becomes:

Zn + H 2 SO 4 = ZnSO 4 + 2 H.
There remains another question, however; that is, whether it is

EQUATION-BUILDING 307

quite proper to represent the atom of zinc or hydrogen as remaining
single in the elementary state. The answer is that zinc is known
so to exist, but that hydrogen gas consists of molecules H 2 . There-
fore, amended, the equation finally becomes:

Zn + H a SO 4 = ZnS0 4 + H*

When a student has his equation he is inclined to ask whether
it explains everything. Then he has to be told that it explains
nothing. A chemical equation does not explain a chemical reaction;
it expresses it, with certain limitations. It is plain that the
equation gives but a limited expression to the reaction between
zinc and sulphuric acid; for it does not express the fact that the
acid must be considerably diluted with water before any hydrogen
can be obtained.

The student will be aware that there are other circumstances
of this reaction of which this simple equation gives no account.
Nevertheless, it does express, qualitatively and quantitatively, a
reaction between the metal and the acid.

If concentrated instead of dilute acid is added to zinc, there
is little if any evolution of hydrogen; but when the mixture is
heated a vigorous reaction takes place, and much sulphur dioxide
gas is evolved.

If this quite different reaction is to be expressed by an equation,
it must first be understood. Sulphur dioxide necessarily comes
from the sulphuric acid, whence it is derived by reduction, which
may be represented thus:

H 2 S0 4 -0 = H 2 + S0 2 ,
or H 2 S0 4 = H 2 + S0 2 + O.

It is pertinent to ask how this reduction is effected. Now,
since hydrogen is a reducing agent, a facile theory of the process
is the following:

Zn + H 2 SO 4 = ZnSO 4 + 2 H.
2 H + H 2 S0 4 = 2 H 2 O + S0 2 .

The student may now add these two equations together, in this
case treating them as if they were algebra; then, since hydrogen
is eliminated, the final equation becomes:

Zn + 2 H 2 S0 4 = ZnSO 4 + 2 H 2 O + S0 2 .
Now this is the equation for the reaction in question whether

308 CHEMICAL THEORY

hydrogen is liberated, and in the nascent state acts as a reducing
agent, or not.

But it is certainly open to question whether hydrogen is liberated
at all in this reaction; and there is no need to assume that it is.
Zinc itself is a reducing agent, so why should it not react directly?
Now, whilst zinc oxide cannot appear in presence of excess of acid,
it is symbolic of the state of oxidation assumed by the zinc; so that
if the reaction is to be represented without assuming the agency of
liberated hydrogen, the following device may be adopted:

Zn + H 2 SO 4 = ZnO + H 2 O + SO 2
ZnO + H 2 SO 4 = ZnSQ 4 + H 2 O
and adding Zn + 2 H 2 SO 4 = ZnSO 4 + 2 H 2 O + SO 2 .

The same result is shown as above, and if it is not assumed that
ZnO is actually separated, this method of arriving at the result is
freer from assumption than the former one.

Equations which often puzzle beginners are those representing
the interaction of hydrogen sulphide with metallic salt solutions.
The reaction

CuSO 4 + H 2 S = CuS + H 2 SO 4

is simple enough, because the molecules are matched; but how
is SbCl s to be represented as reacting with H 2 S? The answer is
simple if it is remembered that the reaction is one of double
decomposition; antimony chloride + hydrogen sulphide give anti-
mony sulphide + hydrogen chloride. That is, hydrogen and chlorine
atoms must be in equal numbers on the left-hand side of the equa-
tion to produce hydrogen chloride molecules on the right. This
can result only if the equation is written:

2 SbCl 3 + 3 H 2 S = Sb 2 S 3 + 6 HC1;

thus not only showing six complete molecules of hydrogen chloride,
but also satisfying the trivalency of antimony.
The reaction

3 Cu + 8 HNO S = 3 Cu(NO s ) 2 + 4 H 2 O + 2 NO

may now be considered.

Since copper does not displace hydrogen from dilute acid, the
nascent hydrogen theory will not be assumed to account for the
reduction of the nitric acid. The copper will be represented as
being directly oxidized by the nitric acid, but as appearing as
nitrate instead of oxide, because of the excess of acid present. It is

EQUATION-BUILDING 309

important, however, to realize how the nitric acid is reduced to the
state of nitric oxide, and this reduction may be represented thus:

2HN0 3 = H 2 + N 2 6 = H 2 + 2NO + 30.

To remove these 3 atoms of oxygen from 2 molecules of nitric
acid 3 atoms of copper are needed, thus:

30u + 3O =CuO;

but nitrate is formed rather than oxide, thus:

3 CuO + 6 HNO 3 = 3 Cu(NO 3 ) 2 + 3 H 2 O.

And now, adding together these three equations, the following
equation is obtained as representing the actual reaction that takes

place:

3 Cu + 8 HNO 3 = 3 Cu(NO 3 ) 2 + 4 H 2 O + 2 NO.

Thus, by this method of construction it becomes apparent why
3 atoms of copper and 8 molecules of nitric acid are required; and
the student should now recognize that the equation has been built
by a sound method based on an understanding of the scientific
principles involved. He may ask himself how otherwise he would
produce the equation if he did not remember it. Only by hap-
hazard and guesswork could he make the attempt.

In the chapter on the Classification of Chemical Changes the
following equations are found:

i. 4Zn 4- 10 HNO 3 = 4Zn(NO3) 2 + 5 H 2 O + N 2 O,
and ii. 4Zn + 10HNO 3 = 4Zn(NO 3 ) 2 + 3 H 2 O + NH 4 NO 3 .

They may be built as follows:

i. 2 HNO 3 = H 2 O + N 2 O + 4O
4Zn-f 4O = 4ZnO

4 ZnO + 8 HNO 3 = 4 Zn(NQ 3 ) 2 + 4 H 2 Q

adding 4 Zn + 10 HNO 3 = 4 Zn(NO 3 ) 2 + 5 H 2 O + N 2 O;

or, by the nascent hydrogen theory, which may be preferable here:

2HNO 3 = H 9 O +
4Zn + 8HNO 3 = 4Zn(NO 3 ;
8H + 4O = 4H 2 O

adding 4 Zn + 10 HNO 3 = 4 Zn(NO 3 ) 2 + 5 H 2 O + N 2 O

ii. HNO 3 + 8H = 3H 2 O-f NH,
4Zn + 8HNO 3 = 4Zn(NO 3 ), + 8H

NH 3 + HNO, = NH 4 NO 3

Adding 4 Zn + 10 HNO 3 = 4 Zn(NO 3 ) 2 + 3 H 2 O + NH 4 NO 3
1 See, however, note in Appendix on Ionic and Electronic Equations.

