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of the conception of the chemical atom. At once he saw that the
hypothesis, if true, furnished a marvellous key to secrets of
matter hitherto insoluble—questions relating to the relative
proportions of the atoms themselves. It is known, for example,
that a certain bulk of hydrogen gas unites with a certain bulk of
oxygen gas to form water. If it be true that this combination
consists essentially of the union of atoms one with another (each
single atom of hydrogen united to a single atom of oxygen), then
the relative weights of the original masses of hydrogen and of
oxygen must be also the relative weights of each of their
respective atoms. If one pound of hydrogen unites with five and
one-half pounds of oxygen (as, according to Dalton’s experiments,
it did), then the weight of the oxygen atom must be five and
one-half times that of the hydrogen atom. Other compounds may
plainly be tested in the same way. Dalton made numerous tests
before he published his theory. He found that hydrogen enters
into compounds in smaller proportions than any other element
known to him, and so, for convenience, determined to take the
weight of the hydrogen atom as unity. The atomic weight of
oxygen then becomes (as given in Dalton’s first table of 1803)
5.5; that of water (hydrogen plus oxygen) being of course 6.5.
The atomic weights of about a score of substances are given in
Dalton’s first paper, which was read before the Literary and
Philosophical Society of Manchester, October 21, 1803. I wonder
if Dalton himself, great and acute intellect though he had,
suspected, when he read that paper, that he was inaugurating one
of the most fertile movements ever entered on in the whole
history of science?
Be that as it may, it is certain enough that Dalton’s
contemporaries were at first little impressed with the novel
atomic theory. Just at this time, as it chanced, a dispute was
waging in the field of chemistry regarding a matter of empirical
fact which must necessarily be settled before such a theory as
that of Dalton could even hope for a bearing. This was the
question whether or not chemical elements unite with one another
always in definite proportions. Berthollet, the great co-worker
with Lavoisier, and now the most authoritative of living
chemists, contended that substances combine in almost
indefinitely graded proportions between fixed extremes. He held
that solution is really a form of chemical combination—a
position which, if accepted, left no room for argument.
But this contention of the master was most actively disputed, in
particular by Louis Joseph Proust, and all chemists of repute
were obliged to take sides with one or the other. For a time the
authority of Berthollet held out against the facts, but at last
accumulated evidence told for Proust and his followers, and
towards the close of the first decade of our century it came to
be generally conceded that chemical elements combine with one
another in fixed and definite proportions.
More than that. As the analysts were led to weigh carefully the
quantities of combining elements, it was observed that the
proportions are not only definite, but that they bear a very
curious relation to one another. If element A combines with two
different proportions of element B to form two compounds, it
appears that the weight of the larger quantity of B is an exact
multiple of that of the smaller quantity. This curious relation
was noticed by Dr. Wollaston, one of the most accurate of
observers, and a little later it was confirmed by Johan Jakob
Berzelius, the great Swedish chemist, who was to be a dominating
influence in the chemical world for a generation to come. But
this combination of elements in numerical proportions was exactly
what Dalton had noticed as early as 1802, and what bad led him
directly to the atomic weights. So the confirmation of this
essential point by chemists of such authority gave the strongest
confirmation to the atomic theory.
During these same years the rising authority of the French
chemical world, Joseph Louis Gay-Lussac, was conducting
experiments with gases, which he had undertaken at first in
conjunction with Humboldt, but which later on were conducted
independently. In 1809, the next year after the publication of
the first volume of Dalton’s New System of Chemical Philosophy,
Gay-Lussac published the results of his observations, and among
other things brought out the remarkable fact that gases, under
the same conditions as to temperature and pressure, combine
always in definite numerical proportions as to volume. Exactly
two volumes of hydrogen, for example, combine with one volume of
oxygen to form water. Moreover, the resulting compound gas
always bears a simple relation to the combining volumes. In the
case just cited, the union of two volumes of hydrogen and one of
oxygen results in precisely two volumes of water vapor.
Naturally enough, the champions of the atomic theory seized upon
these observations of Gay-Lussac as lending strong support to
their hypothesis—all of them, that is, but the curiously
self-reliant and self-sufficient author of the atomic theory
himself, who declined to accept the observations of the French
chemist as valid. Yet the observations of Gay-Lussac were
correct, as countless chemists since then have demonstrated anew,
and his theory of combination by volumes became one of the
foundation-stones of the atomic theory, despite the opposition of
the author of that theory.
The true explanation of Gay-Lussac’s law of combination by
volumes was thought out almost immediately by an Italian savant,
Amadeo, Avogadro, and expressed in terms of the atomic theory.
The fact must be, said Avogadro, that under similar physical
conditions every form of gas contains exactly the same number of
ultimate particles in a given volume. Each of these ultimate
physical particles may be composed of two or more atoms (as in
the case of water vapor), but such a compound atom conducts
itself as if it were a simple and indivisible atom, as regards
the amount of space that separates it from its fellows under
given conditions of pressure and temperature. The compound atom,
composed of two or more elementary atoms, Avogadro proposed to
distinguish, for purposes of convenience, by the name molecule.
