A History of Science, vol 4 by Henry Smith Williams (the two towers ebook .TXT) 📕

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.

During 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