310 CHEMICAL THEOEY

The student will see that the method of procedure is to fix on
the essential reduction product, and consider how it can be formed.
Now, in the above example, a molecule of ammonia is formed from
one of nitric acid by the action of 8 atoms of hydrogen. These
are therefore produced from 8 other molecules of the acid. So the
final equation is easily and surely obtained.

Allied to these reactions are those concerned in the use of
permanganate and dichromate in volumetric analysis. Consider
the following equations representing the oxidation of ferrous iron:

2KMnO 4 +10FeSO 4 + 8H 2 SO 4 =
K 2 Cr 2 O 7 + 6 FeSO 4 + 7 H 2 S0 4 = K 2 SO 4 + Cr 2 (SO 4 ) 3 + 3 Fe 2 (SO 4 ) 3 + 7 H 2 0.

It may be observed in passing that the numbers of molecules
of sulphuric acid required and of water formed correspond with
the number of oxygen atoms in the oxidizing agent. But this
circumstance does not really furnish a key to the reactions. This
is supplied by considering the available oxygen of the oxidizing
agent and the oxygen essential for oxidizing the iron.

Expressed in terms of oxides, these are shown thus in the case
of permanganate:

2 KMnO 4 = K 2 O + Mn 2 O 7 = K 2 O + 2 MnO + 5 O.
2FeO + O = Fe 2 3 .

Thus, since 2 molecules of permanganate in presence of acid yield
5 atoms of oxygen for oxidizing purposes, together with potassium
and manganous salts, to which the above oxides correspond; and
2 atoms of iron require 1 atom of oxygen to oxidize them from
the ferrous to the ferric state, it appears that 10 molecules of ferrous
sulphate are oxidized by 2 molecules of permanganate in presence
of sufficient acid, the quantity of which participating in the reaction
is easily arrived at by adding together the following equations:

2KMnO 4 = K 2 O + 2MnO + 5O
10 FeSO 4 + 5 O + 5 H 2 SO 4 = 5 Fe^SOJg + 5 H 2 O

K 2 O + H 2 S0 4 = K 2 SO 4 + H 2 O

__ 2 MnQ + 2 H 2 SO 4 = 2 MnSO 4 + 2 H 2 O _
Adding 2KMnO 4 +10FeSO 4 +8H 2 SO 4 = K 2 SO 4 +2MnSO 4 +5Fe 2 (SO 4 ) 3 +8H 2 O.

The equation for the dichromate reaction is similarly obtained,
the chromium being reduced from the state of Cr0 8 to Cr 2 3 , thus:

2Cr0 3 = Cr 2 O 3 + 30;

EQUATION-BUILDING 311

and the iron oxidized by the available oxygen as before:

K 2 2 r = K 2 + Cr 2 3 + 30
6 FeSO 4 + 3 O + 3 H 2 SO 4 = 3 Fe 2 (SO 4 ) 3 + 3 H 2 O

K 2 + H 2 S0 4 = K 2 S0 4 + H 2
Cr 2 O 3 + 3 H 2 S0 4 = Cr 2 (SO 4 ) 3 + 3 H 2 O

Adding K 2 Cr 2 O 7 +6FeSO 4 +7H 2 SO 4 = K 2 SO 4 +Cr 2 (SO 4 ) 3 +3Fe 2 (SO 4 ) 3 +7H 2 O ;

Other equations for other reactions of permanganate and dichro-
mate are built up similarly, e.g. the following:

2 KMnO 4 + 5 H 2 C 2 O 4 + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4 + 8 H 2 O + 10 CO 2
2 KMnO 4 + 5 H,A + 3 H 2 SO 4 = K 2 SO 4 + 2 MnSO 4 + 8 H 2 O + 5 2 .

K,Cr 2 O 7 + 3 H 2 S + 4 H. 2 SO 4 = K 2 SO 4 + O 2 (SO 4 ) 3 + 7 H 2 O + 3 S.
K 2 O 2 O 7 + 6 HI + 4 H,SO 4 = K 2 ISO 4 + Cr 2 (SO 4 ) 3 + 7 H 2 O + 3 I 2 .

The reactions of potassium chlorate with sulphuric and hydro-
chloric acids are interesting, and the building of equations to repre-
sent them is instructive.

When sulphuric acid is added to potassium chlorate chloric acid
is liberated thus:

2 KC1O 3 + H 2 SO 4 = K 2 SO 4 + 2 HC1O 3 ;

but this acid is dehydrated when it is heated with sulphuric
acid.

The anhydride C1 2 O 5 , however, does not exist; and in place of
it 2ClO 2 -f O are produced. The oxygen, which would otherwise
accompany the chlorine dioxide, does not appear as gas, but oxidizes
some of chloric to perchloric acid, which is stable in presence of
sulphuric acid; or, otherwise expressed, chloric acid undergoes self-
oxidation and reduction thus:

3 HC1O 3 = 2 ClOj + H 2 O + HC1O 4 .

Consequently, the equation representing the reaction between
potassium chlorate and sulphuric acid may be built up thus:

2 KC1O 3 + H 2 SO 4 = K 2 SO 4 + 2 HC1O 3
2 HC1O 3 = H 2 O + 2 C1O 2 + O
KC1O 3 + O = KC1O 4 _
Adding 3 KC10 3 + H 2 S0 4 = K 2 SO 4 + KC1O 4 + H 2 + 2 C10 2 .

If potassium chlorate is heated with hydrochloric instead of

312 CHEMICAL THEORY

sulphuric acid, chlorine accompanies the chlorine dioxide, since
hydrochloric acid is oxidized in preference to chloric acid:

2 KClOg + 2 HC1 = 2 KC1 + 2 HC1O 3

2 HC1O 3 = H 2 O + 2 C1O 2 + O
2HC1 + = H 2 + C1 2

Adding 2 KC10 3 4- 4 HC1 = 2 KC1 + 2 H 2 O + 2 C1O 2 + C1 2 .