It is to the molecule, considered as the unit of physical
structure, that Avogadro’s law applies.
This vastly important distinction between atoms and molecules,
implied in the law just expressed, was published in 1811. Four
years later, the famous French physicist Ampere outlined a
similar theory, and utilized the law in his mathematical
calculations. And with that the law of Avogadro dropped out of
sight for a full generation. Little suspecting that it was the
very key to the inner mysteries of the atoms for which they were
seeking, the chemists of the time cast it aside, and let it fade
from the memory of their science.
This, however, was not strange, for of course the law of Avogadro
is based on the atomic theory, and in 1811 the atomic theory was
itself still being weighed in the balance. The law of multiple
proportions found general acceptance as an empirical fact; but
many of the leading lights of chemistry still looked askance at
Dalton’s explanation of this law. Thus Wollaston, though from the
first he inclined to acceptance of the Daltonian view, cautiously
suggested that it would be well to use the non-committal word
“equivalent” instead of “atom”; and Davy, for a similar reason,
in his book of 1812, speaks only of “proportions,” binding
himself to no theory as to what might be the nature of these
proportions.
At least two great chemists of the time, however, adopted the
atomic view with less reservation. One of these was Thomas
Thomson, professor at Edinburgh, who, in 1807, had given an
outline of Dalton’s theory in a widely circulated book, which
first brought the theory to the general attention of the chemical
world. The other and even more noted advocate of the atomic
theory was Johan Jakob Berzelius. This great Swedish chemist at
once set to work to put the atomic theory to such tests as might
be applied in the laboratory. He was an analyst of the utmost
skill, and for years be devoted himself to the determination of
the combining weights, “equivalents” or “proportions,” of the
different elements. These determinations, in so far as they were
accurately made, were simple expressions of empirical facts,
independent of any theory; but gradually it became more and more
plain that these facts all harmonize with the atomic theory of
Dalton. So by common consent the proportionate combining weights
of the elements came to be known as atomic weights—the name
Dalton had given them from the first—and the tangible conception
of the chemical atom as a body of definite constitution and
weight gained steadily in favor.
From the outset the idea had had the utmost tangibility in the
mind of Dalton. He had all along represented the different atoms
by geometrical symbols—as a circle for oxygen, a circle
enclosing a dot for hydrogen, and the like—and had represented
compounds by placing these symbols of the elements in
juxtaposition. Berzelius proposed to improve upon this method by
substituting for the geometrical symbol the initial of the Latin
name of the element represented—O for oxygen, H for hydrogen,
and so on—a numerical coefficient to follow the letter as an
indication of the number of atoms present in any given compound.
This simple system soon gained general acceptance, and with
slight modifications it is still universally employed. Every
school-boy now is aware that H2O is the chemical way of
expressing the union of two atoms of hydrogen with one of oxygen
to form a molecule of water. But such a formula would have had
no meaning for the wisest chemist before the day of Berzelius.
The universal fame of the great Swedish authority served to give
general currency to his symbols and atomic weights, and the new
point of view thus developed led presently to two important
discoveries which removed the last lingering doubts as to the
validity of the atomic theory. In 1819 two French physicists,
Dulong and Petit, while experimenting with heat, discovered that
the specific heats of solids (that is to say, the amount of heat
required to raise the temperature of a given mass to a given
degree) vary inversely as their atomic weights. In the same year
Eilhard Mitscherlich, a German investigator, observed that
compounds having the same number of atoms to the molecule are
disposed to form the same angles of crystallization—a property
which he called isomorphism.
Here, then, were two utterly novel and independent sets of
empirical facts which harmonize strangely with the supposition
that substances are composed of chemical atoms of a determinate
weight. This surely could not be coincidence—it tells of law.
And so as soon as the claims of Dulong and Petit and of
Mitscherlich had been substantiated by other observers, the laws
of the specific heat of atoms, and of isomorphism, took their
place as new levers of chemical science. With the aid of these
new tools an impregnable breastwork of facts was soon piled about
the atomic theory. And John Dalton, the author of that theory,
plain, provincial Quaker, working on to the end in
semi-retirement, became known to all the world and for all time
as a master of masters.
HUMPHRY DAVY AND ELECTROCHEMISTRYDuring those early years of the nineteenth century, when Dalton
was grinding away at chemical fact and theory in his obscure
Manchester laboratory, another Englishman held the attention of
the chemical world with a series of the most brilliant and widely
heralded researches. This was Humphry Davy, a young man who had
conic to London in 1801, at the instance of Count Rumford, to
assume the chair of chemical philosophy in the Royal Institution,
which the famous American had just founded.
Here, under Davy’s
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