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The Elements of 
^Physical and General Chemistry 



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

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





By R. M. Caven, 'D.Sc., F.I.C. 

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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 

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; 


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 

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 

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- 

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 

GLASGOW, Uctobti , 


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 

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



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. 

GLASGOW, October, 1925. 


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. 

GLASGOW, February, 1929. 



CHAP. Page 


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


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. 


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. 


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. 


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. 



CHAP. Page 


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. 



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. 


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. 


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. 


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. 


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


CHAP. p age 


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 



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


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


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. 


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


Chemical Nomenclature and Symbols 315 

Hydrion Concentration and pH. Value 318 

Amphoteric Hydroxides 316 

Ionic and Electronic Equations 320 

INDEX 325 




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 


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 

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 

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 


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 


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 

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. 


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 

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 


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 

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: ^ 


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 

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 



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- 

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 


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: 




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. 




Hydrogen ... 


An atom of water or steam, 



composed of 1 of oxygen 

Carbon or Charcrn 



+ 1 of hydrogen 




T*l 1 



An atom of ammonia, com- 

Magnesia ... 



posed of 1 of azote + 1 
of hydrogen 







An atom of carbonic oxide, 
composed of 1 of carbon 

JLILHl . . 




+ 1 of oxygen 




A 4- e v. 'i 



An atom ot caroomc aciclj 
1 carbon + 2 oxygen 




An atom of sulphuric acid, 

Mercury ... 


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 

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 


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 


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. 


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 

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 



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: 








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, 



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. 









n 2 






N 2 




















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. 


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 

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. 


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? 


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 

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 

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 


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 

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. 



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 

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 


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. 


2. Determination of Equivalent Weights 

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

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 


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 


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 

008 g 


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. 


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 

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 


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: 






Approximate Density) 
(0 = 16) / 






Approximate Molecular \ 
Height f 






Molecular Proportion of \ 
Carbon / 







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 


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 


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. ^ 


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. 


(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: 


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. 


Weight of air displaced when sealed globe is weighed 

_ 0-001293 X 76 X 273 


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 

0-205 = 37.!. 


(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 



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, 

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. 


= 131-8. 

: E 




(c) Victor Meyer's Method of Vapour-density 

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. 


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. 

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, 


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. 


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 

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. 



Specific Heat. 

Atomic Weight. 

Atomic Weight 
x Specific Heat 
=3 Atomic Heat. 





Lead ... 


207 2 


Gold ... 








Tin .. 




Silver .. 













Nickel .. 








Cobalt .. 








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 



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

(D60) 4 


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. 


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 


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 


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- 


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. 


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. 


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 


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. 



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 


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

100 KS 


100 X 50 X 0-317 


i. M = 

ii. M = 


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 


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 

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. 


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 


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 

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. 


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 


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 


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 

(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 


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. 


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 . 


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. 


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. 


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 . 


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. 


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. 


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. 


(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. 

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. 

specific heats of the solid elements are inversely proportional to 


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 

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

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. 



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. 



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 

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. 


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 

The following hydrides exhibit the valency of a number of 

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: 


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 







BC1 3 

CC1 4 

PF & 


ZnCl 2 

PC1 3 

SiCl 4 

AsF fi 



A1C1 3 

SnCl 4 

SbF 6 

Valency 1 








Na 2 


B 2 3 

CO 2 


S0 3 


OsO 4 

K 2 O 


A1 2 S 

SiO 2 

P 2 6 

Cr0 3 

(I 2 7 ) 

Eu0 4 

Ag 2 


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. 


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 

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 

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 


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. 



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. 


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: 



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 
OH 2 
FH - 

SiH 4 SiO 2 
PH 3 PA 
SHo SO 3 

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. 


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 

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. 



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 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, 



Pig 9 


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 

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. 


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. 


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. 


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


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 


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, 

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


Atomic Weight. 


Mean of Extreme 
Atomic Weights. 



i f if\ 




It)' 1U 











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

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


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 

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, 

H Li Be B C N O 

F Na Mg Al Si P S 

Cl K Ca Cr Ti Mn Fe, <kc. 


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; 


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 

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 

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 






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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. 












