It is evident, from what has been already said, that chemistry and physics are near akin—indeed, they can hardly be separated. Avogadro’s law and spectrum analysis are but two illustrations of the relationship, but many other examples are equal to them in importance. Take, for instance, the action of light upon chemical substances; it may provoke union of elements, or effect the decomposition of compounds; upon the latter phenomenon the art of photography depends. That salts of silver are chemically changed by light was the fundamental observation, and upon this fact most photographic processes, though not all, are founded. Thus light, working as a chemist in the laboratory of the photographic plate, has become the useful servant of all arts, all sciences, and all industries—an indispensable aid to invention and research. On this theme a volume might be written; a bare reference to it must be sufficient here.
Still another branch of chemistry, recently developed but essentially an extension of the theory of valence, is also due to the study of optical relations. That different crystalline bodies differ in their behavior toward polarized light has long been known, and the polariscope is recognized as an instrument of great value in chemical research. To the analysis and valuation of sugars and sirups it is most effectively applied, and commercial transactions of great magnitude depend in part upon its testimony. Here is practical utility, but the development of theory is what concerns us now.
The discovery of isomerism, of the fact that very different compounds might contain the same elements united in the same proportions, was easily interpreted by the theory of valence in a fairly complete and satisfactory way. In the structural formulæ the different atomic groupings were clearly shown, but with one essential limitation—the arrangement was in a single plane. That is, the linking of the atoms was considered, but not their relations to tridimensional space. For the study of reactions, for the classification of compounds, the structural symbols sufficed; but human thought is not so easily satisfied, and more was soon required. One class of isomers was unexplained, and an explanation was demanded.
A typical example of the difficulty was offered by tartaric acid, which exists in two forms differing crystallographically and optically. One form, dissolved in water, twisted a ray of polarized light to the left, the other produced a rotation to the right, while the crystals of the two acids, similar in all other respects, also showed a right- and left-handedness in the arrangement of their planes. The crystal of one variety resembled the other as would its reflection in a mirror—the same, but reversed. These differences, discovered by Pasteur as long ago as 1848, the theory of valence could not explain; to interpret them, and other similar cases, the arrangement of the atoms in space had to be considered.
In 1874 two chemists, Van t’Hoff and Lebel, working independently, offered a solution of the problem, and stereochemistry, the chemistry of molecular structure in three dimensions, was founded. They proposed, in effect, to treat the carbon atom essentially as a tetrahedron, the four angles corresponding to the four units of valence or bonds of affinity. They then studied the linking or union of such tetrahedra, and found that with their aid the formulæ for tartaric acid could be developed in different ways, showing right- and left-handed atomic groupings. Other similar compounds were equally explicable. Thus the definite conception of a tridimensional, geometric atom led to a new development of structural formulæ, from which many discoveries have already proceeded. The fruitfulness of the speculation vindicates its use, but it is only the first step in a method of research which must in time be applied to all of the chemical elements. Probably the study of crystalline form will be connected with these chemico-structural expressions, and from the union some greater generalizations will be born. From the geometry of the crystal to the geometry of the molecule there must be some legitimate transition. With all their utility, our present conceptions of chemical structure are incomplete; they represent only portions or special phases of some great general law, but so far as they go, properly used, they are valid.
But light is not the only physical force involved in chemical changes; heat and electricity are far more important. Heat, in particular, is essential to every chemical operation; it provokes combination and effects decomposition; it appears in one reaction and vanishes in another; apart from thermal phenomena the science of chemistry could not exist. From the very beginnings of chemistry this interdependence has been recognized, and its study has led to notable discoveries and to great enlargements of resource. In the theory of phlogiston the connection between heat and chemical change was crudely stated, and when Lavoisier saw that combustion was oxidation, thermochemistry began to exist.
In every chemical change a definite amount of heat is either liberated or absorbed—a distinct, measurable quantity. This fact was established by Hess in 1840, and since then the thermal values of many reactions have been determined, notably by Thomsen in Denmark and Berthelot in France. The data are already numerous, but as yet they have not been co-ordinated into any general law. They are in great measure the raw material with which some future scholar is to build. One fact, however, is already clear—namely, that the heat of formation or of combustion of any compound is conditioned by its structure. Two isomeric substances may differ widely in their calorific constants, an observation which has repeatedly been verified. Thus the conception of structure, of atomic grouping, appears again a chief factor in a set of unsolved problems.
In the relations of chemistry to heat perhaps the greatest advances have been made in the extension of our resources, particularly in regard to the development and control of temperatures. At the beginning of the century the range of temperatures available to the chemist was narrowly limited—from the freezing point of mercury at one end to the heat of a blast furnace at the other. His command of heat and cold are now vastly greater than then, and the steps which have been taken are worth tracing.
At the lower end of the scale the greatest progress has been made through the liquefaction of gases. When a liquid evaporates, heat is absorbed, or, reversely stated, cold is produced, and the more rapid the evaporation the greater is the cooling effect. A command of more volatile liquids is therefore a command of cold, and the liquefied gases represent the extreme limit of our power in that direction.
Near the beginning of the century, by combining cold and pressure, sulphurous acid and chlorine were reduced to the liquid state. In 1823 Faraday succeeded in liquefying still other gases, and in 1835 Thilorier went even further and reduced carbonic acid to a snowlike solid. Liquid chlorine, sulphurous acid, and carbonic acid, stored in strong cylinders of steel, are now commercial products, manufactured and sold in large quantities like any other merchandise. They can be transported to long distances and kept indefinitely, to the great convenience of chemists and the furtherance of research.