ON THE DEVELOPMENT OF PHYSICAL CHEMISTRY SINCE 1850 171
Molecular volumes of liquids. Nature of solution. Van ’t
Hoff’s application of the gas laws to phenomena of solution.
Osmosis and osmotic pressure. Traube. Pfeffer. Semi-permeable
membranes. Measurement of osmotic pressure. Arrhenius.
Doctrine of ionisation. Its applicability to the explanation
of chemical phenomena. Thermo-chemistry. Mass action.
Nature of reversible reactions. Thermal and Electrolytic
dissociation. Relation between chemical nature and opticity,
magnetic rotation and viscosity. Theory of phases. Catalysis.
Enzyme action. Relations between valency and volume.
Photochemistry.
BIBLIOGRAPHY 187
INDEX 191
ILLUSTRATIONS
PAGE
JUSTUS VON LIEBIG _Frontispiece_
JEAN BAPTISTE ANDRÉ DUMAS 9
THOMAS GRAHAM 13
From a painting by G. F. Watts, R.A., in the possession
of the Royal Society
ALEXANDER WILLIAM WILLIAMSON 15
BUNSEN, KIRCHHOFF, AND ROSCOE 19
SIR WILLIAM RAMSAY 48
MARIE CURIE (_née_ SKLODOWSKA) 56
STANISLAO CANNIZZARO 65
SIR JAMES DEWAR 98
DMITRI IVANOWITSCH MENDELÉEFF 110
AUGUST KEKULÉ VON STRADONITZ 126
JACOBUS HENRICUS VAN ’T HOFF 142
EMIL FISCHER 166
SVANTE AUGUST ARRHENIUS 180
HISTORY OF CHEMISTRY
CHAPTER I
STATE OF CHEMISTRY IN THE MIDDLE OF THE NINETEENTH CENTURY
In the preceding volume an attempt was made to outline the significant
features in the development of chemistry, as an art and as a science,
from the earliest times down to about the middle of the last century.
Since that time chemistry has progressed at a rate and to an extent
unparalleled at any period of its history. Not only have the number
and variety of chemical products—inorganic and organic—been enormously
increased, but the study of their modes of origin, properties, and
relations has greatly extended our means of gaining an insight into
the internal structure and constitution of bodies. This extraordinary
development has carried the science beyond the limits of its own
special field of inquiry, and has influenced every department of
natural knowledge. Concurrently there has been a no less striking
extension of its applications to the prosperity and material welfare of
mankind.
With the death of Davy the era of brilliant discovery in chemistry,
wrote Edward Turner, appeared for the moment to have terminated.
Although the number of workers in the science steadily increased, the
output of chemical literature in England actually diminished for some
years; and, as regards inorganic chemistry, few first-rate discoveries
were made during the two decades prior to 1850. Chemists seemed to
be of Turner’s opinion that the time had arrived for reviewing their
stock of information, and for submitting the principal facts and
fundamental doctrines to the severest scrutiny. Their activity was
employed not so much in searching for new compounds or new elements as
in examining those already discovered. The foundations of the atomic
theory were being securely laid. The ratios in which the elements of
known compounds are united were being more exactly ascertained. The
efforts of workers, Graham excepted, seemed to be spent more on points
of detail, on the filling-in of little gaps in the chemical structure,
as it then existed, than in attempts at new developments. For a
time—during the early ’thirties—chemists struggled with the claims
of rival methods of notation, and it was only gradually that the
system of Berzelius gained general acceptance. At none of the British
universities was there anything in the nature of practical tuition in
chemistry. Thomson, at Glasgow, occasionally permitted a student to
work under him, but no systematic instruction was ever attempted. The
first impulses came from Graham in 1837, when he took charge of the
chemical teaching at the University of London, and when, in 1841, he
assisted to create the Chemical Society of London. Four years later
the Royal College of Chemistry in London was founded and placed under
the direction of August Wilhelm Hofmann—one of the most distinguished
pupils of Liebig. Under his inspiration the study of practical
chemistry made extraordinary progress, and discovery succeeded
discovery in rapid succession. In bringing Hofmann to England we had,
in fact, imported something of the spirit and power of his master,
Liebig.
Among the pupils and co-workers of Hofmann were Warren de la Rue, Abel,
Nicholson, Mansfield, Medlock, Crookes, Church, Griess, Martius, Sell,
Divers, and Perkin. Whilst at Giessen he had begun the study of the
organic bases in coal-tar with a view more especially of establishing
the identity of Fritzsche’s _anilin_ with the _benzidam_ of Zinin and
the _krystallin_ of Unverdorben. Hofmann continued to cultivate with
unremitting zeal the field thus entered. With Muspratt he discovered
_paratoluidine_ and _nitraniline_; with Cahours _allyl alcohol_. His
pupil Mansfield worked out, at the cost of his life, the methods
for the technical extraction of benzene and toluene from coal-tar,
and thereby made the coal-tar colour-industry possible. It was in
attempting to synthesise quinine by the oxidation of aniline that
Perkin, then an assistant at the college, obtained, in 1856, _aniline
purple_, or _mauve_, as it came to be called by the French, the first
of the so-called coal-tar colouring matters. In 1859 this was followed
by the discovery of _magenta_, or _fuchsine_, by Verquin. For its
manufacture Medlock, one of Hofmann’s pupils, in 1860 devised a process
by which for a time it was almost exclusively made. Hofmann studied the
products thus obtained, and showed that they were derivatives of a base
he called _rosaniline_; and he demonstrated that the colouring matters
were only produced through the concurrent presence of aniline and
toluidine. He also proved that the base of the dye, known as _aniline
blue_, was _triphenylrosaniline_. As the result of these inquiries he
obtained the violet or purple colouring matters known by his name.
Lastly, all his classical work on the amines, ammonium compounds, and
the analogous phosphorus derivatives was done at the Royal College of
Chemistry.
Prior to the establishment by Liebig, in 1826, of the Giessen
laboratory, the state of chemistry in Germany was not much, if at all,
better than with us. The creation of the Giessen school initiated a
movement which has culminated in the pre-eminent position which Germany
now occupies in the chemical world. Students from every civilised
country came to study and to work under its leader, and to carry away
with them the influence of his example, the inspiration of his genius,
and the stimulating power of his enthusiasm.
│Justus von Liebig│, was born at Darmstadt on May 12, 1803, and
after graduating at Erlangen, where he worked on the fulminates, he
repaired to Paris and entered the laboratory of Gay Lussac, with whom
he continued his inquiries. Returning to Germany, he was appointed
Professor of Chemistry at Giessen in 1826, and began those remarkable
series of scientific contributions upon which the superstructure of
organic chemistry largely rests. He investigated the _cyanates_,
_cyanides_, _ferrocyanides_, _thiocyanates_, and their derivatives.
In conjunction with Wöhler he discovered the group of the _benzoic
compounds_ and created the _radical theory_. With Wöhler also he
investigated _uric acid and its derivatives_. He discovered _hippuric
acid_, _fulminuric acid_, _chloral_, _chloroform_, _aldehyde_,
_thialdine_, _benzil_, and elucidated the _constitution of the organic
acids_ and the _amides_. He greatly improved the methods of organic
analysis, and was thereby enabled to determine the empirical formulæ of
a number of carbon compounds of which the composition was imperfectly
known. He practically laid the foundations of modern agricultural
chemistry, and to his teaching is due the establishment of an important
branch of technology—the manufacture of chemical fertilisers. He worked
on physiological chemistry, especially on the elaboration of fat, on
the nature of blood, bile, and on the juice of flesh. He studied the
processes of fermentation, and of the decay of organised matter. He was
a most prolific writer. The Royal Society’s Catalogue of Scientific
Papers enumerates no fewer than 317 contributions from his pen. He was
the founder of the _Annalen der Chemie_, which is now associated with
his name, and of the _Jahresbericht_; he published an _Encyclopædia
of Pure and Applied Chemistry_ and a _Handbook of Organic Chemistry_.
His _Familiar Letters on Chemistry_ was translated into every modern
language, and exercised a powerful influence in developing popular
appreciation of the value and utility of science. Liebig left Giessen
in 1852 to become Professor of Chemistry at the University of Munich
and President of the Academy of Sciences. He died at Munich on April
18, 1874.
With the name of Liebig that of Wöhler is indissolubly connected.
Although the greater part of their work was not published in
conjunction, what they did together exercised a profound influence on
the development of chemical theory.
│Friedrich Wöhler│ was born at Eschersheim, near Frankfort, on July 31,
1800. After studying at Marburg, where he discovered, independently of
Davy, _cyanogen iodide_, and worked on _mercuric thiocyanate_, he went
to Heidelberg and investigated _cyanic acid_ and its compounds, under
the direction of Gmelin. In 1823 he worked with Berzelius at Stockholm,
where he prepared some new tungsten compounds and practised mineral
analysis. In 1825 he became a teacher of chemistry in the Berlin Trade
School. Here he succeeded for the first time in preparing the metal
_aluminium_ and in effecting the _synthesis of urea_—one of the first
organic compounds to be prepared from inorganic materials. Jointly with
Liebig he worked upon _mellitic_ and _cyanic_ and _cyanuric acids_. In
1832 Wöhler, now appointed to the Polytechnic at Cassel, began with
Liebig their memorable investigation on _bitter-almond oil_. In 1836 he
was called to the Chair of Chemistry in the University of Göttingen,
and with Liebig attacked the constitution of _uric acid and its
derivatives_—the last great investigation the friends did in common.
Wöhler subsequently devoted himself mainly to inorganic chemistry. He
isolated _crystalline boron_, and prepared its _nitrides_, discovered
the spontaneously inflammable _silicon hydride_, _titanium nitride_,
and analysed great numbers of minerals and meteorites and compounds of
the rarer metals. He made Göttingen famous as a school of chemistry.
At the time of the one and twentieth year of his connection with the
University it was found that upwards of 8000 students had listened to
his lectures or worked in his laboratory. He died on September 23, 1882.
In France, Dumas exercised a no less powerful influence. If Liebig
could reckon among his pupils Redtenbacher, Bromeis, Varrentrapp,
Gregory, Playfair, Williamson, Gilbert, Brodie, Anderson, Gladstone,
Hofmann, Will, and Fresenius; Dumas could point to Boullay, Piria,
Stas, Melsens, Wurtz, and Leblanc—all of whom did yeoman service in
developing the rapidly expanding branch of organic chemistry.