The interaction of sodium hydrogen sulphite and sodium iodate
in aqueous solution, by which iodine is liberated quantitatively from
the iodate, is a reaction the equation for which may be built up
from first principles, thus:

2NaIO 3 = Na 2 O + I 2 O 6
I 2 6 + 5 NaHS0 3 = 5 NaHSO 4 + I 2
2 NaHSO 4 + Na 2 O = 2 Na 2 SO 4 + H 2 Q _
Adding 2 NaIO 3 + 5 NaHSO 3 = 2 Na,^ + 3 NaHSO 4 + H 2 O + I 2 .

The equation for the cyanide process for gold extraction may be
built up as follows:

2 Au + O 2 + H 2 O = Au 2 O + H 2 O 2
2 Au + H 2 O 2 = Au 2 O + H 2 O
2Au 2 O + 8 KCN + 2 H 2 O = 4 KAu(CN) 2 + 4 KOH

Adding 4Au + 8 KCN +2 H 2 0+ O 2 = 4 KAu(CN) 2 + 4 KOH.

The parts of this reaction are remarkable; for whilst the first
two equations represent what cannot possibly take place alone, since
gold is not susceptible of atmospheric oxidation, they likewise show
the fact of the intermediate formation of hydrogen peroxide through
the partition of the oxygen molecule between gold and water, which
is not shown in the final equation.

The preparation of phosphine by heating white phosphorus
with sodium hydroxide solution is a reaction which is represented
by the following equation:

P 4 + 3 NaOH + 3 H 2 O = 3 NaH 2 P0 2 + PH 3 ,

the product in solution being sodium *hypophosphite.

In endeavouring to build this equation the first question that
occurs is: Where does the hydrogen come from to produce the
phosphine? The answer is that it comes from the sodium
hydroxide, and the simplest representation of this fact is:

NaOH + P = NaOP + H

EQUATION-BUILDING 313

NaOP is not sodium hypophosphite, however; but if the elements
of water are added the formula for this salt, NaH 2 P0 2 , is obtained,
A further stage, therefore, is:

NaOH + H 2 O + P = NaH 2 PO 2 + IL

If now this equation is multiplied by three, and an atom of phos-
phorus is added to give PH 3 , the above equation results thus:

3 NaOH + 3 H 2 O + P 4 = 3 NaH 2 PO 2 + PH 3 .

Phosphonium iodide, PH 4 I, is prepared by dropping water on to
an intimate mixture of finely-divided phosphorus and iodine in an
inert atmosphere; and the following equation, representing the
reaction, is one of the most difficult in inorganic chemistry:

4 I + 4H 3 PO 4 .

It presents two problems: First, to understand why these products
result; and second, to discover the reason for the remarkable mole-
cular proportions exhibited by the equation.

There is apparently some connection between this reaction and
the foregoing, though water alone seems to perform the function of
the sodium hydroxide solution above. Then if hypophosphorous
acid itself is formed instead of its sodium salt, it is unstable, and
yields phosphoric acid and phosphine. Hydriodic acid is required,
however, to combine with the phosphine, and this may be produced
by the interaction of iodine and phosphorus to produce the iodide
which is subsequently hydrolyzed, yielding hydriodic and phos-
phoric acids. In this way the reaction may be accounted for, and
the equation built as follows:

8P+12H 2 O = 6 H 3 ?O 2 + 2 PH 3
6H 3 PO 2 = 3H 3 PO 4 + 3PH 3
P + 5I = PI 6

PI 6 + 4 H 2 O = H 3 P0 4 + 5 HI
5PH 3 + 5HI = 5PH 4 I

Adding 9 P + 5 1 + 16 H 2 = 4 H 3 PO 4 + 5 PH 4 I.

This is a rather complicated theory to account for the formation of
phosphonium iodide, but it may stand in default of a better; it has
the merit of accounting for a complicated equation.

314 CHEMICAL THEORY

These examples are sufficient to illustrate the art of equation-
building. Many other examples will occur to the student; and he
will discover by practice that any reaction of which he has an in-
telligent knowledge can be expressed by an equation constructed by
the application of the corresponding chemical principles.

APPENDIX

CHEMICAL NOMENCLATURE AND SYMBOLS

With the spread of chemical knowledge, and the accumulation
of chemical substances, a system of naming, i.e. of nomenclature,
became necessary. When Lavoisier broke the phlogiston theory,
and abolished the fantastic names connected with it, he proposed
names for elements and compounds to accord with the newer views;
and, with some modifications and additions, Lavoisier's nomencla-
ture has been retained.

Names of Elements.

If a list of the elements is examined it will be found that the
names of many metals end in -um or -ium the Latin neuter
suffix and that some names of non-metals, e.g. carbon, boron,
silicon, end in -on. These terminations are regarded as the proper
ones for the names of metals and non-metals respectively, and
newly discovered elements have generally received names in ac-
cordance with this convention. Recent examples among metals
are masurium, rhenium, and hafnium; and among non-metals,
argon, neon, &c.; although the name helium is an exception, so
that it has been proposed to change it to helion. Nevertheless, in
English, the common names of well-known metals, such as gold>
silver, copper, iron, are retained, almost of necessity, as well as
those of noil -metals, such as sulphur and phosphorus, though
arsenic is sometimes known as arseiiium. In the choice of symbols,
however, which must be international, the Latin rather than the
national name is used. Thus we write Au (aurum) not G or Go
for gold, and the Germans H for hydrogen, although their name for
this element is wasserstoff. Nevertheless, the French still sometimes
use Az as symbol for what we call nitrogen and symbolize by N, since

315

316 APPENDIX

Lavoisier named this gas azote. The names oxygen and hydrogen
we owe to Lavoisier, as is readily understood, because of this great
chemist's interest in these elements.

Symbols of the Elements.

Berzelius originated symbols for the elements which have been
retained and added to. To be international these were the initial
letters or two letters of the Latin names. Non-metals were given
the preference; e.g. B stood for boron, not for barium, C stood for
carbon, and not for calcium. When more than one element had
the same initial letter the symbol for the second element consisted
of its initial letter with the next letter not common to both elements.
Thus C = carbon, 01 = chlorine, Cr = chromium, Cu = copper,
Co = cobalt. Some non-metals also had symbols consisting of two
letters, because single-letter symbols were already appropriated;
thus C, B, S, necessitated Cl, Br, Si.