A B 

A B 

A B 

A B 

A B 

A B 

A B 


Series 1 





























Fe Co Ni 

















Hu Ilh Pd 



























_ _ 

Os Ir Pt 













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 


(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 


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 

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: 


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 


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. 


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- 

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 

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 : 



1 F 

S 3 IN fllOA O I W OJL V 


the atomic volume is this value multiplied by the atomic weight 

Atomic volume = fttomic weight 

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

Atomic volume Cu = ?1? = 7-15. 

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 


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 

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- 

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 

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 


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 


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: 








Li 2 O 



C0 2 



B 2 H fl , <fec 

. CH 4 

NH 3 

OH 2 



Na 2 O 



SiO 2 


S0 3 



SiH 4 

PH 3 

SH 2 



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. 


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 


and ZnO 


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. 



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. 



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 

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. 


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 


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 


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 


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

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 

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 


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. 


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 


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. 


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. 


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 



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 


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 


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. 


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 


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 


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 


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. 


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. 


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 

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 properties of the elements are periodic functions of their atomic 

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. 
























) <M 



ee n 















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- 


graphic plate in positions which depend only on their individual 

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 



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




of Isotopes. 

Mass Numbers of 
Isotopes in order 
of Intensity. 





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





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





39, 41 





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 


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 



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: 



Ordinary Lead. 

Atomic weight 
1 )ensity 
Atomic volume 




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 


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 

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 



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 








+ 6 



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. 


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 

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 

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. 



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 


Ti 2 3 
(Ti0 2 ) 



V 2 3 
V 2 4 


Cr 2 3 

Cr0 3 

Mn 2 O 3 
MnO 2 

MnO 3 
Mn 2 O 7 

Fe 2 3 

FeO 3 

Co 2 3 


(Cu 2 0) 

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 











1 Science, July, 1921. 



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) 





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. 


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 

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. 


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. 


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. 


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 



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 



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 



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: 



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. 



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 



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 

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 


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: 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 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: 






















the valencies of the separate atoms, and of the ions, their algebraic 
sum, being: 


Si = +4 P = +5 
O = -8 4O = -8 


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 






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: 


+ o=s=o + 


this reaction now becomes: 
:C1: H:6: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- 

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 


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. 


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 


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. 


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 

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 


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: 



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. 




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. 



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 

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 


(= 0*00367) part of its value at C. for every degree rise of 

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 


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 



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. 


Density : 
Air = 1. 


V density ' 

Velocity of Dif- 
fusion : Air = 1. 









Carbon monox 

















Hydrogen sulphide 
Nitrous oxide 




Carbon dioxide 




Sulphur dioxide 




*/ 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. 


and this aspect of the law is illustrated by the following results of 


Time of Molecular 
Passage into Air 
at 100 mm. Pressure. 

V Density : 
Oxygen = 1. 

Time of Mole- 
cular Passage into 
a Vacuum. 

Carbon dioxide ... 






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 

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 



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 

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- 

(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). 


(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: 


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 


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 


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. 


B.-P. Atmospheric 





(-253 C.) 




(-196 C.) 




(about -190C.) 



Carbon monoxide 

(-190 C.) 




(-183 C.) 



Nitrous oxide 

(-89-8 C.) 




(-33-5 C.) 


1 -00632 

1 Lord Rayleigh, Proc. Roy. Soc. t 1905, 74, 446. 




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. 


Fig. 14 


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 




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 


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. 


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 


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 



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. 


Sulphur dioxide 



Critical Pressure. 


B.-P. under 760 mm. 






Hydrogen sulphide 



<;i -8 

Hydrogen chloride 


. .. . 81-6 


Nitrous oxide 

36 5 






sublimes at 83-6 

Carbon dioxide 







Nitric oxide 
Carbon monoxide 



-252 -5 

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 


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; 



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 


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 


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 



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 


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, 



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 


from the surroundings; so water, placed within the cup, may be 

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 


coil in 
Brine Tank 


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. 


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 

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. 



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 

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 

The critical volume is the specific volume of the fluid in its 
critical state. 


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. 


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, 



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 


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 

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. 


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 

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 


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 

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 



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 



& 900 

.9 800 

I 70 


10 20" 30 40 50 60 70 80" 90" 100*110 120" 130* 

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, 



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- 

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 



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. 


C r H 16 


normal Hydrocarbons, 

B.-P. C. Diff. 