[Illustration: JEAN BAPTISTE ANDRÉ DUMAS]
│Jean Baptiste André Dumas│ was born on July 14, 1800, at Alais, where
he was apprenticed to an apothecary. In his sixteenth year he went to
Geneva and entered the pharmaceutical laboratory of Le Royer. Without,
apparently, having received any systematic instruction in chemistry,
he commenced the work of investigation. With Coindet he established
the therapeutic value of iodine in the treatment of _goître_; with
Prevost he attempted to isolate the active principle of _digitalis_,
and studied the chemical changes in the development of the chick in
the egg. In his twenty-fourth year Dumas went to Paris and became
_Répétiteur de Chimie_ at the École Polytechnique. He joined Audouin
and Brongniart in founding the _Annales des Sciences Naturelles_, and
began his great work on _Chemistry Applied to the Arts_, of which
the first volume appeared in 1828. At about this time he devised his
_method of determining vapour densities_, and published the results
of a number of estimations made by means of it. With Boullay he
began an inquiry on the _compound ethers_, out of which grew the
_etherin theory_, which served as a stepping-stone to the theory of
compound radicals—subsequently elaborated by Liebig and Wöhler. Dumas
discovered the nature of _oxamide_ and of _ethyl oxamate_, isolated
_methyl alcohol_, and established the generic connection of groups of
similarly constituted organic substances, or, in a word, the doctrine
of _homology_. His work on the _metaleptic action of chlorine_
upon organic substances eventually effected the overthrow of the
electro-chemical theory of Berzelius and led to the theory of types,
which, in the hands of Williamson, Laurent, Gerhardt, and Odling, was
of great service in explaining the analogies and relationships of whole
groups of organic compounds. He worked in every field of chemistry.
He invented many analytical processes, established the _gravimetric
composition of water and of air_, and revised the _atomic weights_ of
the greater number of the elements then known. Dumas exercised great
influence in scientific and academic circles in France. He was an
admirable speaker, and had rare literary gifts. On the creation of the
Empire he was made a Senator, and was elected a member of the Municipal
Council of Paris, of which he became president in 1859. He died on
April 11, 1884.
It was largely through the influence of these master-minds that
chemistry took a new departure. Prior to their time organic chemistry
hardly existed as a branch of science: organic products, as a rule,
were interesting only to the pharmacist mainly by reason of their
technical or medicinal importance. But by the middle of the nineteenth
century the richness of this hitherto untilled field became manifest,
and scores of workers hastened to sow and to reap in it. The most
striking feature, indeed, of the history of chemistry during the past
sixty years has been the extraordinary expansion of the organic section
of the science. The chemical literature relating to the compounds
of carbon now exceeds in volume that devoted to all the rest of the
elements.
In the middle of the nineteenth century chemists began to concern
themselves with the systematisation of the results of the study of
organic compounds, and something like a theory of organic chemistry
gradually took shape. From this period we may date the attempts
at expressing the internal nature, constitution, and relations of
substances which, step by step, have culminated in our present
representations of the structure and spatial arrangement of molecules.
In 1850 the dualistic conceptions of Berzelius ceased to influence
the doctrines of organic chemistry. The enunciation by Dumas of the
principle of substitution, and its logical outcome in the nucleus
theory and in the theory of types, had not only effected the overthrow
of dualism, but was undermining the position of the radical theory of
Liebig and Wöhler. The teaching of Gerhardt and Laurent had spread
over Europe, and was influencing those younger chemists who, while
renouncing dualism, were not wholly satisfied with a belief in compound
radicals. Williamson’s discovery, in 1850, of the true nature of ether
and of its relation to alcohol, and his subsequent preparation of mixed
ethers, served not only to reconcile conflicting interpretations of
the process of etherification, but also to reconcile the theory of
types with that of radicals. Lastly, his method of representing the
constitution of the ethers and their mode of origin gave a powerful
stimulus to the use of type-formulæ in expressing the nature and
relations of organic compounds.
[Illustration: THOMAS GRAHAM.
From a painting by G. F. Watts, R.A., in the possession of the Royal
Society.]
Other representative men of the middle period of the nineteenth
century, in addition to Williamson, were Graham and Bunsen. The
three men were investigators of very different type, and their work
had little in common. But each was identified with discoveries of a
fundamental character, constituting turning-points in the history of
chemical progress, valuable either as regards their bearing on chemical
doctrine or as regards their influence on operative chemistry.
│Thomas Graham│ was born in Glasgow on December 21, 1805, and, after
studying under Thomas Thomson at the University of that city, attended
the lectures of Hope and Leslie in Edinburgh. In 1830 he succeeded Ure
as teacher of chemistry at Anderson’s College in Glasgow, and in 1837
was called to the Chair of Chemistry in the newly-founded University of
London, in succession to Edward Turner. In 1854 he was made Master of
the Mint. He died in London on September 16, 1869.
Graham’s work was mainly devoted to that section of the science now
known as physical chemistry. His contributions to pure chemistry
are few in number. By far the most important is his discovery of
_metaphosphoric_ acid and its relations to the other modifications of
phosphoric acid. Ortho- or ordinary phosphoric acid was known to
Boyle; pyrophosphoric acid was discovered by Clark. Graham’s work is
noteworthy as first definitely indicating the inherent property of the
acids to combine with variable but definite amounts of basic substances
by successive replacement of hydroxyl groups—the property we now term
_basicity_, and was of fundamental importance in regard to its bearing
on the constitution of acids and salts.
Graham’s fame chiefly rests upon his discovery of _the law of gaseous
diffusion_ (1829–1831), upon his work on the _diffusion of liquids_,
and upon his recognition of the condensed form of hydrogen he termed
_hydrogenium_. Questions involving the conception of molecular
mobility, indeed, constituted the main feature of his inquiries. We owe
to him, among others, the terms _crystalloid_, _colloid_, _dialysis_,
_atmolysis_, _occlusion_—all of which have taken a permanent place in
the terminology of science.
[Illustration: ALEXANDER WILLIAM WILLIAMSON.]
│Alexander William Williamson│ was born at Wandsworth, London, on
May 1, 1824. His father, a Scotchman and a fellow-clerk of James
Mill (the father of John Stuart Mill) in the East India House, took
an active share in the foundation, in 1826, of the University of
London, subsequently known as University College. In 1840 the younger
Williamson entered the University of Heidelberg with the intention
of studying medicine; but, under the influence of Leopold Gmelin, he
turned to chemistry. In 1844 he went to Giessen, to work under Liebig,
and there made his first contributions to chemical science—viz., on the
_decomposition of oxides and salts by chlorine_; _on ozone_; and _on
the blue compounds of cyanogen and iron_. After graduating at Giessen
he went, in 1846, to Paris, where he came under the influence of Comte,
with whom he studied mathematics. In 1850, at Graham’s solicitation,
he was appointed to the Chair of Practical Chemistry at University
College, vacant by the death of Fownes. He at once embarked upon those
researches which constitute his main contribution to science. In the
attempt to build up the homologous series of the aliphatic alcohols
from ordinary alcohol he succeeded in demonstrating the real nature
of ether and its genetic relation to alcohol, and in explaining the
process of etherification. The memoirs (1850–52) in which he embodied
the facts had an immediate influence on the development of chemical
theory. His explanation of the process of etherification familiarised
chemists with the idea of the essentially dynamical nature of chemical
change. He imported the conception of molecular mobility not only
into the explanation of such metathetical reactions as the formation
of the ethers, but into the interpretation of the phenomena of
chemical change in general. In these papers, as also in one on the
constitution of salts, published in 1851, he attempted to systematise
the representation of the constitution and relations of oxidised
substances—organic and inorganic—by showing how they may be regarded as
built up upon the type of water considered as
H
O,
H
in which the hydrogen atoms are replaced, wholly or in part, by
other chemically equivalent atoms. This idea was immediately adopted
by Gerhardt, was further elaborated by Odling and Kekulé, and was
eventually developed into a theory of chemistry.
Williamson continued to direct the laboratory of University College
until 1887, when he retired to the country. He died at Hindhead on May
6, 1904.
│Robert Wilhelm Bunsen│ was born at Göttingen on March 31, 1811, and
after studying chemistry under Stromeyer, the discoverer of cadmium,
went to Paris and worked with Gay Lussac. In 1836 he succeeded Wöhler
as teacher of chemistry in the Polytechnic School of Cassel, and in
1842 became Professor of Chemistry in the University of Marburg. In
1852 he was called to Heidelberg, and occupied the Chair of Chemistry
there until his retirement in 1889. He died at Heidelberg on August 16,
1899.
[Illustration: BUNSEN, KIRCHHOFF, AND ROSCOE.]
Bunsen first distinguished himself by his classical work on the
_cacodyl compounds_, obtained as the result of an inquiry into the
nature of the so-called “fuming liquor of Cadet,” an evil-smelling,
highly poisonous, inflammable liquid formed by heating arsenious oxide
with an alkaline acetate. The investigation (1837–1845) is noteworthy,
not only for the skill it exhibits in dealing with a difficult and
highly dangerous manipulative problem, but also for the remarkable
nature of its results and on account of their influence on contemporary
chemical theory. The research, in the words of Berzelius, was the
foundation-stone of the theory of compound radicals. The name _cacodyl_
or _kakodyl_ was suggested by Berzelius in allusion to the nauseous
smell of the compounds of the new radical _arsinedimethyl, As (CH3)2_,
as it was subsequently termed by Kolbe.
Bunsen greatly improved the methods of gasometric analysis; these he
applied, in conjunction with Playfair, to an examination of the gaseous
products of the blast furnace in the manufacture of iron, and thereby
demonstrated the enormous waste of energy occasioned by allowing
the gases to escape unused into the air, as was then the universal
practice. This inquiry effected a revolution in the manufacture of iron
as important, indeed, as that due to the introduction of the hot blast.
Bunsen devised methods for determining the _solubility of gases_
in liquids, for ascertaining the _specific gravity of gases_, their
_rates of diffusion_, and of combination or _inflammation_. In 1841 he
invented the _carbon-zinc battery_, and applied it to the electrolytic
production of metals, notably of _magnesium_, the properties of which
he first accurately described. In 1844 he contrived the _grease-spot
photo-meter_, which was long in general use for ascertaining the
photometric value of illuminating gas. His methods of ascertaining
_the specific heats of solids and liquids_ were simple, ingenious, and
accurate. In 1855–1863 he carried out, in conjunction with Roscoe, a
long series of investigations on the _chemical action of light_. In
1859, in association with Kirchhoff, he devised the first methods of
_spectrum analysis_, and explained the origin and significance of the
Fraunhofer lines in the solar spectrum, thus laying the foundations of
solar and stellar chemistry. The application of the _spectroscope_ to
analytical chemistry almost immediately resulted in his discovery of
_cæsium_ and _rubidium_.