A symbol has always stood for an atom of an element, though
it may be used as an abbreviation for the element's name.

Berzelius also introduced the accepted way of indicating the
numbers of atoms in a compound, as e.g. CH 4 , C 2 H 4 , although his
numbers were superscript thus: C 2 H 4 , instead of subscript, a custom
still retained by the French.

Names of Compounds.

Names of chemical compounds, specially in organic chemistry,
are sometimes long and polysyllabic; but each syllable of such a
name has a meaning in relation to the constitution of the compound,
so that the name is not only justified, but far more useful than a
fanciful name, although shorter, could be.

Our system of nomenclature of inorganic compounds, the incep-
tion of which we owe to Lavoisier and his colleagues (Mtthode de
Nomenclature chimique, by MM. de Morveau, Lavoisier, Berthollet,
and de Fourcroy, Paris, 1787) may be set forth briefly as follows

A compound of oxygen with another element is called an oxide
(oxyde); consequently the name of the more non-metallic constituent
of a compound of two elements (a binary compound) ends in -ide t
e.g. chloride, sulphide, phosphide, &c. Hydrogen occupies an inter-
mediate position among the elements; consequently, although we
have hydrogen oxide, H 2 O, and sulphide H 2 S, we speak of calcium
hydride CaH 2 , because calcium is more metallic than hydrogen.

APPENDIX 317

If an element forms more than one compound with oxygen, or
other element more non-metallic than itself, the terminations -oua
and -ic imply respectively the less or greater amounts of oxygen,
&c., thus:

FeO is ferrous oxide ; FeCl 2 is ferrous chloride ;
Fe 2 O 3 is ferric oxide; FeCl 3 is ferric chloride.

The same suffixes are used with oxy acids containing less or
more oxygen; and for the corresponding salts the suffixes ite and
ate, thus:

H 2 SO 3 sulphurous acid; Na 2 SO 3 sodium sulphite;

H 2 SO 4 sulphuric acid ; Na 2 SO 4 sodium sulphate.

Sometimes this nomenclature is not sufficient. There are, for
example, four oxyacids of chlorine having the formulae; HC1O,
HC1O 2 , HC1O 3 , HC1O 4 . To meet such a case the prefixes hypo
(below), a,ndper (hyper above), are employed thus:

HC1O hypochlorous acid forming hypochlorites ;
HC1O 2 chlorous chlorites ;

H(J1O 3 chloric chlorates;

HC1O 4 perchloric perchlorates.

Similarly a peroxide contains excess of oxygen, e.g. BaO 2 ;
though sometimes the prefix per is used loosely and unnecessarily,
as when ferric chloride, FeCl 3 , is called perchloride of iron.

The prefix sesqui occasionally appears, though it is generally
(superfluous. It means one-half more, and relates to the relative
numbers of atoms of an element in a compound; e.g. Fe 2 O 3 is
sometimes called iron scsqui-oxide, from the relative numbers
of iron and oxygen atoms.

Numerical prefixes may be of Latin or Greek origin, thus:

1234 5 678

Latin: Uni Bi Ter Quadri Quinque Sexa Septa Octa.
Greek: Mono Di Tri Tetra Penta Hexa Hepta Octa.

There has been no rule or rigid custom regarding the use of
these prefixes, though it is generally agreed that hybrid words,
made of Latin and Greek, such as tetravalent, are bad.

Usually Greek prefixes appear in chemical names, e.g. in the
oxides of nitrogen: nitrogen monoxide, dioxide, trioxide, tetroxide,
pentoxide; and this seems right, since the word oxide is of Greek
origin; though we have carbon bi- or di-sulphide, and potassium
quadr- or tetr-oxalate. As mentioned in the text (p. 57) a pro-

318 APPENDIX

posal has been made recently to standardize the prefixes for
valency thus;

Uni-, bi-, tri-, quadri-, quinque-, sexi-, septi-, octi-valent.

HYDRION CONCENTRATION AND pH VALUE

Hydrion or hydrogen ion is the hydrogen atom minus its electron.
It is the acidifying principle of acids, and is present in water to a
minute extent, being produced, together with an equivalent amount
of hydroxidion by the reaction:

HOH ^s H- + OH'(orH + + OH~).

As stated on p. 213 the concentration of hydrion and hydroxidion
in pure water is 1 grm. molecule in 10 million litres, i.e. 10~ 7
normal. The concentration of an ion is expressed in physical chem-
istry by enclosing its symbol in square brackets; thus in water:

[H + ] = [OH-] = io- 7 and[H + ] [OH-] = i<r u

Now, according to the principles of chemical dynamics, in an
acid solution, in which hydrion concentration is increased, as well as
in an alkaline solution, in which hydroxidion concentration is
increased, hydrion and hydroxidion concentrations must stand in
geometrical ratio, so that their product has the same constant value
as in water, thus:

[H + ] [on'] = K = io- u

For example, if [H+] becomes 10 ~ 6 [OH~] correspondingly becomes
10 ~ 8 , and so on.

Since the value for pure water is [H+] = [OH~] = 10 ~ 7 , acidity
or alkalinity is inferred from the value of the negative index of
hydrion concentration alone, it being understood that hydroxidion
concentration necessarily has the complementary value.

The hydrion concentration of a solution is spoken of as its pW.
value (p = power in the mathematical sense), and this, as already
shown, is the negative value of the logarithm of the concentration.
Thus since for water [H + ] = 10 ~ 7 the pH value for water is 7; so
that this value corresponds with absolute neutrality in which hydrion
and hydroxidion concentration values are equal; whilst acidity is
indicated by values less and alkalinity by values greater than 7.

The determination of pH. values is important in bio-chemistry
and industrial chemistry, and much research has been done upon

APPENDIX 319

the subject. In general, two methods are available: colorimetric
and electrometric. In the colorimetric method various indicators or
mixtures of indicators are used, and colour charts are prepared
showing standards to be matched corresponding with various jr?H
values; but this method is applicable only to solutions originally
colourless. The electrometric method depends on the reading of
electropotential differences in solution and, being independent of
colour, can be applied to highly coloured liquids. For these methods
special textbooks must be consulted. 1

AMPHOTERIC HYDROXIDES

Since an amphoteric hydroxide is one which may behave either
as an acid or a base, it is one which in solution may provide either
hydrogen or hydroxide ions. Thus A1(OH) 3 undergoes ionization in
the two ways shown in the following scheme:

H- + A10(OH)' 2 (or A1OV H 2 O)

Representative of the two functions of A1(OH) 3 are aluminium
chloride, A1C1 3 , and sodium metaluminate NaAlO 2 , and the basic
and acidic strengths of the hydroxide are indicated by the resistance
to hydrolysis of these two salts respectively.

