A History of Science, vol 4 - Henry Smith Williams (epub e ink reader .txt) 📗

that could not be imitated in the domain of

the non-living. It was regarded almost as an axiom of chemistry

that no organic compound whatever could be put together from its

elements—synthesized—in the laboratory. To effect the synthesis

of even the simplest organic compound, it was thought that the

“vital force” must be in operation.

 

Therefore a veritable sensation was created in the chemical world

when, in the year 1828, it was announced that the young German

chemist, Friedrich Wohler, formerly pupil of Berzelius, and

already known as a coming master, had actually synthesized the

well-known organic product urea in his laboratory at Sacrow. The

“exception which proves the rule” is something never heard of in

the domain of logical science. Natural law knows no exceptions.

So the synthesis of a single organic compound sufficed at a blow

to break down the chemical barrier which the imagination of the

fathers of the science had erected between animate and inanimate

nature. Thenceforth the philosophical chemist would regard the

plant and animal organisms as chemical laboratories in which

conditions are peculiarly favorable for building up complex

compounds of a few familiar elements, under the operation of

universal chemical laws. The chimera “vital force” could no

longer gain recognition in the domain of chemistry.

 

Now a wave of interest in organic chemistry swept over the

chemical world, and soon the study of carbon compounds became as

much the fashion as electrochemistry had been in the, preceding

generation.

 

Foremost among the workers who rendered this epoch of organic

chemistry memorable were Justus Liebig in Germany and Jean

Baptiste Andre Dumas in France, and their respective pupils,

Charles Frederic Gerhardt and Augustus Laurent. Wohler, too,

must be named in the same breath, as also must Louis Pasteur,

who, though somewhat younger than the others, came upon the scene

in time to take chief part in the most important of the

controversies that grew out of their labors.

 

Several years earlier than this the way had been paved for the

study of organic substances by Gay-Lussac’s discovery, made in

1815, that a certain compound of carbon and nitrogen, which he

named cyanogen, has a peculiar degree of stability which enables

it to retain its identity and enter into chemical relations after

the manner of a simple body. A year later Ampere discovered that

nitrogen and hydrogen, when combined in certain proportions to

form what he called ammonium, have the same property. Berzelius

had seized upon this discovery of the compound radical, as it was

called, because it seemed to lend aid to his dualistic theory. He

conceived the idea that all organic compounds are binary unions

of various compound radicals with an atom of oxygen, announcing

this theory in 1818. Ten years later, Liebig and Wohler undertook

a joint investigation which resulted in proving that compound

radicals are indeed very abundant among organic substances. Thus

the theory of Berzelius seemed to be substantiated, and organic

chemistry came to be defined as the chemistry of compound

radicals.

 

But even in the day of its seeming triumph the dualistic theory

was destined to receive a rude shock. This came about through

the investigations of Dumas, who proved that in a certain organic

substance an atom of hydrogen may be removed and an atom of

chlorine substituted in its place without destroying the

integrity of the original compound—much as a child might

substitute one block for another in its play-house. Such a

substitution would be quite consistent with the dualistic theory,

were it not for the very essential fact that hydrogen is a

powerfully electro-positive element, while chlorine is as

strongly electro-negative. Hence the compound radical which

united successively with these two elements must itself be at one

time electro-positive, at another electro-negative—a seeming

inconsistency which threw the entire Berzelian theory into

disfavor.

 

In its place there was elaborated, chiefly through the efforts of

Laurent and Gerhardt, a conception of the molecule as a unitary

structure, built up through the aggregation of various atoms, in

accordance with “elective affinities” whose nature is not yet

understood A doctrine of “nuclei” and a doctrine of “types” of

molecular structure were much exploited, and, like the doctrine

of compound radicals, became useful as aids to memory and guides

for the analyst, indicating some of the plans of molecular

construction, though by no means penetrating the mysteries of

chemical affinity. They are classifications rather than

explanations of chemical unions. But at least they served an

important purpose in giving definiteness to the idea of a

molecular structure built of atoms as the basis of all

substances. Now at last the word molecule came to have a distinct

meaning, as distinct from “atom,” in the minds of the generality

of chemists, as it had had for Avogadro a third of a century

before. Avogadro’s hypothesis that there are equal numbers of

these molecules in equal volumes of gases, under fixed

conditions, was revived by Gerhardt, and a little later, under

the championship of Cannizzaro, was exalted to the plane of a

fixed law. Thenceforth the conception of the molecule was to be

as dominant a thought in chemistry as the idea of the atom had

become in a previous epoch.

CHEMICAL AFFINITY

Of course the atom itself was in no sense displaced, but

Avogadro’s law soon made it plain that the atom had often usurped

territory that did not really belong to it. In many cases the

chemists had supposed themselves dealing with atoms as units

where the true unit was the molecule. In the case of elementary

gases, such as hydrogen and oxygen, for example, the law of equal

numbers of molecules in equal spaces made it clear that the atoms

do not exist isolated, as had been supposed. Since two volumes

of hydrogen unite with one volume of oxygen to form two volumes

of water vapor, the simplest mathematics show, in the light of

Avogadro’s law, not only that each molecule of water must contain

two hydrogen atoms (a point previously in dispute), but that the

original molecules of hydrogen and oxygen must have been composed

in each case of two atoms–else how could one volume of oxygen

supply an atom for every molecule of two volumes of water?