Bunsen worked on problems of _chemical geology_, and made a long series
of analyses of _volcanic products_. With Schischkoff, he examined, in
1857, the products of fired gunpowder. He effected many improvements in
analytical chemistry; devised the _iodiometric method_ of volumetric
analysis, and systematised the processes of _water analysis_. Lastly,
he invented the _gas-burner_—a piece of apparatus with which his name
is inseparably associated, and which has been of inestimable service to
operative chemistry and in the arts. Bunsen was no theorist, and purely
speculative questions had little or no interest for him. At the same
time he was a great teacher, and made the chemical school of Heidelberg
no less famous than the schools of Giessen and Göttingen.
* * * * *
The mass of material relating to the development of chemistry which
has been accumulated during the past sixty years is so vast that it
would be hopeless to attempt to survey it in detail within the limits
of such a work as this. Nor, indeed, is this required in a history of
this character. Those who desire information concerning the origin and
sequence of the facts which collectively make up the superstructure
of modern chemistry must be referred to the encyclopædias or larger
treatises—or, preferably, to the numerous monographs, dealing with
special sections, which the volume and complexity of the matter to be
dealt with seem to render increasingly necessary. All we can do here
is to attempt to show what has been the main outcome of this sixty
years of incessant effort to elucidate the mysteries of chemical
phenomena and to ascertain the nature of the conditions which control,
modify, or determine them. All this effort is ultimately directed
to the solution of the fundamental problem of the constitution of
matter. The most significant result of this endeavour has been the
elaboration and consolidation of the doctrine of chemical atoms, not
necessarily of atoms in the limited Daltonian sense, but of atoms
considered as associations of particles, or corpuscles—that is, of
entities which _may_ be divisible, but which, in the main, are not
divided in the vast number of the transformations in which they are
concerned. This modification of the original conception of Dalton has
been thought by some to destroy the basis upon which his theory really
rests. There is no necessity for such an assumption. So pronounced an
atomist as Graham, as far back as 1863, in a suggestive paper entitled
_Speculative Ideas on the Constitution of Matter_, enlarged the
conception of the Daltonian atom in precisely the sense which recent
experimental work appears to require. The present position, too, as it
affects chemists, was equally well stated by Kekulé, in 1867, in the
following terms:
The question whether atoms exist or not has but little
significance from a chemical point; its discussion belongs rather
to metaphysics. In chemistry we have only to decide whether the
assumption of atoms is an hypothesis adapted to the explanation
of chemical phenomena. More especially have we to consider the
question whether a further development of the atomic hypothesis
promises to advance our knowledge of the mechanism of chemical
phenomena.
I have no hesitation in saying that, from a philosophical point of
view, I do not believe in the actual existence of atoms, taking
the word in its literal signification of indivisible particles of
matter; I rather expect that we shall some day find for what we now
call atoms a mathematico-mechanical explanation which will render
an account of atomic weight, of atomicity, and of numerous other
properties of the so-called atoms. As a chemist, however, I regard
the assumption of atoms not only as advisable, but as absolutely
necessary, in chemistry. I will even go further, and declare my
belief that _chemical atoms exist_, provided the term be understood
to denote those particles of matter which undergo no further
division in chemical metamorphoses. Should the progress of science
lead to a theory of the constitution of chemical atoms—important as
such a knowledge might be for the general philosophy of matter—it
would make but little alteration in chemistry itself. The chemical
atoms will always remain the chemical unit; and for the specially
chemical considerations we may always start from the constitution
of atoms, and avail ourselves of the simplified expression thus
obtained—that is to say, of the atomic hypothesis. We may, in
fact, adopt the view of Dumas and of Faraday—that, _whether matter
be atomic or not, thus much is certain: that, granting it to be
atomic, it would appear as it now does_.[1]
[1] _The Study of Chemical Composition_, by Ida Freund
(Cambridge University Press), 1904.
The greater part of that which follows will be devoted, therefore,
to an exposition of certain of the great advances in knowledge—many
of them of primary importance—which have been made during the last
fifty or sixty years and which have served to strengthen this extended
conception of the atomic theory, and to establish its position as an
article of the scientific faith of the twentieth century.
CHAPTER II
THE CHEMICAL ELEMENTS DISCOVERED SINCE 1850
In 1850 the number of substances generally recognised as chemical
elements, in the sense in which that term was first employed by Boyle,
was sixty-two. Two members—viz., the _pelopium_ of Rose and the
_ilmenium_ of Hermann—were, however, subsequently shown to be identical
with metals already known. At the present time (1910) the number of the
chemical elements definitely recognised as such is eighty-two. In 1850,
as now, they were broadly classified as metals and non-metals, although
it was felt then, no less strongly than now, that no very clear line of
demarcation was traceable between the two groups. Sixty years ago the
elements usually styled non-metals were thirteen in number; to-day the
number is nineteen—the increase being due to the inclusion of arsenic
and the discovery of the so-called inactive elements, helium, argon,
krypton, neon, and xenon. In 1850 there were forty-seven elements
definitely classed as metals; in 1910 the number is sixty-three.
At all periods in the history of chemistry as a science the general
tendency has been to name substances, whenever possible, in accordance
with the theoretical conceptions of the time, and hence it has happened
that the same body at successive periods has had very dissimilar
names. But in naming the substances we term elements, theoretical
conceptions are not usually applicable. Oxygen, it is true, derives its
name from such a conception; and, etymologically, the name connotes
an error. Hydrogen, too, has no more right to be called the _water
former_ than oxygen. Davy, who invented the term chlorine, advocated
that the chemical elements should be named from some distinguishing
peculiarity, either of origin or of physical property. In the main this
principle has been adopted especially in later years although there
are numerous instances of names derived from pure arbitrary sources.
It is largely for the reason that the names of the elements are, with
rare exceptions, unconnected with theories that they have remained
unchanged, whereas names of compounds, which are far more frequently
dependent upon speculative ideas, have constantly been altered in
order to comply with the prevailing hypotheses of the period. At the
same time it is not always clear that the etymology of certain of
the elements is well ascertained. It has been recently shown, for
example that the commonly accepted origin of the word “antimony” from
_antimoine_, based on the alleged experiences of mediæval ecclesiastics
has no valid foundation. The word is, in reality, derived from the
Arabic _alhmoud_: this became latinised to _althimodium_ and eventually
to _antimonium_.
By the middle of the nineteenth century the system of symbolical
notation suggested by Berzelius was everywhere current; and, stripped
largely of its dualistic associations, this system still remains
the most generally convenient method of expressing the composition,
analogies, and numerical relations of substances. During the middle
of the last century philosophic chemists, although subscribing, with
hardly an exception, to the doctrine of definite combining proportions,
were by no means agreed as to the sufficiency of Dalton’s explanation
of the experimental laws of chemical combination; and the hypothesis of
atoms in the Daltonian sense was not universally accepted. To some the
atomic theory of Dalton, which assumed that the combining proportion
was identical with the relative weight of the atom, was unnecessary as
an explanation of the laws of combination. Or at most it was only one
out of a variety of molecular conditions in which matter might exist.
Consequently some chemists were in the habit of drawing a distinction
between _chemical_ atoms and _physical_ atoms. The chemical atom was
identical with the Daltonian atom but this was by no means the same
as the physical atom of Democritus or Leucippus. The view in 1850, in
fact, was not very dissimilar from that to which recent experimental
inquiry has led. But it can hardly be said that the doubts were
dependent upon valid experimental evidence; they arose rather from the
erroneous interpretation of imperfectly ascertained facts—upon the
supposed inconsistencies of the law of Gay Lussac with the hypotheses
of Avogadro and Ampère. As soon as the facts were clearly perceived
and the inconsistencies reconciled we heard less of the supposed
distinction between the chemical and the physical atom. It is only
within quite recent time, and as the result of entirely new lines of
inquiry, that the distinction has been revived.
In the early part of the last century attempts were made by Berzelius
to classify the chemical elements according to their electro-chemical
relations, and by Thomson according as they were “supporters” or
“non-supporters of combustion.” It was soon perceived that Thomson’s
system had no philosophical basis, and it quickly fell into disuse.
After the discovery of isomorphism, an endeavour was made by Graham to
arrange the simple bodies in accordance with their natural relations,
and even before 1850 the various elements were grouped by him very much
as now.
This scheme of classification, somewhat modified by considerations
of valency, and occasionally corrected by more accurate information
concerning true analogies (as when vanadium was transferred by Roscoe
to the nitrogen group), was in general use for practically a quarter of
a century—in fact, until it was superseded by the gradual adoption of
Mendeléeff’s arrangement based on periodicity. There can, however, be
little doubt that this attempt by Graham at a natural classification
paved the way along which Newlands and eventually Mendeléeff were led
to devise our present rational system of grouping the chemical elements.
The numerical relationships existing among the equivalents and atomic
weights of the elements of certain of these groups, pointed out by
Dumas, Pettenkofer, Odling, Gladstone, and others, gave rise to much
speculation. The values of the gradational differences, of course,
depended upon whether equivalents or atomic weights were employed; but
the immediate point is that, whichever basis was adopted, definite
numerical relations were to be perceived. Thus, in the case of the
group of the halogens, it was pointed out that the individual members
are connected together as follows:
Fluorine. Chlorine. Bromine. Iodine.
19 35.5 80 127
a a + d a + 2d + d´ 2a + 2d + 2d´
where a = 19; d = 16.5; d´ = 28.
Thus, too, in the case of the nitrogen group:
Nitrogen. Phosphorus. Arsenic. Antimony. Bismuth.
14 31 75 119 207
a a + d a + d + d´ a + d + 2d´ a + d + 4d´
where a = 14; d = 17; d´ = 44.
* * * * *
On the basis of these and similar numerical relationships it was
surmised that, just as the successive members of a group of homologous
organic radicals are formed by increments of CH2, so the substances
in the several groups of the elements may be produced by successive
additions of some form of matter common to them all. This has its
counterpart, somewhat modified, in the modern hypothesis of the
disintegration of the elements. Dumas conceived the elements in any
particular group to be built up by successive accretions of particular
forms of matter; Rutherford and Soddy suppose them to be derived by the
successive elimination of matter from some unstable parent substance.
Since 1850 the existence of at least twenty-two new elements may
be said to have been established. Of course, many more than this
number have been announced, more or less tentatively; but subsequent
investigation has either not confirmed their existence, or has
definitely disproved it. The names, symbols, and atomic weights of the
twenty-two, arranged in alphabetical order, are as follows:
Argon A 39.9
Cæsium Cs 132.8
Dysprosium Dy 162.5
Europium Eu 152.0
Gadolinium Gd 157.3
Gallium Ga 69.9
Germanium Ge 72.5
Helium He 4.0
Indium In 114.8
Krypton Kr 83.0
Lutecium Lu 174.0
Neodymium Nd 144.3
Neon Ne 20.0
Praseodymium Pr 140.6
Radium Ra 226.4
Rubidium Rb 85.4
Samarium Sa 150.4
Scandium Sc 44.1
Thallium Tl 204.0
Thulium Tm 168.5
Xenon Xe 130.7
Ytterbium } Yb 172.0
(Neoytterbium) }
The additions have been due, to some extent, to the refinement of
processes of analysis already in use, but more especially to the
employment of new analytical methods; or, lastly, to the application
of a generalisation concerning the mutual relations of the elements
which has served to indicate not only the existence of new and specific
members of families of elements already known, but to point out the
probable mode of their occurrence.[2]
[2] The substances which appear to be formed by the
disintegration of uranium, radium, thorium—the so-called
radio-active elements—such as _ionium_, _actinium_,
_polonium_, and the various _emanations_ to which they
give rise, are not here enumerated. They are dealt with in
Chapter III.
