THE POPULAR SCIENCE MONTHLY
which has been the philosophy of organic chemistry for the past thirty-five years, and the guiding thought in practically all of the best work in organic chemistry from 1874 to the present time.
The second great work of Van't Hoff had also to do primarily with the chemistry of carbon. In 1867 the Norwegian physicist, Guldberg, and his son-in-law Waage, the chemist, both of Christiania, announced the law of the effect of mass or quantity on chemical reaction—the law of mass action. This was published in the "Announcements" of the University of Christiania, and very little attention was paid to it for some time. Guldberg and Waage applied their law to comparatively few reactions.
Van't Hoff, shortly after the publication of his brief paper of eleven pages in Dutch, on "The Arrangement of the Atoms in Space," took up experimentally the study of the velocities of chemical reactions and the conditions of chemical equilibria, from the standpoint of the law of mass action. He, his assistants and students, carried out an elaborate series of investigations in which the law of mass action was applied to a large number of chemical reactions, and shown to hold. The results of this work were published in French, under the French equivalent of "Studies in Chemical Dynamics." In this work the whole subject of chemical dynamics and chemical equilibrium was placed upon a scientific basis, and for the first time.
The third and greatest work of Van't Hoff had to with the relation between solutions and gases. Through his colleague—the botanist, De Vries, the attention of Van't Hoff was called to the osmotic pressure measurements that had been made by the botanist Wilhelm Pfeffer. A comparison of the results obtained by Pfeffer with the gas pressures exerted by gases containing the same number of gaseous molecules in a given volume that the solution contained dissolved molecules in the same volume, showed that the gas-pressure was exactly equal to the osmotic pressure-—in a word, the laws of gas-pressure apply to the osmotic pressure of solutions.
Van't Hoff showed that the laws of gas-pressure apply to the osmotic pressure of solutions of non-electrolytes, i. e., those substances whose aqueous solutions do not conduct the current. He also pointed out that the laws of gas-pressure do not apply to the osmotic pressure of a single electrolyte—a single acid, base or salt. Arrhenius explained the apparent discrepancy in the case of electrolytes by means of the theory of electrolytic dissociation, which says that acids, bases and salts in aqueous solution are broken down into charged parts or ions.
The question arises why is it so important to have shown that the laws of gas-pressure apply to the osmotic pressure of solutions? We know more about matter in the gaseous state than in any other state of aggregation. We can deal with gases from the standpoint of the