Treatment of a solution of arsenic acid with a concentrated acid, yielding a large concentration of hydrogen-ion, would carry the series of actions, represented in the above ionization equation for arsenic acid, decidedly toward the right—suppressing the arseniate-ion AsO_{4}^{3−} and increasing the concentration of the arsenic-ion As^{5+}. Since we cannot apply the equilibrium laws (or the principle of the solubility-product) to solutions as concentrated as the one under discussion, a quantitative theoretical treatment of the subject cannot be given. The following may be suggested: The action of the acid would be favorable to the precipitation of As_{2}S_{5} by suppressing the arseniate-ion AsO_{4}^{3−} and thus increasing the concentration of the hydroxide As(OH)_{5}, available for ionization as a base and for the production of the ion As^{5+}. But the further favorable effect of the hydrochloric acid, in converting the hydroxide into a salt AsCl_{5} and increasing thereby the concentration of As^{5+}, would be ‹very largely offset› by the action of the acid in suppressing the [p250] sulphide-ion ([S^{2−}] = ‹k› / [H^{+}]^2; see p. 201). For systems to which the equilibrium laws could be applied, the concentration of As^{5+} (except for the suppression of the ion AsO_{4}^{3−}) would grow, approximately, with the fifth power[512] of the concentration of the hydrogen-ion, and the concentration of the sulphide-ion would decrease, approximately, proportionally to the square of the concentration of the hydrogen-ion. Further, the precipitation of As_{2}S_{5}, in a system to which the principle of the solubility-product were applicable, would depend on the relation of the product [As^{5+}]^2 × [S^{2−}]^5 to the solubility-product constant; it is evident that the value for [As^{5+}]^2 would ‹increase› proportionally to the tenth power of [H^{+}] and the value of [S^{2−}] ‹decrease› proportionally to the tenth power of the same factor [H^{+}]. The two effects would consequently offset each other under such conditions. However, the equilibrium laws cannot legitimately be applied to such concentrated solutions and the relation has been developed only to indicate opposing factors, which must be taken into account. An experimental study of the problem would be extremely interesting.[513] Since it involves the question of the ‹minute basic ionization of a moderately strong acid› (H_{3}AsO_{4}), which may be open to ‹measurement› (see Chap. XVI), the problem is one of particular interest and importance.