Although the existence of the element _fluorine_ was surmised as far
back as 1771, when Scheele first recognised that the product of the
action of oil of vitriol upon fluor-spar contained a hitherto unknown
substance, it was not until 1886 that this substance was definitely
isolated by Moissan by the electrolysis of the acid potassium fluoride
in solution in hydrogen fluoride. Cerium tetrafluoride, CeF4, and
lead tetrafluoride, PbF4, when heated, were observed by Brauner to
evolve a gas having a smell resembling that of hypochlorous acid,
which was probably free fluorine. Certain violet-coloured varieties
of fluor-spar, when powdered, emit a peculiar smell, which has been
attributed to free fluorine.
Gore observed that anhydrous hydrogen fluoride would not conduct
electricity—a fact confirmed by Moissan. Moissan found, however,
that on adding potassium fluoride to the liquid it readily suffered
electrolysis with the liberation of free fluorine as a light greenish
yellow gas with a pungent, irritating smell resembling that of
hypochlorous acid. It has a vapour density corresponding with an
atomic weight 19. By the application of cold and pressure it may be
liquefied. At still lower temperatures it may be frozen to a white
solid. Fluorine is characterised by an extraordinary chemical activity,
and combines, even at ordinary temperatures, with a large number of
substances. Sulphur, phosphorus, arsenic, antimony, boron, iodine, and
silicon inflame or become incandescent in contact with it. It combines
with hydrogen with explosive violence, even in the dark and at the
lowest temperature. It unites also with the metals, occasionally with
incandescence, and decomposes water with liberation of oxygen.
The application, by Bunsen, of the _spectroscope_ to chemical analysis
almost immediately resulted in his discovery, in 1860, of _cæsium_,
and, in 1861, of _rubidium_. Cæsium was first detected in the mineral
water of Dürkheim in the Palatinate and in the mineral petalite, by
the two blue lines it forms in the spectrum, whence its name from the
Latin _cæsius_, used to designate the blue of the clear sky. Rubidium
was found in a lepidolite by means of a number of lines in different
parts of the spectrum not previously observed, two being especially
remarkable in the outermost region of the visible red portion—whence
the name of the element from the Latin _rubidus_, used to designate
the darkest red colour. The new metals were found to have the closest
analogies to potassium, with which they usually occur associated
in nature. Rubidium is found in a number of lepidolites, leucite,
spodumene, triphylite, mica, and orthoclase, and in the Stassfurt
carnallite; in sea-water and in many mineral waters. It occurs also
in the ashes of many plants such as those of beetroot, tobacco, tea,
coffee, etc. It is doubtful if it is a normal constituent of plant
food, attempts to introduce it in place of potash having failed. It is
not improbable that these elements would have remained unknown except
for spectrum analysis. At all events, one of them—cæsium—was missed
in 1846 by Plattner, in the course of the analysis of the mineral
_pollucite_, in which it occurs to the extent of one third of its
weight. After the discovery of cæsium by Bunsen, this mineral was again
analysed by Pisani, when it was found that the alkali which Plattner
had mistaken for potassium was in reality cæsium. Cæsium is found to a
very small extent in many mineral waters, in a variety of minerals, and
in the ashes of plants.
In 1861 Sir William Crookes made known the existence of a new
element which he called _thallium_. He found it in a seleniferous
deposit obtained from an oil of vitriol factory in the Harz. It was
characterised by giving a bright green line in the spectroscope—whence
its name from θαλλός, a green or budding twig. The discovery was
confirmed in the following year by Lamy. Thallium, in its general
chemical relations, has many analogies to the metals of the alkalis
although in the metallic state it has the closest resemblance to lead.
It occurs in many varieties of pyrites, in a few minerals, such as
crookesite, lorandite, zinc-blende and copper pyrites, etc., and in
certain mineral waters.
In 1863 Reich and Richter, by means of the spectroscope, detected
the presence of a new element in the zinc-blende of Freiberg.
The observation that it afforded two indigo-blue lines in the
spark-spectrum led them to give it the name _indium_. It has since been
found in numerous blendes, in various zinc and tungsten ores, and in
many iron ores. It is a silver-white, ductile, and malleable metal,
melting at 174°, and burning when heated with a violet flame. It is
related in chemical characters to aluminium and zinc. Its true place in
the natural scheme of classification of the elements was indicated by
Mendeléeff.
In 1875 Lecoq de Boisbaudran discovered a new element in the
zinc-blende of Pierrefitte in the Pyrenees, also by means of spectrum
analysis. The spark-spectrum of its salts affords two characteristic
violet lines quite different in position from those given by indium.
To the new element its discoverer gave the name of _gallium_. It has
been found in very small amounts in other blendes, but is still one
of the rarest of the chemical elements. It is a bluish-white, hard,
and slightly malleable metal fusing at a temperature not much higher
than that of a hot summer day. Its existence and main properties, as
well as its more significant chemical relationships, were predicted by
Mendeléeff in 1869 from considerations based upon his periodic law.
(See _ante_.)
In the same year Mendeléeff also predicted the existence of a new
element belonging to the group of which boron is the first member,
which he provisionally termed _eka-boron_, and described its main
properties. Mendeléeff’s prediction was verified in 1879 by Nilson’s
discovery of the element _scandium_. Scandium occurs associated with
yttrium, ytterbium, etc., in many Swedish minerals, such as _euxenite_,
_gadolinite_, _yttrotitanite_, etc. The metal itself has not been
isolated, but the properties of its compounds correspond closely
with those of the corresponding ekaboron compounds, as predicted by
Mendeléeff.
A further illustration of the value of the principle of periodicity, as
developed by Mendeléeff, in indicating the existence of new elements,
is seen in the discovery of _germanium_. In 1885 Weisbach discovered a
new Freiberg silver mineral, to which he gave the name _argyrodite_.
This on analysis by Winkler was found to contain a new element to the
extent of about seven per cent. with properties identical with those
predicted by Mendeléeff for a missing element in the fourth group of
the periodic series, consisting of silicon, tin, and lead, and which
he had provisionally termed _eka-silicon_. _Argyrodite_, in fact, is
a double sulphide of silver and germanium, 2Ag2S.GeS2. Germanium is a
greyish-white, lustrous metal of sp.gr. 5.5., melting at about 900°,
and resembling silicon and tin in its general chemical relations.
_Dysprosium_, _europium_, _gadolinium_, _lutecium_, _neodymium_,
_praseodymium_, _samarium_, _thulium_, and _ytterbium_ (_neoytterbium_)
belong, like scandium, to the group of the so-called rare earth metals.
These substances have been detected in a great variety of minerals,
many of which are extremely rare. The elements most frequently occur in
nature associated with yttrium, cerium, thorium, and zirconium.
_Dysprosium_ was first detected, in 1886, by Lecoq de Boisbaudran in
the so-called erbium earth of Mosander, in which Cleve had previously
(1880) announced the existence of two other elements, _holmium_ and
_thulium_. There is some reason to believe that the holmium of Cleve is
identical with dysprosium. _Ytterbium_ was discovered by Marignac, in
1878, in the mineral _gadolinite_. In 1906 Auer von Welsbach announced
that Marignac’s ytterbia was a mixture, which was confirmed in the
following year by Urbain, who separated it into two elements, which
he named _neoytterbium_ and _lutecium_. _Europium_ was discovered by
Demarçay in 1901. All these earths are met with in small quantities
associated with yttria in _gadolinite_, _euxenite_, _samarskite_,
_xenotime_, _cerite_, _orthite_, and other similar minerals. Their
compounds, or such of them as have been described, resemble the
corresponding compounds of yttria. They are recognised by differences
in their spectroscopic behaviour. _Gadolinium_ was detected,
independently, in 1886, by Marignac and Lecoq de Boisbaudran in the
terbium earth of Mosander.
What was long known as _didymium_ (διδυμος = a twin) was discovered by
Mosander in 1841. It owes its name to its close chemical relationship
to, and almost constant association with, _lanthanum_—both elements
occurring in many minerals, more particularly in _cerite_, _allanite_,
and _monazite_. In 1885 Auer von Welsbach announced that the didymium
of Mosander was, in reality, a mixture of two elements which could be
separated by the systematic fractional crystallisation of the double
ammonium nitrates; to these elements he gave the names _praseodymium_
(πράσινος, leek-green) and _neodymium_ (νέος, new). Neodymium salts
are rose-coloured, whereas those of praseodymium are green, and the
elements are further characterised by differences in their absorption
and spark-spectra. When mixed, the substances give the spectrum
originally considered to be characteristic of didymium.
_Samarium_ was discovered in 1879 by Lecoq de Boisbaudran in
_samarskite_. Its salts are yellow, and afford in solution
characteristic absorption bands.
It is not improbable that many of the minerals from which the so-called
rare earths are obtained contain elements hitherto unrecognised, and it
is possible that certain of the substances now assumed to be elements
may, like didymium, turn out to be mixtures. In fact, additional
elements have from time to time been announced, as for example, the
_decipium_ of Delafontaine (1878) and the _monium_ or _victorium_ of
Crookes (1899), pronounced by Urbain to be identical with gadolinium:
their individuality cannot as yet be said to be established. Didymium
itself was stated by Krüss and Nilson (1888) to be even more
complicated than the work of Auer von Welsbach would seem to indicate,
and to contain no fewer than eight elementary substances. As yet,
however, no confirmation of this surmise has been obtained.
The chemistry of the rare earths has of late years been greatly
extended owing to the employment of certain of the members of the
group in the manufacture of the “mantles” used in gas-lighting, and
which consist substantially of thoria, mixed with about one per cent.
of ceria. Large quantities of _monazite_, _thorianite_, _thorite_,
_cerite_, and other minerals, are now worked up for the sake of the
thoria and ceria they contain, and considerable amounts of residual
products, consisting largely of other members of the family, are now
available for investigation. It is reasonably certain, therefore, that
our knowledge of this section of inorganic chemistry will be largely
augmented in the immediate future. Indeed, the application of thoria
to the construction of gas-mantles may be said to have removed that
substance from the category of the rare elements. No sooner was it
discovered that it was capable of useful application than unexpected
sources of supply were found.