 

What, then, does this imply? Why, that the elementary atom has

an avidity for other atoms, a longing for companionship, an

“affinity”—call it what you will—which is bound to be satisfied

if other atoms are in the neighborhood. Placed solely among

atoms of its own kind, the oxygen atom seizes on a fellow oxygen

atom, and in all their mad dancings these two mates cling

together—possibly revolving about each other in miniature

planetary orbits. Precisely the same thing occurs among the

hydrogen atoms. But now suppose the various pairs of oxygen atoms

come near other pairs of hydrogen atoms (under proper conditions

which need not detain us here), then each oxygen atom loses its

attachment for its fellow, and flings itself madly into the

circuit of one of the hydrogen couplets, and—presto!—there are

only two molecules for every three there were before, and free

oxygen and hydrogen have become water. The whole process, stated

in chemical phraseology, is summed up in the statement that under

the given conditions the oxygen atoms had a greater affinity for

the hydrogen atoms than for one another.

 

As chemists studied the actions of various kinds of atoms, in

regard to their unions with one another to form molecules, it

gradually dawned upon them that not all elements are satisfied

with the same number of companions. Some elements ask only one,

and refuse to take more; while others link themselves, when

occasion offers, with two, three, four, or more. Thus we saw that

oxygen forsook a single atom of its own kind and linked itself

with two atoms of hydrogen. Clearly, then, the oxygen atom, like

a creature with two hands, is able to clutch two other atoms.

But we have no proof that under any circumstances it could hold

more than two. Its affinities seem satisfied when it has two

bonds. But, on the other hand, the atom of nitrogen is able to

hold three atoms of hydrogen, and does so in the molecule of

ammonium (NH3); while the carbon atom can hold four atoms of

hydrogen or two atoms of oxygen.

 

Evidently, then, one atom is not always equivalent to another

atom of a different kind in combining powers. A recognition of

this fact by Frankland about 1852, and its further investigation

by others (notably A. Kekule and A. S. Couper), led to the

introduction of the word equivalent into chemical terminology in

a new sense, and in particular to an understanding of the

affinities or “valency” of different elements, which proved of

the most fundamental importance. Thus it was shown that, of the

four elements that enter most prominently into organic compounds,

hydrogen can link itself with only a single bond to any other

element—it has, so to speak, but a single hand with which to

grasp—while oxygen has capacity for two bonds, nitrogen for

three (possibly for five), and carbon for four. The words

monovalent, divalent, trivalent, tretrava-lent, etc., were coined

to express this most important fact, and the various elements

came to be known as monads, diads, triads, etc. Just why

different elements should differ thus in valency no one as yet

knows; it is an empirical fact that they do. And once the nature

of any element has been determined as regards its valency, a most

important insight into the possible behavior of that element has

been secured. Thus a consideration of the fact that hydrogen is

monovalent, while oxygen is divalent, makes it plain that we must

expect to find no more than three compounds of these two

elements—namely, H—O—(written HO by the chemist, and called

hydroxyl); H—O—H (H2O, or water), and H—O—O—H (H2O2, or

hydrogen peroxide). It will be observed that in the first of

these compounds the atom of oxygen stands, so to speak, with one

of its hands free, eagerly reaching out, therefore, for another

companion, and hence, in the language of chemistry, forming an

unstable compound. Again, in the third compound, though all hands

are clasped, yet one pair links oxygen with oxygen; and this also

must be an unstable union, since the avidity of an atom for its

own kind is relatively weak. Thus the well-known properties of

hydrogen peroxide are explained, its easy decomposition, and the

eagerness with which it seizes upon the elements of other

compounds.

 

But the molecule of water, on the other hand, has its atoms

arranged in a state of stable equilibrium, all their affinities

being satisfied. Each hydrogen atom has satisfied its own

affinity by clutching the oxygen atom; and the oxygen atom has

both its bonds satisfied by clutching back at the two hydrogen

atoms. Therefore the trio, linked in this close bond, have no

tendency to reach out for any other companion, nor, indeed, any

power to hold another should it thrust itself upon them. They

form a “stable” compound, which under all ordinary circumstances

will retain its identity as a molecule of water, even though the

physical mass of which it is a part changes its condition from a

solid to a gas from ice to vapor.

 

But a consideration of this condition of stable equilibrium in

the molecule at once suggests a new question: How can an

aggregation of atoms, having all their affinities satisfied, take

any further part in chemical reactions? Seemingly such a

molecule, whatever its physical properties, must be chemically

inert, incapable of any atomic readjustments. And so in point of

fact it is, so long as its component atoms cling to one another

unremittingly. But this, it appears, is precisely what the atoms

are little prone to do. It seems that they are fickle to the last

degree in their individual attachments, and are as prone to break

away