Determinations of molecular freezing-point depressions by Raoult and others showed that certain substances exerted only about half the osmotic pressure calculated from their known formulæ, whereas others have abnormally high osmotic pressures. The explanation of the discrepancies in the latter case was given in 1887 by Arrhenius, who pointed out that _only those solutions which have abnormally high osmotic pressures are electrically conductive_. This pregnant observation proved to be very fruitful in suggestiveness; and the connection between conductivity and Van ’t Hoff’s theory of solution was developed by Arrhenius into the doctrine of _electrolytic dissociation_ or _ionisation_—one of the most important consequences of Faraday’s electrolytic laws, the work of Hittorf, and the kinetic conceptions of Williamson and Clausius to which the last quarter of a century has given rise. Arrhenius showed that not only were free ions present in an electrically conductive solution before electrolysis, as maintained by Clausius, but that the proportion of molecules dissociated into ions could be calculated from measurements of electrical conductivity, as well as from measurements of osmotic pressure. Both methods give concordant results—a strong confirmation of the validity of the theory. In a solution of common salt, containing a gramme equivalent of that substance in a litre, Arrhenius calculated that only about three tenths of the salt exists as NaCl, the remaining seven tenths being resolved into independent ions of chlorine (chloridion) and sodium (sodion): NaCl⇄[Na·] + Cl´, each moving freely in all directions, like gaseous molecules. On passing the current, electrodes placed in the solution exert a directive action on the free ions, these alone being concerned in determining the conductivity, the un-ionised molecules or the solvent itself exercising no influence. Methods of determining the migration velocity of the ions have been worked out by Hittorf, Kohlrausch, Lodge, and others.
The theory of ionisation affords a satisfactory explanation of many chemical phenomena. It accounts for the characteristic properties of acids, and explains why different acids have varying “strengths” and why a “weak” acid has the same “strength” as the “strong” acid at high equivalent dilutions: in each case the acid is nearly completely ionised—in other words, the “strength” of an acid depends on the concentration of its hydrogen ions. So, too, the “strength” of a base is related to the number of its hydroxyl ions. Aqueous ammonia is relatively a “weak” base—its solution contains few hydroxyl ions. On the other hand, caustic potash is a “strong” base—its solution, on moderate dilution, is almost completely ionised: KOH = K· + OH´, the positive ion being represented by one or more dots, and the negative ion by one or more dashes. The theory accounts, too, for many phenomena in analytical chemistry—such as why magnesia is precipitated by ammonia only in the absence of ammonium chloride, and why sulphuretted hydrogen throws down zinc sulphide in the absence of hydrochloric acid. It also serves to explain many thermo-chemical facts observed by Hess, Thomsen, and others, such as the fact that the heat of neutralisation of the “strong” acids and bases is independent of their nature, and has the uniform value of 13,700 calories, in agreement with the value, as calculated by Van ’t Hoff, for the reaction H· + OH´ = H2O, deduced from Kohlrausch’s measurements of the conductivity of water at varying temperatures.