The same result has followed in other cases. One of the most
significant developments of modern chemistry is seen in the efforts
which are constantly being made to turn the so-called rare elements to
useful account; and when they are found to be technically valuable it
is generally observed that hitherto unknown sources of supply are soon
available. Cerium salts have been found to be useful in the colouring
of glass and porcelain, as mordants in dyeing, in photography, and in
medicine. Zirconium has been used in incandescent electric lighting,
and thallium has been employed in the manufacture of highly refractive
optical glass. Titanium, molybdenum, and vanadium are used in the
manufacture of steel of high tensile strength. Tantalum and tungsten
are employed in the construction of filaments in incandescent electric
lighting. Tantalum, indeed, has been found to occur in considerable
quantities, and to be more largely distributed than was hitherto
supposed. Alloys of tungsten and aluminium are used in automobile
construction, and alloys of tungsten, aluminium, and copper in the
manufacture of propeller blades. Tungsten steel is used in armour
plates, and to stiffen the springs of cars; in the manufacture of
piano-wire, and to increase the permanency of magnets. Even the rarer
metals of the platinum group are finding many important applications.
Osmium-iridium is used for the bearings of compasses, for the tips
of gold pens, and in the construction of standard weights. Osmium
and ruthenium enter into the composition of filaments for electric
lighting. The extraordinary influence of light on the electric
conductivity of selenium has been made use of in the transmission of
photographs by telegraph and telephone wires, and for measuring the
light intensity of the Röntgen rays in clinical work.
CHAPTER III
THE INACTIVE ELEMENTS: RADIUM AND RADIO-ACTIVITY
_Argon_, _helium_, _krypton_, _neon_, and _xenon_ belong to the group
of the so-called inactive elements, and constitute what are known as
the rare gases of the atmosphere. The existence of these bodies is
of great theoretical value and few discoveries of recent times have
exacted more interest and curiosity. Twenty years ago it was generally
assumed that practically all that was to be known concerning the
composition of atmospheric air had been ascertained. Priestley and
Cavendish had recognised that it was mainly composed of oxygen and
nitrogen, and Cavendish had definitely stated that these gases are
present in practically constant proportion, independent of season,
climate, or locality. Thénard, Saussure, and others, had determined
the limits of variation in the amount of carbon dioxide. Bunsen and
Regnault had established that the quantities of oxygen and nitrogen are
subject to slight alteration, the extent of which could be readily
determined by the exact eudiometric processes they had devised.
Lastly, it was proved beyond a doubt that the gases of the atmosphere
are simply mechanically mixed, and can be separated by a variety of
physical methods. In fact, of no single subject could it be more
confidently assumed that finality of knowledge had apparently been
reached.
In 1892, in the course of a series of determinations of the densities
of the common gases, Lord Rayleigh found that the density of nitrogen
obtained from the air was slightly greater than the density of that
gas prepared by the decomposition of ammonia and of nitric acid, the
difference in weight being about 1 part in 200—an amount far greater
than could be accounted for by errors of weighing. Various suppositions
were made in explanation of the discrepancy; but these, when tested,
were found not to account for the facts. By heating the atmospheric
nitrogen with metallic magnesium, whereby the greater portion of the
gas is absorbed to form the nitride, Sir William Ramsay found that the
density of the residual gas was still further increased, which rendered
it probable that the relatively high density of atmospheric nitrogen
as compared with that derived from ammonia, and, as Lord Rayleigh
found, from other sources also, was due to the presence of a gaseous
substance in the air of considerably greater density than nitrogen or
oxygen. Lord Rayleigh also subjected atmospheric nitrogen mixed with
oxygen to the electric discharge over a solution of caustic soda, in a
manner similar to that already employed by Cavendish, and found also
that the residual gas was considerably increased in density. At the
Oxford meeting of the British Association in August, 1894, the two
investigators were in a position to announce that the discrepancy was
actually due to the presence of a hitherto unknown gaseous constituent
of atmospheric air, considerably more soluble in water than nitrogen,
and to which, on account of its chemical inertness, the name of
_argon_ (ἀργον, idle) was given. By a special apparatus devised by
Lord Rayleigh, in which a mixture of air and oxygen is submitted to an
electric flame produced by a powerful, rapidly alternating current,
considerable quantities of argon were separated from the air. It has
also been found that by the use of metallic calcium or a mixture of
magnesium and lime, the atmospheric nitrogen is absorbed at a lower
temperature, and more rapidly than by magnesium alone.
Argon has been found to exist in the gases from springs and mineral
waters, notably in those of Bath, Cauterets, Wildbad, and Harrogate. It
has also been found in a meteorite, in the gas occluded in rock-salt,
and in the minerals _malacone_, _uraninite_, _brōggerite_, etc. No
animal or vegetable substance appears to contain it. It is present in
atmospheric air to the extent of about one per cent. by volume. It is
a colourless gas of an atomic weight of 39.9: one litre of it at the
standard temperature and pressure weighs 1.7815 grams. Experiments made
by the method of Kundt and Warburg—_i.e._, by determining the ratio
of the specific heats at constant pressure and constant volume by the
velocity of sound in the gas—prove that argon, like mercury gas, is
monatomic. This of itself indicates that argon is an element, since a
monatomic compound is a contradiction in terms. The calculations from
the experimental data presuppose that argon obeys the laws of Boyle and
Dalton, which was found on trial to be the case. By the application
of cold and pressure argon can be liquefied. The liquid boils at
-186°.1 and freezes at -187°.9. The spectrum of the gas is exceedingly
complicated, consisting of a great number of lines extending throughout
the visible portion and far into the extreme red and ultra-violet. The
colour of the light emitted on sparking the gas changes with increase
of temperature from a brilliant red to a bright blue—depending on the
intensity of the discharge. All attempts to induce argon to enter
into combination with other substances have failed. The methods of
its preparation show that it does not combine with oxygen, although
Troost and Ouvrard state that it unites with magnesium vapour. It forms
no compounds with hydrogen, chlorine, phosphorus, sulphur, sodium,
tellurium, etc. Even fluorine, probably the most generally active of
the chemical elements, shows no tendency to unite with it.
In 1888 Dr. Hillebrand, of the U.S. Geological Survey, in examining
a form of uraninite known as _cleveite_, so named from the late
Professor Cleve, found that on treatment with dilute sulphuric acid
it gave off considerable quantities of a gas which was assumed to
consist only of nitrogen, as it gave the spectroscopic reactions of
that element. To test whether this gas contained argon, Ramsay, in
1895, further examined it spectroscopically. After sparking it with
oxygen in the presence of caustic soda solution, in the way already
described, it gave no indications of argon. The main characteristic
of its spectrum was a bright yellow line, known as D3, not coincident
with that afforded by sodium, but identical in position with a line
detected in the chromosphere during the solar eclipse of 1868, which
line, on examination by Frankland and Lockyer, could not be ascribed
to any known element. For this supposed new element the name _helium_,
from ἥλιος, the sun, had been suggested. This was the first occasion
on which an element observed originally only in the sun was found to
occur also on the earth. The presence of the new element in the gas
from cleveite was subsequently confirmed by Langlet working in Cleve’s
laboratory.
[Illustration: SIR WILLIAM RAMSAY.]
Helium is a monatomic gas having the atomic weight 4. It is less
soluble in water than argon. Like argon, it shows no tendency to
enter into chemical union with any other substance. It has been found
in many minerals, particularly in those containing uranium and the
so-called rare earth metals. It also occurs among the gases issuing
from certain mineral springs, such as those of Bath and at Cauterets
in the Pyrenees, and also at Adano near Padua. The spectrum of helium
contains, in addition to the characteristic yellow line—by which its
presence had been recognised not only in the solar chromosphere, but
also in certain of the fixed stars—two lines in the red, and lines in
the green, blue, and violet. The character of the light emitted by
the spark-discharge is modified by the intensity of the discharge in
a manner similar to that of argon. It has been shown by Collie that
its spectrum is altered by the presence of mercury vapour. It is the
least refractive of all the gases. Helium was liquefied by Kammerlingh
Onnes in 1908. It forms a colourless liquid of sp. gr. 0.154, boiling
at -268.5; that is, 4°.5 above the absolute zero of temperature. Its
critical temperature is about 5° absolute, and its critical pressure
above 2¼ atmospheres.
The methods now in use for obtaining liquid air, referred to in a
subsequent chapter, enable large quantities of that material to be
obtained readily; and it was in investigating spectroscopically the
residues left after volatilising a quantity of liquid air that Ramsay
and Travers, in 1898, detected the existence of two new monatomic
gaseous constituents of the air which they named respectively _krypton_
(χρυπτός, hidden) and _neon_ (νέος, new), the former heavier and the
latter lighter than argon. By fractional distillation of the argon,
simultaneously procured, a gas was obtained which in the spectroscope
showed the characteristic lines of helium—previously recognised in
atmospheric argon by Kayser and Friedländer—together with a complicated
spectrum consisting of a number of lines in the red, orange, and yellow
due to the new element neon. On cooling this mixture to -252° by means
of liquid hydrogen, the neon solidified, while the helium remained
gaseous and could thus be separated.
Krypton was obtained from the residues left on the evaporation of a
large quantity of liquid air. Mixed with the krypton was a third
gaseous constituent of air, to which the name _xenon_ (ξενος, the
stranger) was given. The boiling-point of krypton at atmospheric
pressure was found to be -152°, and its melting-point -169°; the
boiling-point of xenon was -109° and its melting-point -140°. Their
critical temperatures were respectively -62°.5 and +14°.7. Hence
xenon could be liquefied by pressure a very little below the mean
temperature of the air. Neon boils at -243° and freezes at -253°. They
form colourless liquids freezing to ice-like solids. All of them, with
the exception of argon, which is present to the extent of about 1 part
in 107 parts of air, are contained in extremely small amounts in the
atmosphere, approximately in the following proportions:
Helium 1 part in 245,300 parts by volume.
Neon 1 ” ” 80,800 ” ”
Krypton 1 ” ” 20 millions ” ”
Xenon 1 ” ” 170 ” ” ”
Many tons of liquefied air have since been systematically fractionated,
but no other gas than those above named has been obtained.
Julius Thomsen, of Copenhagen, in a paper published in 1895, entitled
_On the Probability of the Existence of a Group of Inactive Elements_,
pointed out, in relation to Mendeléeff’s Law of Periodicity (see
_ante_), that in periodic functions the change from negative to
positive value, or the reverse, can take place only by a passage
through zero or through infinity; in the first case the change
is gradual, and in the second case it is sudden. The first case
corresponds with the gradual change in electrical character with rising
atomic weight in the separate series of the periodic system, and the
second case corresponds with a passage from one series to the next.