Now it has been shown by Wood (Chem. Soc. Trans., 1908, 93,
428) that while aluminium chloride in decinormal solution is hydro-
lysed to the extent of about 4 per cent at 25 C, being transformed
probably into the basic chloride A1(OH)C1 2 , sodium metaluminate
in decinormal solution is hydrolysed to the extent of about 35 per
cent at the same temperature. Therefore it appears that although
A1(OH) 3 is amphoteric, its basic exceeds its acidic strength.

Another amphoteric hydroxide studied by Wood (loc. cit.) is
arsenious hydroxide, whose acidic and basic functions are indicated
by the dual ionization:

H- + ILAsCV

OH' + H 2 AsO 2 -

It is known that in this case the acidic function predominates,
indeed, that arsenious hydroxide is an. acid comparable in strength
with boric acid.

1 e.g. The Theory and Use of Indicators. Prideaux.
The Determination, of Hydrogen Ions. Clark.

320 APPENDIX

It may be asked whether an ainphoteric hydroxide in solution,
such as arsenious hydroxide, ionizes in these two ways simulta-
neously. The answer is that it does not, for if the [H>] of its
solution exceeds that of water, the [OH'] must necessarily bo less
than that of water, so that, since the following relation is necessarily
true, whether in the case of water itself or of aqueous acid and
alkalis:

[H-] X [OH'] = K* = 1(T U ,

an amphoteric substance, dissolving in water to form a feebly acid
solution, produces no hydroxidion of its own but actually reduces
the [OH'] of water itself. Consequently, the essential thing about
an amphoteric hydroxide, which distinguishes it from compounds
which are always either acids or bases, is that it can produce either
anions or cations according to whether base or acid is added to its
solution.

IONIC AND ELECTRONIC EQUATIONS

Consider the simplest example of a chemical equation:

Zn + 2HC1 = Zn01 2 + H 2 .
Expressed in terms of ions this becomes:

Zn + 2H + + 2(Jl~ = Zn ++ + 2Cl~+H 2 .

The chloride ions on both sides need not be shown, so that the
equation becomes:

Zn + 2H + = Zn ++ + H 2 .

It thus appears that the reaction between zinc and hydrochloric
acid is the displacement of hydrogen ions by zinc owing to the
greater solution pressure of the metal; moreover, one zinc atom
displaces two hydrogen ions and appropriates their charges, to
become a bivalent ion carrying two + charges.

Now consider a more difficult example, the equation:
3Cu + 8HNO 3 = 3Cu(NO 3 ) 2 + 4H 2 O + 2NO

As shown on p. 309, such an equation is generally built up by
means of oxides. The result is sure, but the method is not very
convincing; for oxides, in this case CuO and N 2 O 6 , never come into
existence, and have to be apologized for.

Moreover, as above, the existence and reactivity of ions
rather than molecules must be recognized, as follows:

3Cu + 8H + + 8NO 3 ~ = 3GV" 1 " + 6NO 3 " + 4H 2 O + 2NO;

APPENDIX 321

and this, by removal of 6NO 3 ~ common to both sides, becomes:
3Cu + 8H + + 2lSr<V = 3Cu ++ + 4H 2 O + 2NO.

It will be observed that in this simplified equation there are
neutral atoms or molecules on both sides, together with charged
ions; and that the electric charges balance, the algebraic sum on
both sides being in this case 6 +.

If we look below the surface for the reason of this reaction, it
seems at first sight that six of the H-ions transfer their charges
to three Cu-atoms so as to ionize them; yet this cannot be the way
in which the reaction begins, for if chloride or sulphate ions were
present in place of nitrate ions such a transfer would not occur.
The inception of the reaction must therefore lie between the
nitrate ions and the metallic copper. 1

In the former method of building the equation it was recognized
that nitrogen is reduced from the condition represented by N 2 O 5 to
that represented by NO. Now without symbolizing N 2 O 5 its
reduction may be represented by employing the idea of valency.
By reduction from the quinquivalent to the bivalent state each of
the two nitrogen atoms in 2NO 3 ~ loses 3 positive valency units,
these 6 units being transferred to 3 copper atoms, converting them
into ions; 4 oxygen atoms are consequently liberated as ions with
double negative charges, the two further charges required for them
being already available from the singly charged nitrate ions.
These 4 oxygen ions then neutralize 8 hydrogen ions forming
4 water molecules. The NO molecules remain uncharged, so that
the whole procedure is represented thus:

+ 2NO 3 ~~ =

4O" + 8H + = 4H 2 O.

Then adding:

3Cu + 8H* + 2NO 3 "~ = SCu" 4 " + + 4H 2 O + 2NO, as before.

Here is the idea of valency transfer from atom to atom, but this
may be carried further, for is not valency transfer electron transfer?
Decrease of positive valency is gain of electrons, decrease of
negative valency loss of electrons, and vice versa. Thus if E
symbolizes an electron:

1 As pointed out on p. 291, it appears that the initiation of the reaction is due to a trace
of nitrous acid; yet this does not invalidate the main equation.

322 APPENDIX

2NO 3 ~ + 6E = 2NO + 4O =

3Cu-6E = 3Cu ++
and 4CT + 8H + = 4H 2 O, as before.

Therefore the whole reaction is reducible to the transfer of 6
electrons from 2 nitrate ions to 3 copper atoms, with necessary
consequences.

When equations are built up by considering oxides, this is
tantamount to saying that the reactions involved are examples
of oxidation with corresponding reduction. Thus in the above
example the metallic copper may be said to be oxidized to the
the cupric state while nitric acid is reduced to the state of nitric
oxide. From this, two related conclusions follow:

Oxidation = gain of + charges and valency = loss of electrons.
Reduction = loss of + charges and valency = gain of electrons.

Clearly, then, oxidation and reduction are reciprocal, and consist
essentially in the transfer of electrons from the reducing to the
oxidizing agent, from the substance oxidized to the substance
reduced.