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 
B.-P. 0. 













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. 



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 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 

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 

The boiling-points of successive members of a homologous series 
of chemical compounds rise regularly with successive increments of 
molecular weight 


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. 


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 




tube it yields crystals of iodine without passing through the liquid 

In the former case the iodine behaves thus: 




in the latter case thus: 




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 

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 


Temp. C. Press, mg. Hg. 

58-1 . 


75-2 . 


91-9 . 


102-7 . 



) 90-0 

120-4 . 


125-5 . 




177-6 . 


184-35 (b. p 

) 760-0 


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. 


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 

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 

The influence of pressure upon melting-point is shown for ic& 
and phosphorus in the following table: 



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. 


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.). 



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. 






Helium be 

ow 271 



TTvrl rofifpn 










under pressure. 














- 38-87 




- 7-3 



















Sulphur (rhombic) 




Sulphur (monoclinic) 







2200 8 -2500 
























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 


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: 

CH 3 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 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 

- 17-9 
- 5 
+ 7 

+ 18-6 

The first few members of the series show exceptional relation- 


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. 

(c) Formation of Solids from Solution Crystalline and Amorphous States. 

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 


separate within it, just as crystals may separate within a liquid 

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 



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 

In the following table are given a few data regarding cryo- 


Cryohydric Point, 

Percentage of 
Anhydrous Salt in 

Sodium chloride... 



Potassium iodide 



Sodium nitrate ... 

-17 -5 r 


Ammonium sulphate . 
Ammonium chloride 



Magnesium sulphate . 
Sodium sulphate 

- 5 
- 0-7 


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 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 

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: 






<3 "0 


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. 


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. 



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 



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 



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 

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. 



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. 


Hemihedral forms: calcite, corundum, graphite. 

NOTE. The hemihedral forms of the hexagonal system are 
often classified as a separate system, the Bhombohedral System 


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 

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- 


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 


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, 


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 

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 

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. 


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 


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: 

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 


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 

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.); 

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. 



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 ^ 
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. 


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. 


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 

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 

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 



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. 


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 


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 

The pure, dry gas is contained over mercury in the graduated 



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 


by 1 volume 

Gas under 

of Water at C. 

Atmospheric Pressure . 



-252 -^ 

Carbon monoxide 






Nitrous oxide 


- 89-8^ 

Carbon dioxide 


- 80 

Hydrogen sulphide 


- 61-8 
- 33-7 

Sulphur dioxide 


- 10-1 

Hydrogen chloride 


- 83 



- 33-5 

It will be noticed that these gases fall conveniently into three 


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 

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. 


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 



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 





























There is a notable difference between hydrogen chloride and 
ammonia as regards effect of temperature on solubility. 






Temperature, < 
Centigrade. * 
























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 



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 



"3 < 




















) 10 20 30 40 50 60 

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, 


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. 


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. 



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. 





of Gas 

Carbon dioxide ... 





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 of Liquids in Liquids 

Pairs of liquids may be divided into three categories, according 
to their mutual solvent action. 

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. 

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 


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 

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 

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 

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 


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 



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 

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 


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. 


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 


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: 


Unsaturation : solid * solution. 


Supersaturation: solid - solution. 

Saturation: solid -^-^ solution. 

The state of equilibrium between a solid and a solution is 
thus analogous to the state of chemical equilibrium between the 


factors and products of a reversible chemical reaction, and to the 
state of thermal equilibrium between two bodies at the same 

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 



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. 


Tempei a- 

KC10 3 . 



Ca(OH) 2 . 


2 H 2 0. 

10 J 
























30 -3 














































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 


Fig. 42 

The solubilities of crystallized sodium sulphate appear to be 
anomalous. They are as follow: 






























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 


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. 

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.; 


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. 


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 

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. 


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 


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. 




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; 



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 2 !SO 4 2SO 4 
: + 2H 2 O 

= 2H 2 SO 4 



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. 


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 

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 


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. 


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. 


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. 


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, 



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: 





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? 


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 


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: 


Cu," blue. 
Co", pink. 
Ni", bright green. 
Cr'", purple. 
Mn", pale pink. 
Fe", green. 



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- 


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: 


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. 


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 

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 . 