It therefore appears that the passage from one series to the next in
the periodic system should take place through an element which is
electrically indifferent. The valency of such an element would be zero,
and therefore in this respect also it would represent a transitional
stage in the passage from the univalent electronegative elements of the
seventh to the univalent electropositive elements of the first group.
This indicates the possible existence of a group of inactive elements
with the atomic weights 4, 20, 36, 84, 132—numbers corresponding fairly
closely with the atomic weights respectively of helium, neon, argon,
krypton, and xenon.
* * * * *
No discovery of recent years has created more widespread interest than
that of the radio-active elements.
In 1896 Henri Becquerel found that uranium salts emitted an invisible
radiation which had the power of affecting a photographic plate,
even though not directly exposed to it, exactly in the same way as
the Röntgen or X-rays. Since that time a number of substances have
been shown to possess a similar property. Such substances are said
to be radio-active. The radiation emitted by them is not uniform in
character. It has been found to be of three distinct types, known
respectively as the α, β, and γ radiations. The α rays consist of
positively electrified particles moving with a velocity equal to about
a fifteenth of that of light. These rays have little penetrative power,
and are capable of being deflected by a magnet.
The β rays consist of negatively electrified particles of a mass not
greater than one thousandth of that of the hydrogen atom, and they
move with a velocity approximating to that of light. The β rays have a
greater penetrative power than the α rays, and are even more readily
deflected by a magnet.
The γ rays are analogous to, if not identical with, the X or Röntgen
rays; they move with the velocity of light, have a high penetrative
power, but are not affected by the magnet. All three forms of
radiation render gases electrically conductive, excite luminescence or
fluorescence in certain substances, change the colour of glass, convert
oxygen into ozone and yellow phosphorus into red phosphorus, and act
upon photographic plates.
According to the disintegration theory of Rutherford and Soddy,
the radio-active elements are forms of matter undergoing changes
resulting in the formation of new forms possessing chemical and
physical properties differing from those of the parent substance,
these changes being accompanied by the production of sensible
heat, or some other manifestation of energy, due to the process of
transformation of the changing atoms. The rate of change is found to
be different for each radio-active element, but to be constant for
the same element irrespective of its particular form of combination.
The relative radio-activity of the various chemical combinations of a
given radio-active element is directly proportional to the quantity
of the element contained in them. The process of disintegration
may be carried through a number of intermediate products until
a stable form is produced. Uranium, in which the phenomenon of
radio-activity was first perceived, is supposed to give rise to no
fewer than seventeen different forms of matter, including _radium_,
_actinium_, and _polonium_. Thorium, another radio-active element,
is supposed to disintegrate into eight different forms of matter.
Uranium disintegrates with extreme slowness; it is calculated that
in a year not more than one ten-billionth part of the uranium is
transformed. The first disintegration product is termed uranium │x│.
If a quantity of dehydrated uranium nitrate be treated with ordinary
ether, a slight residue is obtained which is found to contain uranium
│x│. It emits β and γ rays, and is relatively rapidly transformed
into other substances. Ordinary uranium, freed from uranium │x│, only
emits α rays. Uranium salts can be freed from uranium │x│ by repeated
crystallisation, uranium │x│ remaining in the mother liquors.
[Illustration: MARIE CURIE (_née_ SKLODOWSKA).]
The existence of radium was first made known by │Mme. Curie│ in 1898.
In examining certain uranium minerals and uranium products, Mme. Curie
observed that their radio-activity was apparently greater than that
corresponding with the amount of uranium contained in them, and she
was led to surmise that this might be due to the presence of some
constituent more strongly radio-active than uranium. This supposition
proved to be well founded, and she eventually succeeded in isolating
a new element termed _radium_, forming compounds with characters and
relationships akin to those of barium. The richest source of radium at
present known consists of certain residues occurring at Joachimsthal,
in Bohemia, left after the extraction of uranium from pitch-blende, in
which radium occurs to the extent of 0.2 gram per ton. These residues
are mainly sulphates of lead and calcium, mixed with a great variety
of other metallic compounds. To obtain the radium the mixture is
heated with concentrated caustic soda solution, the residue washed with
water and treated with hydrochloric acid which dissolves the greater
portion of the material. Nearly the whole of the radium is left in
the insoluble portion. This, after washing with water, is boiled with
a solution of sodium carbonate so as to transform the alkali-earths
into carbonates. These are converted into chlorides or bromides from
which, by repeated crystallisation, barium chloride or bromide is
obtained, containing the greater portion of the radium as a halide
salt. The radium and barium salts are then separated by fractional
crystallisation, the radium salts being slightly less soluble in water
and alcohol, and in solutions containing the halogen acid, than the
barium salt.
Pure radium chloride (RaCl2) is a white crystalline salt, resembling
barium chloride, with which it appears to be isomorphous. Radium,
like barium, forms an insoluble carbonate and sulphate, but a soluble
nitrate and bromide. The bromide is much less stable than the chloride;
on standing it evolves bromine and becomes basic. Radium has as yet
been obtained in such small quantities that very few of its compounds
have been prepared.
The rays from radium salts burn the skin, and are found to be useful
in the destruction of rodent ulcers; they appear to act upon proteids,
destroy bacteria, bleach chlorophyll, and affect the germinative power
of seeds. A pure and freshly-prepared salt of radium seems to emit only
α rays, but it soon forms disintegration products, and then gives out,
in addition, the β and γ rays.
In the process of disintegration the salts emit heat corresponding to
about 75 gram calories per hour for each gram of radium present; their
temperature is thus uniformly higher than that of their environment.
One product of the change probably connected with the emission of the α
rays, is the gas helium.
Radium has an atomic weight of 226.5. It is regarded as a product
of the disintegration of uranium, the atomic weight of which is
238.5. It is believed to have been formed through an intermediate
product known as _ionium_, a radio-active element discovered by
Boltwood in the mineral _carnolite_. The atomic weight of ionium is
surmised to be about 230. Radium itself is supposed to form at least
eight disintegration products, the first of which is the so-called
_emanation_, discovered by Dorn in 1900, an inactive gas with an
atomic weight of about 180, giving a bright line spectrum, decomposing
into helium, liberating oxygen and hydrogen from water, and capable
of being condensed to a liquid and solidified at a low temperature.
Ramsay and Gray have determined its physical constants. The liquid
is phosphorescent and shines with a colour depending on the nature
of the glass of the vessel which contains it. The solid is also
phosphorescent, the colour varying with the temperature. It gives out
only α rays and in its disintegration, like radium, evolves heat. Its
position in the Periodic Table is probably above that of xenon. Other
products are known as _radio-lead_ and _polonium_. The latter substance
was identified by M. and Mme. Curie in 1898, and was the first of the
strongly radio-active substances to be recognised. In the periodic
system it seems to follow bismuth and to be a member of Group VI., with
a possible atomic weight of 210. Its spectroscopic characters have
recently been examined by Mme. Curie and Debierne, who have shown that
in its decay it evolves helium.
The rate of disintegration of radium is relatively slow; it has been
calculated that the time required for half of any given quantity of
radium to change completely into other products is about 2000 years.
Rutherford has calculated that in 26,000 years a kilogram of radium
would be reduced to one milligram of active substance, the remainder
having passed into degradation products.
In 1899 Debierne announced the existence of another radio-active
element contained in uranium minerals, which he termed _actinium_. This
is probably a disintegration product of uranium and identical with the
_emanium_ of Giesel. It occurs associated with the rare earths which
can be separated from the pitch-blende residues, and is eventually
found in the lanthanum salts. Nothing is known as to its atomic
weight or its chemical relationships. It undergoes change, and forms,
apparently, a gaseous emanation which rapidly disintegrates and can
be condensed to a liquid at a low temperature. Four other successive
products have been identified by the character of the radiation they
emit, their degradation constants, and the time required for one half
of any given quantity to disintegrate into other forms of matter.
_Thorium_ was shown to contain a radio-active element by Mme. Curie
and Schmidt, independently, in 1898. Whether thorium is itself
active is doubtful. The rate of disintegration of _radio-thorium_
is probably greater than that of uranium. It, too, seems to form a
gaseous emanation which can be condensed at the temperature of liquid
air and appears to be an inert gas of high molecular weight with the
characteristics of the argon family.
The type of radiation emitted by the several products has been
observed, and their constants of change and half-value periods
calculated; but little or nothing is known at present concerning their
atomic weights, spectroscopic or chemical characters.
CHAPTER IV
ATOMS AND MOLECULES: ATOMIC WEIGHTS AND EQUIVALENTS
It has already been pointed out that the discovery by Gay Lussac, and
independently by Dalton, that gases combine in simple proportions by
volume, and that the volume of the gaseous product, measured under
comparable conditions of temperature and pressure, stand in simple
relation to the volumes of the constituents, seemed to most of Dalton’s
contemporaries, but not to Dalton himself, to afford strong evidence
of the validity of his explanation of the essential nature of chemical
combination. It appeared obvious from the facts that there must exist
some simple relation between the densities, or specific gravities,
of the elementary gases and their atomic weights. When, however, the
principle underlying Gay Lussac’s law was extended so as to include
gases in general—both simple and compound—difficulties were met with
which were only satisfactorily cleared away during the latter half of
the nineteenth century. The first rational attempt to explain the
facts observed by Dalton and Gay Lussac, concerning the volumetric
relations of gases, was made in 1813 by Amedeo Avogadro by the
assumption that a given volume of all gases—simple or compound—contains
the same number of integral molecules; hence the relative weights
of these volumes represent the relative weights of the molecules.
According to Avogadro, in the case of the simple gases the integral
molecules are composed of a certain number of elementary molecules
of _the same kind_, whereas the integral molecules of compound gases
and vapours are made up of elementary molecules _of different kinds_.
The _elementary molecule_ of Avogadro is now termed the _atom_; his
_integral molecule_ we call simply a _molecule_. Similar conceptions
were published independently by Ampère in 1814. It follows from the
doctrine of Avogadro and Ampère that, as the number of integral
molecules is the same in equal volumes of all gases, these molecules
must be equidistant from each other, their mutual distances depending
upon pressure and temperature. This at once serves to explain the laws
of Boyle and Dalton that gases, no matter what their chemical nature,
behave identically, as regards change of volume, when compressed by
pressure or expanded by heat.
The true significance of the hypotheses of Avogadro and Ampère was
long obscured, first on account of their imperfect appreciation by
the great leaders of chemical thought during the first half of the
nineteenth century—Berzelius, Gay Lussac, Wollaston, and Gmelin—and,
secondly, on account of the almost universal practice of deducing
atomic weights from purely chemical considerations of equivalence.