On p. 297 the following reaction is cited to show the reciprocity
of oxidation and reduction:

6FeS0 4 + 3H 2 S0 4 + 2HN0 3 = 3Fe 2 (SO 4 ) 3 + 4H 2 O + 2NO.
Written in ions and halved this equation becomes:

3Fe ++ + 4H + + NO 3 ~ = 3Fe +++ + 2H 2 O + NO,
and written in parts to show electron transfer:

NO 3 ~ +3E = NO + 2O"
3Fe ++ - 3E = 3Fe + ++
and 2O == + 4H + = 2H 2 O.

Consider the well-known equation showing oxidation of ferrous
sulphate in acid solution by potassium permanganate:
2KMnO 4 + 10FeSO 4 + 8H 2 SO 4 = K 2 SO 4 + 2MnSO 4 + 5Fe 2 (SO 4 ) 3 + 8H 2 O.
Expressed in terms of ions this bec'omes:

2MnO 4 ~ + 10Fe + + + 16H + = 2Mn ++ + 10Fe + + + + 8H 2 O,
or halving:

Mn0 4 ~ + 5Fe ++ + 8H + = Mn ++ + 5Fe +++ + 4H 2 O.
To understand this reaction it is necessary to recognize that the

APPENDIX 323

Mn-atom is reduced in valency from 4- 7 to + 2 by an accession of
five electrons from iron thus:

MnO 4

6Fe ++ -5E =
4CT +8H + = 4H 2 O

_
Adding MnO 4 ~ +5Fe ++ + 8H + = Mn ++

It will be observed that the ionic charges balance, there being
17 + on both sides of the equation.

Consider next the oxidation of cuprous sulphide to cupric
nitrate and sulphur by the reaction:

= 6Cu(NO 3 ) 2 + 3S + 4NO + 8H 2 O,

or, expressed in ions:

Electrons are lost or gained as follows:

6Cu + -6E = 6Cu ++
SS^-GE = 38

4NO 3 ~+12E == 4NO + 8O
= 8H,O.

Adding eliminates electrons, and gives the above ionic equation.

If it be asked how such an ionic equation can be built up, the
following is the answer:

The reactants and products must be known, i.e. Cu 2 S, HNO 3 ,
Cu(N0 3 ) 2 , S, NO.

Then the following electronic equations can be written:

2E = 2Cu ++ (i)

S = - 2E = S (ii)

NO 3 ~" + 3E = NO + 2O = (Hi)

4E are required, and NO 3 - provides 3E; therefore 12 E must be
transferred in the course of the reaction; so, multiply (i) and (ii)
by 3, and (iii) by 4, and add, eliminating E, also supplying sufficient
H+ to combine with O=, and the equation becomes:

One other example will suffice: the reaction representing the

324 APPENDIX

oxidation by alkali hypochlorite of the arsenic deposited in Marsh's

test:

5NaOCl + 2As + 6NaOH = 5NaCl + 2Na 3 AsO 4 + 3H 2 O;

or 5O(jr+2As + 60H~ = 5C1~ + 2AsO 4 = +3H 2 O.
To build the equation:

As-5E = As +++++ ; OC1~+2E = C1"+O = ;
or 2As-10E = 2As ++++ ~ h ; 5OC1" + 10E = 5(jr+50 = .

2As++ ++ + requires 8O = of which 5 are obtained from 50C1~;
the remaining 3 come from OH" ions thus:

6OH" = 3
Consequently the equation is built thus;

5OCT + 10E = 5
2As-10E =

60H" = 3H 2

Adding: 50C1" +2Aa + 6OH" = 5U1

The student will recognize that equations built in this manner
by the use of modern electronic theory give a new insight into the
nature of chemical change,

INDEX

a-particles, 95.

Aoegg, in.

Absorptiometer, 191.

Acid, definition of, 237.

Acidic ions, 218.

Acidic oxides, 238.

Acids and metals, interaction of, 288.

Acids, bases, and salts, 236.

Active mass, 260.

Affinity, chemical, 260.

Affinity, units of, 59.

Air, liquefaction of, 156.

Air, liquid, 153.

Allotropic elements, list of, 187.

Allotropy, 182.

Allotropy, definition of, 183.

Ammonia, 51.

Ammonia, liquid, 154.

Ammonium chloride, dissociation of, 275.

Ammonium hydroxide, 281.

Amorphous state, 172.

Amphoteric hydroxides, 319.

Amphoteric oxides, 239.

Andrews, 146, 149.

Anions, 209.

Anode, 209.

Arrhenius, 98.

Aston, 106.

Atom as a planetary system, 1 10.

Atom, modern view of, 93.

Atomic nucleus, no.

Atomic number, 101.

Atomic sheath, 112.

Atomic theory, 4, 17.

Atomic theory, Dalton's, 9.

Atomic value, 56.

Atomic volume, 78.

Atomic volume curve, 79.

Atomic weight, 18, 54.

Atomic weight standards, 20.

Atomic weight values, correction of, 86.

Atomic weights, Dalton's, ip.

Atomic \veights, determination of, 24.

Atomicity, 15.

Atoms, 13.

Atoms, Greek, 2.

Auto-oxidation, 304.

Avogadro, 14.

Avogadro's theory, 15, 17.

Avogadro's theory, method of, 25.

825

/3-particles, 95.

Barium peroxide, dissociation of, 277.

Base, definition of, 237.

Basic ions, 218.

Basic oxides, 236.

Berthollet, 5.

Berzelius, 16, 70, 94, 118, 120.

Bohr, 114.

Boiling-points of liquids, 164.

Bonds, 59.

Boyle, 3.

Boyle's law, deviations from, 143.

Bragg, Sir W. H., 128.

Bragg, W. L., 128.

Bredig, 225.

Bury, 114.

Calcium carbonate, dissociation of, 276.

Cannizsaro, 16.

Carbon and silicon compared, 130.

Carbon, atomic weight of, 37.

Carbon dioxide, liquid, 155.

Carbon monoxide, 52.

Carbonates, 246.

Carbonic anhydride, 49.

Cascade method of cooling, 149.

Catalysis, 266.

Catalysis, theory of, 270.

Catalyst, definition of, 271.

Cataphoresis, 229.

Cathode, 209.

Cations, 209.

Chemical affinity, 260.

Chemical change, 254.

Chemical change, limits of, 265.

Chemical change, rate of, 265.

