«The Solubility- or Ion-Product Principle.»—We may formulate this important conclusion by stating, that, ‹in saturated solutions of silver acetate, the product of the concentrations of its ions has a constant value at a given temperature›. Analogous relations may be developed for the saturated solutions of other difficultly soluble ionogens. The constant has been called the ‹solubility-product constant› or the ‹ion-product constant› of the ‹ionogen›. For salts like lead iodide PbI_{2}, silver chromate Ag_{2}CrO_{4}, etc., each molecule of which forms more than one of a given ion, the concentration of such an ion is raised, in the solubility-product, to the power corresponding to the number of ions of this kind formed from a single molecule of the electrolyte.[292] Thus, lead iodide ionizes according to the equation PbI_{2} ⇄ Pb^{2+} + 2 I^{−} and, for a saturated solution of lead iodide[293] at a given temperature, [Pb^{2+}] × [I^{−}]^2 = K. For silver chromate, ionizing according to the equation Ag_{2}CrO_{4} ⇄ 2 Ag^{+} + CrO_{4}^{2−}, the form of the solubility-product equation is [Ag^{+}]^2 × [CrO_{4}^{2−}] = K. In general, for a saturated solution of a difficultly soluble salt at a given temperature ‹the product of the ion concentrations, each raised to the power corresponding to the number of that kind of ion formed from the ionization of one molecule of the salt, is a constant›. «Criticism of the Derivation of the Principle.»—Nernst developed this important relation in 1889, shortly after the theory of ionization was formulated. Since then, however, the soundness of the theoretical development, on which it was based, has been rendered open to question in a way that could hardly have been foreseen at that early stage in the development of the theory of ionization. In the first place, it is known now that the ionization of easily ionizing substances (strong electrolytes) does [p142] not conform to the law of chemical equilibrium (‹vide› p. 108); as far as our present knowledge goes, the ratio in equation I is not a constant, but grows ‹larger› with an increasing total concentration of good electrolytes. In the present case, this total concentration may be increased by the introduction of ‹foreign salts›.[294] In the second place, the second fundamental principle used, the principle of the constant solubility of the dissolved molecular or non-ionized salt, as expressed in equation III, was questioned and disproved by Arrhenius in 1899. The molecular solubility depends on the total concentration of salts in the solution and, in general, decreases with increasing concentration of the total dissolved salts. This result does not invalidate the law of physical equilibrium; it merely means that the presence of salts, especially in appreciable quantities, modifies the nature of the solvent and changes its dissolving power, much as we have different dissolving power shown by different pure solvents, such as water and alcohol.