At the same time, it must be admitted that the recognition of the
value of these hypotheses was still further retarded by the seeming
anomalies which resulted from a more extended knowledge of the vapour
densities of elements and compounds. Thus the vapour densities of
mercury, sulphur, phosphorus, and arsenic, as ascertained by Dumas and
Mitscherlich, were plainly inconsistent with chemical analogies and
the law of Dulong and Petit. So, too, what appeared to be the vapour
densities of sal-ammoniac, phosphorus pentachloride, sulphuric acid,
calomel, and of other substances that might be mentioned, were not in
accordance with the values demanded by other well-ascertained facts.
By the middle of the nineteenth century the hypothesis of Avogadro
was practically forgotten and the law of volumes ignored. The atomic
weights of the elements, and the system of notation universally
employed in England and Germany, were based wholly upon equivalents.
The anomalies thus created were clearly pointed out by Gerhardt, and
subsequently by Laurent who showed how a consistent and harmonious
explanation of the facts could be reached by regarding as true
equivalents equal volumes—for example, of steam, ammonia, hydrogen
chloride, carbon dioxide, marsh gas, etc.; by assuming, in other words,
that equal numbers of the molecules of these various substances are
contained in equal volumes of the gases, as contended by Avogadro and
Ampère. The simplicity and consistency of the new notation gradually
won for it the adhesion of chemists. This adhesion was facilitated by
the memorable researches of Williamson on etherification, of Gerhardt
on the anhydrides, and by the work of Frankland on the radicals; and
it reached its logical conclusion when │Cannizzaro│—on the basis of
Avogadro’s hypothesis—discussed, in 1858, the true atomic weights of
the metallic elements as distinguished from their equivalent values.
In certain of its aspects the new table of atomic weights drawn up by
Cannizzaro resembled that originally proposed by Berzelius; but the
numbers adopted by the Swedish chemist were founded on no uniform or
rational basis, and were frequently inconsistent.
[Illustration: STANISLAO CANNIZZARO.]
Our present tables of atomic weight bring, therefore, the values
for the several elements into harmony with the doctrine of Avogadro
and Ampère, with the law of Dulong and Petit, and with the facts of
isomorphism. The values are, in fact, in unison with all the criteria
which serve to indicate the atomic weights of the elements. As a
result, our present system of notation, which is, of course, based upon
these atomic weights, assigns formulæ to compounds which indicate their
true relative molecular weights and simplify the accurate expression
of their relationships and chemical transformations. One by one the
instances of anomalous vapour density, which were so many stumbling
blocks to the universal acceptance of a system based upon the law of
gaseous volumes, have been shown to be not only not inconsistent with
it, but actually so many corroborative proofs. Thus, in the case of
ammonium chloride, the observed vapour density of which was found to
be practically half its calculated value, it has been proved that
the vapour of this salt, when heated to the temperature at which the
observations were made, is mainly resolved into molecules of ammonia
and hydrogen chloride, which together occupy double the space of the
ammonium chloride molecule. Phosphorus pentachloride vapour, on being
sufficiently heated, is similarly more or less resolved into phosphorus
trichloride and chlorine.
Moreover, it has been found, by a more accurate study of the action
of heat upon the vapour of ammonium chloride and of phosphorus
pentachloride, that within certain narrow limits of temperature these
substances can actually exist as such in the gaseous state, and that
the density of their vapours does actually conform to that demanded
by theory. Moreover, phosphorus pentafluoride, the analogue of the
pentachloride, is gaseous at ordinary temperatures, and has a normal
density. It may be heated to a high temperature without showing any
sign of decomposition.
Limitations of space will not allow of a fuller explanation of the
apparent anomalies already alluded to. They have each in turn been
experimentally attacked and satisfactorily explained. No valid
exception is now known to the universal applicability of the principle.
To-day the chemical history of a substance, whether elementary or
compound, if vaporisable, is not complete until its vapour density is
known, since a knowledge of this constant affords the most certain
means of establishing the relative weight of its molecule. Accordingly
many chemists have endeavoured to simplify and render more convenient
the modes of determining vapour densities. Thanks to the efforts of
Hofmann and Victor Meyer, the processes associated with the names of
Dumas, Gay Lussac, Deville, and Troost, which have furnished us with
valuable information in the past, have now given way to comparatively
simple and rapid methods, which, although not necessarily more
accurate, furnish the required information with less expenditure of
time and trouble; that is to say, they serve to indicate which of
two, or more, presumed molecular weights is correct, and so enable
us to establish the molecular formula of the substance. The chemical
formula of a substance is a condensed expression of a number of facts
connected with its history. Thus the expression H2O—the chemical
formula for water—indicates that the substance is composed of hydrogen
and oxygen, in the proportion, using round numbers, of 2 parts by
weight of hydrogen and 16 parts by weight of oxygen; or, in other
words, of 2 atoms of hydrogen, each weighing 1, and 1 atom of oxygen
weighing 16. The formula, moreover, connotes the fact that when the
gases combine 2 volumes of hydrogen unite with 1 volume of oxygen to
form 2 volumes of water-vapour (steam). So, too, the formula HCl—which
represents hydrogen chloride—means that the substance is a compound of
1 atom of hydrogen weighing 1 united with 1 atom of chlorine weighing
35.5; it also denotes the fact that in the act of union 1 volume of
hydrogen combines with 1 volume of chlorine to form 2 volumes of
hydrogen chloride. Lastly, the formula NH3 signifies that the molecule
of ammonia is composed of 1 atom of nitrogen weighing 14, and 3 atoms
of hydrogen, each weighing 1; it further indicates that when ammonia
gas is resolved into its constituents, as it can be when sufficiently
heated, 2 volumes of ammonia gas are increased to 4 volumes of a
mixture made up of 1 volume of nitrogen and 3 volumes of hydrogen.
In short, all the formulæ are what are called _two-volume formulæ_;
that is, the relative molecular weights of the substances occupy
the same volume as two relative parts by weight of hydrogen. Hence
their vapour densities—referred to hydrogen as unity—are the halves
respectively of their molecular weights. Steam is 9 times, hydrogen
chloride 18.25 times, and ammonia 8.5 times heavier than hydrogen, when
measured under identical conditions of temperature and pressure. The
quantities expressed by the formulæ H2, H2O, HCl, NH3, occupy the same
volume. This is equally true of the quantities expressed by O2, N2,
Cl2, etc. These expressions signify that the molecules of the elements
of hydrogen, oxygen, nitrogen, and chlorine consist each of 2 atoms of
the respective substances: the molecule of water consists of 3 atoms—2
of hydrogen and 1 of oxygen; the molecule of hydrogen chloride of 2
atoms—1 of hydrogen and 1 of chlorine; whereas the molecule of ammonia
contains 4 atoms—1 atom of nitrogen and 3 atoms of hydrogen.
Certain of the elementary bodies are, as already stated, capable of
existing in different allotropic states. Thus there is a modification
of oxygen known as _ozone_. This substance has long been recognised as
being formed from air under the influence of the electric discharge. It
was the subject of study by Schönbein as far back as 1839; but that it
was a condensed form of oxygen and not a peroxide of hydrogen—as was at
one time surmised—was first established by Andrews and Tait. The degree
of its condensation was definitely ascertained by Soret by observing
its rate of diffusion, from which, on the basis of Graham’s law, its
density could be inferred. It was found to be 1½ times that of ordinary
oxygen. Hence, if the molecule of oxygen consists of 2 atoms, that of
ozone consists of 3 atoms. The chemical symbol of ozone is, therefore,
O3.
It has been found also that the molecule of sulphur, in the state of
solution, contains eight atoms. This complex molecule in the gaseous
state gradually breaks down as the temperature is increased, and at
temperatures above 850° contains, like its analogue oxygen, only
two atoms. The molecules of phosphorus and arsenic, in the gaseous
state, are each found to consist of four atoms. On the other hand, the
molecules of mercury, zinc, and cadmium each consist of only one atom.
As will be shown later, Kundt and Warburg established, in 1875, the
fact that mercury vapour is a monatomic gas, by determining the rate at
which sound is propagated through it. In the same way it was shown that
helium, argon, and its congeners are also, as already stated, monatomic
gases.
The applicability of the law of Dulong and Petit to the determination
of atomic weights has been frequently exemplified during the last
sixty years; and a number of these constants have been rectified
by its means—_e.g._, of thallium, uranium, glucinum, indium, etc.
The anomalies presented by the cases of elements of low atomic
weight—_e.g._, carbon, boron, silicon—have been further inquired into;
and it has been shown by Weber, and independently by Dewar, that in the
case of these substances the specific heat rapidly increases with the
temperature, and approximates at high temperatures to a value required
by the law of Dulong and Petit.
Within recent years other methods of ascertaining molecular weights
have been put at the disposal of chemists. These methods are especially
valuable in the case of bodies which cannot be volatilised. They depend
upon the influence of the substance (1) upon the freezing-point and
(2) upon the boiling-point of a solvent. It has long been known that a
substance in solution affects the freezing-point of the solvent, and
in the great majority of cases depresses it. Sir Charles Blagden,
as far back as 1788, showed that in aqueous solutions of inorganic
salts the depression was proportional to the amount dissolved. It was
subsequently found by Coppet that, in a number of solutions of similar
salts where these were present in the ratio of their molecular weights,
the solutions froze at practically the same temperature: the molecular
depressions of the freezing-points differ from group to group but are
nearly equal in groups of similar compounds. Raoult further observed
that, when certain quantities of the same substance are successively
dissolved in a solvent on which it exerts no chemical action, there
is a progressive lowering of the point of solidification of the
solution, and this depression is proportional to the weight of the
substance dissolved in a constant weight of the solvent. In the case
of a large number of solvents the depressions of the freezing-point,
calculated for amounts proportional to the molecular weights of the
dissolved substance, were nearly constant. Raoult pointed out that
these relations between the molecular weights and the lowering in
the freezing-point may be employed to determine the molecular weight
of a soluble substance. The molecular weight _m_ is found from the
expression _m_ = K/A, where A is the quotient obtained by dividing
the observed depression in the freezing-point of the solvent by the
percentage content of the solution, and K (the molecular depression)
is a constant dependent on the solvent. Thus in the case of phosphorous
oxide it was found that 0.6760 gram added to 20.698 grams of benzene—in
which the oxide is soluble without change—lowered the freezing-point of
the 3.16 per cent. benzene solution by 0°.68. Since the value of K for
benzene is 49, we have (3.16 × 49)/0.68 = 227, which serves to indicate
that P4O6 is the true molecular formula for phosphorous oxide. This
result is confirmed by vapour-density observations.
The effect of adding a substance to a solvent is to diminish the
vapour pressure of the liquid. Hence, since the boiling-point of a
liquid is that temperature at which the vapour pressure is equal to
the atmospheric pressure, the effect of adding the soluble substance
is to raise the boiling-point, since a higher temperature is required
in order that the pressure of the vapour shall equal that of the
atmosphere. It has been proved that equal volumes of solutions in the
same solvent which have the same boiling-point contain an equal number
of molecules of the dissolved substance.