ELECTROLYTE. An electrolyte is a substance which conducts an 
electric current whilst undergoing chemical decomposition which is 

An electrolyte, when dissolved in a suitable solvent, undergoes 
spontaneous dissociation into ions. This is electrolytic dissociation 
or ionization. 


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. 


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 



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. 


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 

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 


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- 

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. 


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 

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 


classified according to the sizes of their particles, with the following 


Suspensions of non-colloids. 

Colloidal platinum. 
Colloidal ferric hydroxide. 
Colloidal arsenious sulphide 
Colloidal gold. 
1 per cent gelatin. 
Colloidal silicic acid. 

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 


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 



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: 






Carbon particles in iron. Gold in ruby glass. 



Colloidal solutions of metals, gelatin, starch, &c. 



Smoke. Fine dust. Fumes. 



Certain minerals. 






Fog. Mist. Clouds. 



Solidified froths, e.g. pumice. 



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 

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." 


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. 


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- 

A sol passes into a gel by coagulation, a gel into a sol by 

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. 



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, 


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 


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 


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: 


This reaction involves rearrangement of the atoms concerned, i.e. 
intra-molecular rearrangement, probably in the following way: 


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 

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, 


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 

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 

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* 


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. 


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. 



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) 

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. 




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 

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. 


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, 

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 . 


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 

= Pb<g| and i8>PK8n 

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. 


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 


group, is only bivalent. The constitutions of these two oxides are 
therefore ordinarily represented thus: 


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): 


Pb0 2 . 
N0 2 



BaO 2 . 



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 




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 

Silicon tetrachloride, SiCl 4 , the chloride of a non-metal, may be 


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, 


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. 



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, 


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. 


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. 


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 


is. Consider, for example, the following series of chlorides and 

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. 


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 

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. 


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 


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. 


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. 


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 


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, 


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. 

Types of Chemical Compounds 

HYDRIDES. Metallic and non-metallic. 

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. 




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 



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 


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 



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 

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 





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 

360 C. 

Millimetres Hg Pressure. 


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. 



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. 

o 1000 
8 900 
| 800 
s 700 
g 600 


fi 500 
100 ( 


















360 380 400 420 440 460 480 

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- 


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. 


There is one other aspect of this simple chemical reaction between 


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 


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 


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 


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 

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. 


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. 


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. 



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: 


, * 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 


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 


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. 


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 

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 


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 

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 


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 ". 


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. 


It is customary to classify chemical reactions in the following 

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 


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 

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 


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 


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 

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. 


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 


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, 


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, 


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. 


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 


satisfied, the PbO in Pb 3 O 4 is not free to combine with more 

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 


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 



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. 


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, 


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- 


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 

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 


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. 


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 

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. 


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. 


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 


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, 


and produces an acid solution by incipient hydrolysis; when this 
solution is evaporated to dryness and ignited, the oxide A1 2 O 3 

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 

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. 


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 


solution, whilst acetic acid is a weak acid which is but slightly 
ionized, and therefore approaches water in its behaviour towards 

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 

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 

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, 


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- 

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. 


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 


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 

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, 


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 


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- 

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 

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. 


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 

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. 


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 

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 



always the case, however, for carbon is quadrivalent in both alcohol 

and aldehyde: 

H H 

CHy-i-OH, CH 3 -C=0, 

and barium is bivalent in both BaO and Ba0 2 : 



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. 


Free oxygen and ozone. 

Hydrogen peroxide and other per- 

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. 


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. 


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 


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 


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 

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. 


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 


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 

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. 


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 

Nascent hydrogen reduces an aldehyde thus: 
CH 3 -COH + 2H = CH 2 -CH 2 OH; 


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 


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. 


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. 


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 


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 


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 


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 


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 


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. 


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; 


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 

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 


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 


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. 


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. 



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 



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. 


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- 


posal has been made recently to standardize the prefixes for 
valency thus; 

Uni-, bi-, tri-, quadri-, quinque-, sexi-, septi-, octi-valent. 


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 


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 


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. 


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 

[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 


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; 


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. 



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 

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 

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 


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 


oxidation by alkali hypochlorite of the arsenic deposited in Marsh's 


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, 


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. 


/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. 



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> 


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. 



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 : 

*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, 

Phosphorus pentachloride, dissociation of, 


/>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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