Chemical changes classified, 273.

Chemical combination, laws of, 16.

Chemical compounds, types of, 233.

Chemical displacement, 32, 54.

Chemical equilibrium, 257.

Chemical nomenclature and symbols, 315.

Chemical properties, periodicity of, 81.

Chemical reactions in solution, 218.

Chemistry in space, 65.

Chlorine, liquid, 155.

Claude, 153.

Colloidal state, 223.

Colloids, 224.

Colloids, classification of, 230.

326

INDEX

Colloids, gradation of, 228.

Colloids, protective, 229.

Coloured ions, 218.

Colours of salt solutions, 218.

Combustion, heat of, 262.

Complex ions, 219.

Complex salts, 251.

Compound gases, compositions of, 47.

Compounds, names of, 316.

Conditions of oxidation, 298.

Conditions of reduction, 298.

Constitutional formulae, 60.

Covalency, 120.

Co valency, illustrations of, 120.

Critical data, 148.

Critical density, 147.

Critical pressure, 147.

Critical state, 146.

Critical temperature, 146.

Critical volume, 147.

Crookes, 94, 105.

Cryohydrates, 173.

Cryoscopic method, 42.

Crystal, definition of, 176.

Crystal systems, 179.

Crystalline state, 172.

Crystallization, 181.

Crystalloids, 224.

Crystallography, 175.

Crystallo-hydrates, dissociation of, 277.

Crystals, 175.

Dalton, 6, 93, 137.

Dalton and Henry's law, 196.

Davy, H., 67, 219.

Decomposition, thermal, 278.

Definite proportions, law of, 5.

Deliquescence, 278.

Dephlegmator, 200.

Dewar vessel, 152.

Dialysis, 224.

Diffusion of gases, 137.

Dilute solutions, properties of, 208.

Disperse phase, 229.

Dispersion medium, 229.

Dissociation, electrolytic, an.

Dissociation pressure, 258.

Dissociation, thermal, 258, 274.

Distillation, fractional, 199.

Distillation under reduced pressure, 163.

Doberetner, 70, 137.

Double bond in carbon compounds, 64.

Double salts, 249.

Dualistic system, 118.

Dulong and Petit, 32.

Dumas, 27, 120.

Ebulliscopic method, 43.
Efflorescence, 278.
Eka-aluminium, 85.
Electrodes, 208.
Electrolysis, 208.
Electrolysis, laws of, 215.
Electrolyte, 208.
Electrolytic dissociation, 211.
Electron, 94.
Electro-valency, 120.
Element, 16.
Element, definition of, 3.
Elements, classification of, 69.
Elements, Greek, 2.

Elements, melting-points of, 170.
Elements, names of, 315.
Elements, symbols of, 316.
Enantiotropic change, 185.
Endothermic change, 263.
Endothermic compounds, 263.
Energy free, latent, total, 261.
Enzymes, catalytic influence of, 270,
Equation building, 305.
Equation, thermochemical, 261.
Equations, ionic and electronic, 320.
Equilibrium, chemical, 257.
Equivalent weight, 1 8, 54.
Equivalent weights, determination of, 21.
Ethylene, 53.
Eutectic mixture, 173.
Exothermic change, 263.
Exothermic compounds, 263.

Faraday, 94, 149, 215, 225.

Finely divided metals, catalytic influence of>

268.

Fixed proportions, law of, 5.
Formation, heat of, 263.
Fractional distillation, 199.
Fractionating column, 200.
Fusible metal, 175.

Gallium, 85.

Gas laws, 135.

Gaseous diffusion, 137.

Gaseous diffusion, law of, 138.

Gaseous mixtures, 141.

Gaseous mixtures, solubilities of, 196.

Gases, liquefaction of, 147.

Gases, properties of, 134.

Gases, solubilities of, 191.

Gay-Liissac, n.

Gel, 225.

Graham, 140, 225.

Graphic formulas, 59, 60.

Greek elements and atoms, 2.

Hafnium, 102.
Halides, 241.
Heat, action of, 274.
Heat of combustion, 262.
Heat of formation, 263.
Helium, liquefaction of, 154.
Henry's law, 195.
Hess, law of, 263.
Hofmann, 27.
Hydrated salts, 248.
Hydrides, 233.
Hydrion concentration, 318.
Hydrogel, 225.
Hydrogen chloride, 48.
Hydrogen ions, catalytic influence of, 267.
Hydrogen, liquefaction of, 153
Hydrogen sulphide, 49.
k Hydrolysis, 247, 295.
Hydrosol, 225.
Hydroxides, decomposition of, 280.

Indicators, theory of, 220.

Inorganic salts, constitutions of, 127.

Ionic and electronic equations, 320.

lonization, 212.

Ions, 209.

Ions, complex, 219.

Isobares, 99.

Isomorphism, 35.

INDEX

327

Isomorphism, law of, 55.
Isothermals of carbon dioxide, 145.
Isotopes, 100.

Joule-Thomson effect, 152.
Kopp, 1 60.

Ladenburg, 140.

Landsberger apparatus, 44.

Langmuir, 113.

Langmuir's postulate, 113.

Lavoisier, 3, 254.

Law of Boyle, 135.

Law of Charles, 135.

Law of Dalton and Henry, 196.

Law of definite proportions, 5.

Law of Dulong and Petit, 33.

Law of fixed proportions, 5.

Law of gaseous diffusion, 138.

Law of Henry, 195.

Law of Hess, 263.

Law of Isomorphism, 34, 55.

Law of multiple proportions, 6.

Law of octaves, 72.

Law of partial pressures, 141.

Law of reciprocal proportions, 7.

Law of specific heats, 54.

Law of volumes, n, 17.

Laws of chemical combination, 16.

Laws of electrolysis, 275.

Lead, atomic weights of, 109.

Le Chateher, 260.

Lezm, G. N , 112.

Limits of chemical change, 265.

Linde and Hampson, 152.

Liquefaction of gases, 147.

Liquefied gases, practical applications of,

Liquid :
Liquids
Liquids
Liquids
Liquids
Liquids

*ir, 153-
boiling-points of, 161.
densities of, 159.
molecular volumes of, 159.
properties of, 158.
solidification of, 168.

/ 1 . _

Liquids, specific volumes of, 159.
Liquids, vapour pressures of, 161.
Litmus, 221.
Long periods, 77.
Lowry, 271.