The equation for the molecular increment of the boiling-point for
a solvent is _d_ = 0.02T²/_w_, in which _d_ is the increment of
the boiling-point caused by the solution of one gram-molecule of a
substance in 100 grams of the solvent, T the _absolute_ boiling-point
of the solvent, and _w_ the heat of vaporisation of the solvent for one
gram. The molecular rise of the boiling-point is therefore independent
of the nature of the dissolved substance.
The molecular weight of the substance _m_ is obtained from the
formula _m_ = _pd_/Δ, in which _p_ = the percentage weight of the
dissolved substance, _d_ = the molecular increment in boiling-point
(0.02T²/_w_), Δ = the observed rise in boiling-point. If the latent
heat of vaporisation of the liquid is unknown, the value of _d_ may be
obtained by preliminary experiments with a substance of known molecular
weight; in this case _d_ = _m_Δ/_p_.
The calculation of the molecular weight _m_ may also be made by the
formula _m_ = K(_s_/ΔL) × 100, in which Δ is the rise in boiling-point,
_s_ the weight of dissolved substance, L the weight of solvent, and K
the molecular boiling-point increment. Convenient forms of apparatus
for using these methods have been devised by Beckmann, and are now in
general use.
* * * * *
From the time of Berzelius, each successive generation of chemists has
striven to better the example of that master of determinative chemistry
in the effort to obtain accurate values for the atomic weights of the
elements.
Among the immediate successors of Berzelius in this work should be
mentioned Turner, Penny, Dumas, and Marignac. Dumas in 1859 published
the results of an extensive revision of the atomic weights of the
elements. On this he based the far-reaching generalisation that, in
the language of Prout, “the combining or atomic weights of bodies bear
certain simple relations to one another, frequently by multiple, and
consequently that many of them must necessarily be multiples of some
one unit.” Dumas further agreed with Prout that “there seems to be no
reason why bodies still lower in the scale than hydrogen (similarly,
however, related to one another, as well as to those above hydrogen)
may not exist, of which other bodies may be multiples, without being
actually multiples of the intermediate hydrogen.”
The Belgian chemist, Stas, who had been associated with Dumas in a
classical determination of the atomic weight of carbon, set himself
to determine, with the highest degree of precision then possible,
the atomic weights of about a dozen of the elements, with a view of
ascertaining (1) whether an atomic weight is a definite and constant
quantity, or whether, as suggested by Marignac, and subsequently by
Crookes, an atomic weight represents “a mean value around which the
actual weights of the atoms vary within certain narrow limits”;
(2) whether, if the atomic weights of the elements are respectively
definite and invariable, the numbers are commensurable as alleged by
Prout and Dumas; and (3) if it should turn out that the numbers are
severally fixed and commensurable, whether this necessarily indicates
that the elements are built up of a primordial matter, the πρώτη ιλη of
the ancients, referred to by Prout in 1816.
Stas devoted many years to the solution of these questions, working
on a scale and with an accuracy and manipulative skill previously
unapproached. The main results of his labour appeared in 1865. He
concluded (1) that the atomic weights of the elements are absolutely
constant values, and are not affected by the nature of the compounds in
which they occur, or the physical conditions of their existence; (2)
that the numbers so obtained are not commensurable: to quote his own
words: “_On doit considérer la loi de Prout comme une pure illusion._”
Hence the elements must, on the basis of Stas’s experimental evidence,
be regarded as “_individualités à part_,” as he expressed it—each a
primordial and unalterable substance.
The appearance of this monumental work, which will ever remain one
of the classics of chemistry, created a great impression. Its effect
persists to this day. It constituted a model and furnished a standard
which each successive worker has striven to emulate, with the result
that atomic weights to-day are among the best ascertained of physical
constants.
Space will not permit of any detailed account of the work done in
connection with atomic weights during the forty-five years which have
elapsed since the publication of Stas’s memoirs; and the reader who
desires fuller information must be referred to the special treatises on
the subject, such as the _Constants of Nature_ of F. W. Clarke, or the
monographs of Meyer and Seubert, Becker, Sebelien, and Van der Plaats.
Reference, however, must be made to the determinations by Lord
Rayleigh, Leduc, Morley, Noyes, Guye, Dixon, and Edgar of values which,
like those of oxygen, hydrogen, nitrogen, silver, and the halogens, are
largely made use of as fiduciary values in atomic-weight work.
Lastly, it should be mentioned, the re-determination of the atomic
weights of the elements with the highest attainable precision and by
the most refined and most modern methods has for some years past been
a special feature of the work of the Harvard Laboratory, under the
direction of Theodore Richards; and some of the most trustworthy and
best established values we possess have been ascertained by him and
his pupils. Atomic weights are of such fundamental importance that the
various nations interested in the pursuit of chemistry have consented
to the establishment of an International Committee, which will take
cognisance of the work done from time to time in this department of
operative chemistry, examine and assess its value, and draw up an
annual report on the subject.
Despite the accumulated testimony of this work in relation to the
validity of the law of the conservation of mass, the sufficiency of the
evidence has now and again been impugned. This aspect of the matter
has within recent years been directly investigated by Landolt; and
as the result of a painstaking series of experiments, in which every
recognised source of error was removed or allowed for, it would appear
that there is absolutely no ground for the belief that there is any
dissipation of mass in the course of, or as the result of, a chemical
change.
CHAPTER V
THE MOLECULAR THEORY OF GASES
The more obvious physical phenomena of gases were, of course, well
known by the middle of the nineteenth century; and the so-called
gaseous laws—the laws of Boyle, Dalton, and Gay Lussac—were universally
accepted by chemists and physicists at that time as fundamental. That
the first two laws were only approximations to truth in a mathematical
sense was also well known; and the experimental labours of Regnault and
Magnus had not only established limits within which they were inexact,
but had, to some extent, indicated the cause of their departure from
the ideal condition. The hypothesis of Avogadro, as already stated,
was practically ignored at this period, or at least its value was
unappreciated until the time of Gerhardt and Laurent, and more
particularly Cannizzaro, who in 1858 pointed out its real meaning and
made it the keystone in the edifice of modern chemistry.
One of the most significant achievements of the last half-century has
been the demonstration that these gaseous laws are interdependent.
Their further study, and in particular the study of their variation
from exact mathematical expression, has led to a conception of the real
nature of a gas which not only comprehends and knits together these
laws, but affords a rational explanation of them. If the laws of Boyle
and Dalton concerning the relations between pressure, temperature, and
gaseous volume are, and must be from the very nature of the case, only
approximations, it follows that the same is equally true of the laws of
Gay Lussac and Avogadro, since these are dependent on the others. The
definite experimental proof that gases do not actually combine in the
precise ratios demanded by the law of Gay Lussac has been forthcoming
only during the last twenty years. It has been found that, instead
of oxygen and hydrogen combining in the exact ratio of one volume of
oxygen to two volumes of hydrogen to form water, as stated by Gay
Lussac, one volume of oxygen combines with, according to Scott, 2.00245
vols.; according to Leduc, 2.0024 vols.; according to Morley, 2.00268
vols. of hydrogen. What is true of the volume-ratios in which oxygen
and hydrogen are actually found to combine, under ordinary conditions,
is no doubt equally true of analogous instances, such as the union
of hydrogen and chlorine to form hydrogen chloride. It follows also
that the extent of the variation from the mathematical expression of
Gay Lussac’s law must get smaller and smaller as the combining gases
approach the condition of the ideal gas—as they do, for example, under
very low pressures. The precise degree of departure from Gay Lussac’s
law is therefore in a sense accidental, and is dependent upon the
conditions under which combination takes place.
The more exact study of the physical phenomena of gases, and in
particular the clearer recognition of the causes which determine the
extent of their departure from the ideal gaseous laws, have afforded
valuable assistance in ascertaining the atomic weights of certain
elements independently of chemical considerations. The processes of
physical measurement have been so refined within recent years that
physical methods of arriving at molecular—and, inferentially, at
atomic—weights, in the case of all elements and compounds which can
be brought into a condition approaching that of the ideal gas, are to
be preferred to the gravimetric methods of analysis or synthesis as
affording the most probable values of the true atomic weights of the
elements. The work of Lord Rayleigh, Leduc, and of Guye and his pupils
on the densities of the gases has furnished us with a series of values
for the atomic weights of a number of the elements which, in point of
accuracy, are as superior to the values of Stas as the values of Stas
were superior to those of his predecessors. Daniel Berthelot pointed
out in 1898 that the true molecular weight of a gas can be deduced
from its density and its observed variations from Boyle’s law under
atmospheric pressure and at very low pressures. Incidentally, the study
of gaseous phenomena has served to place the theory of atoms upon a far
more stable foundation than it occupied half a century ago. How halting
was the adhesion which even some of the most eminent chemists then gave
to this theory was well exemplified by the remarkable lecture given
before the Chemical Society of London in 1869, in which Williamson—one
of the most sturdy champions of Dalton’s doctrine—set forth its true
value.
That a gas may be looked upon as an association of particles—hard
elastic spheres—moving backwards and forwards in right lines with great
velocity, and possessing in the aggregate a very small proportion of
the space through which they travel, was first conceived by Daniel
Bernoulli in 1738. By means of this hypothesis he explained the direct
proportionality between the density and pressure of a gas. If the
gas consists of moving particles, and the pressure which it exerts
on the sides of the containing vessel is due to the impacts of these
particles, it is obvious that by halving the original volume of the
containing space we halve the space through which the particles travel,
and therefore double the number of their impacts in a given time; in
other words, by compressing the gas to half its initial volume we
double the pressure it exerts, which is nothing else than the law of
Boyle. This conception of the nature of a gas is known as the _kinetic
theory of gases_; it was further developed by Waterston in 1845, still
more fully by Clausius in 1857, and was subsequently placed in its
present position by Maxwell and Boltzmann.
That gases actually do move, and at rates depending on their specific
nature, was rendered probable—apart from this explanation of Boyle’s
law—by many phenomena observed by chemists and physicists in the
eighteenth and early part of the nineteenth century. It was known from
the observations of Leslie in 1804 that specifically light gases moved
or diffused faster than heavy gases. Attempts to determine these rates
were made by Schmidt in 1820, and by Graham in 1846, both of whom found
that the rate of movement of a gas was independent of its chemical
nature, and was determined solely by its mass: _gases move at rates
inversely proportional to the square roots of their densities_. The
following table given by Graham shows the experimentally ascertained
relative rates for a number of gases compared with the rates demanded
by the “_law of gaseous diffusion_.” Column one gives the name of the
gas; column two, the observed rate of diffusion; and column three, the
square root of the density of the gas (air = 1):
_Gas._ _Time of diffusion._ _√density._