Mass, active, 260.

Matter, composition of, i.

Mendeleeff, 73, 93.

Mercuric oxide, dissociation of, 254.

Mercurous chloride, dissociation of, 276.

Metals and acids, interaction of, 288.

Metals and nitric acid, 290.

Meta-stable state, 186.

Methane, 53.

Methyl orange, 221.

Meyer, Victor, 27.

Mitscherlich, 35.

Mixed anhydrides, 240.

Mixtures, solidification of, 173.

Modern view of the atom, 93.

Molecular complexity, 45.

Molecular compositions of compound gases,

Molecular depression, 41.
Molecular elevation, 41.

Molecular formulae, 128.
Molecular theory, 1 1 .

Molecular weights in solution, determina-
tion of, 39.

Molecule, modern view of, 118.
Molecules, 13.
Monotropic change, 186.
Moseley, 101.
Multiple proportions, law of, 6.

Neutral oxides, 234.

Newlands, J. A. R., 71.

Newton, 7, 93.

Nitric acid and metals, 290.

Nitric oxide, 50.

Nitrogen peroxide, dissociation of, 274.

Nitrous oxide, 50.

Nomenclature, chemical, 315.

Nucleus, atomic, no.

Octaves, law of, 72.

Octet, 120.

Octet, theory illustrations of, 125.

Organic compounds, melting-points of , 171.

Osmotic pressure, 212.

Ostwald, 270.

Ostwald, Wo., 230.

Oxidation, 296.

Oxidation, conditions of, 298.

Oxidation in solution, 299.

Oxidation in the dry way, 298.

Oxides, 234.

Oxides, acidic, 238.

Oxides, amphoteric, 239.

Oxides and salts, catalytic influence of, 269.

Oxides, basic, 236.

Oxides, decomposition of, 278.

Oxides, neutral, 234.

Oxides, saline, 240.

Oxidizing agents, 297.

Oxy-salts, 245.

Oxy-salts, decomposition of, 282.

Partition co-efficient, 198.

Peptization, 229.

Periodic law, 73.

Periodic law, method of, 37.

Periodic law modern form, 103.

Periodic law, objections to, 90.

Periodic law, suggestiveness of, 87.

Periodic law, uses of, 84.

Periodic table, 74, 75.

Periodic table modern form, 104.

Periodicity of chemical properties, 81.

Periodicity of physical properties, 78.

Periodicity of valency, 82.

Peroxides, 241.

Phenol-phthalein, 221.

Phosphine, 52.

Phosphorus pentachloride, constitution of,

131-
Phosphorus pentachloride, dissociation of,

275-

/>H value, 318.

Physical properties, periodicity of, 78.
Pictet, 150.
Planck, 115.
Polymerism, 184.
Polymorphism, 182.
Poly oxides, 241.
Positive ray analysis, 106.

328 INDEX

Pressure, osmotic, 2x2.
Priestley, 137, 254.
Protective colloids, 229.
Protons, 96.
Proust, 5.
Prow*, 70, 93.

Quantum theory, 115.

Radioactive change, 98.

Ramsay, Sir William, 95, 97.

Raoult's law, 40, 55.

Rate of chemical change, 265.

Reactions in solution, 218.

Reactions, reversible, 258.

Reciprocal proportions, law of, 7.

Recrystallization, 182.

Reducing agents, 297.

Reduction, 296.

Reduction, conditions of, 298.

Reduction in solution, 299.

Reduction in the dry way, 298.

Reversible reactions, 256.

Richter, 7.

Rontgen, 95.

Rutherford, Sir Ernest, 96.

Rydberg, 103.

Rydberg series, 103.

Saline oxides, 240.

Salt solutions, colours of, 218.

Salts, complex, 251.

Salts, double, 249.

Salts, hydrated, 248.

Saturation, 201.

Saturation capacity, 57.

Self-intensive refrigeration, method of, 151

Sheath, atomic, 112.

Short periods, 77.

Siedentopf, 226.

Simple compression, method of, 148.

Slightly soluble salts, 206.

Smith, J. D. Main, 1 17.

Soddy, 95, 100.

Sodium chloride, crystal unit of, 129.

Sol, 225.

Solids, formation of, 165.

Solids, formation of, from solution, 172.

Solids, melting-points of, 168.

Solids, properties of, 166.

Solubilities, table of, 204.

Solubility, 201.

Solubility and chemical composition, 205.

Solubility, co-efficient of, 190.

Solubility curves, 204.

Solubility, definition of, 201.

Solubility, effect on chemical change, 292.

Solute, 189.

Solution, process of, 203.

Solutions, 189.

Solutions of gases in liquids, 190.

Solutions of liquids in liquids, 197.

Solutions of solids in liquids, 200.

Solvent, 189.

Space lattice, 129.

Specific heats, 33.

Specific heats, law of, 54.

Stoney, Johnstone, 94.

Strecker, 71.

Sublimation, 167.

Suboxides, 236.

Sulphides, 244.

Sulphion, 209.

Sulphur dioxide, liquid, 155.

Sulphuric acid, constitution of, 126.

Sulphurous anhydride, 49.

Superoxides, 241.

Supersaturation, 201.

Symbols, chemical, 316.

Temperature and solubility, 203.
Theory of indicators, 220.
Thermal decomposition, 278.
Thermal dissociation, 258, 274.
Thermochemical equation, 261.
Thermochemistry, 280.
Thomson, SirJ.J.,*)*.
Transition temperature, 185.
Triads, 70.
Tyndall, 226.
Types of chemical compounds, 233.

Ultramicroscope, 226.
Unitary system, 119.
Unsaturation, 201.

Valency, criterion of, 66.

Valency, nature of, 67.

Valency, newer views of, 120.

Valency, older views of, 56.

Valency, periodicity of, 82.

Valency, statement of, 57.

Valency, units of, 59.

Valency, variability of, 61.

Van't Hoff, 66.

Vapour density, determination of, 27.

Vapours, solidification of, 166.

Victor Meyer, 27.

Volatility, effect on chemical change, 292.

Water and steam, 49.

Water, catalytic influence of, 267.

Water, chemical action of, 284.

Werner, A., 77.

Wroblewski, 153.

X-ray spectra, 101.
X-ray spectrography, 126.
X-rays, 95.

Yttrium, 105.
Zsigmondy, 226.

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