[536] Küster, Z. Elecktrochem., 4, 110 and 503 (1897).
[537] In a solution of zinc sulphate in which [Zn2+] = 0.114, the potential εZn, ZnSO4= −0.514 (the minus sign indicates that the metal named first in the subscript has a negative charge). Inserting the values for [Zn2+] and εZn, ZnSO4 in the general equation given on p. [261], and solving for K, we find K = 1017. For [Zn2+] = 0.022 and εZn, ZnSO4 = −0.535, we find K = 1016.8. (Cf. Wilsmore's tables, loc. cit.)
[538] Equilibrium will be established whenever the potential of the system is equal to 0. The potential of the system may be calculated according to the equation (see footnote 1, p. [262])
εCu, Zn = εCu, CuSO4 − εZn, ZnSO4 =
(0.0575 / 2) [log(Cu2+ / KCu) − log(Zn2+ / KZn)].
The potential εCu, Zn is 0 whenever [Cu2+] / KCu = [Zn2+] / KZn, i.e. when [Zn2+] / [Cu2+] = KZn / KCu.
For ions of different valence, such as silver and cupric ions, the equilibrium equation assumes a somewhat less simple form. For Cu ↓ + 2 Ag+ ⇄ 2 Ag ↓ + Cu2+, we have [Ag+]2 / [Cu2+] = (KAg)2 / KCu.
[539] Vide Ostwald's Lehrbuch der allgemeinen Chemie, 2d Ed., Vol. II, p. 874, for the historical data on this action. Vide Küster's experiments, Z. Elektrochem., 4, 503 (1897).
[540] See the footnote, p. [267], in regard to the form the equilibrium ratio assumes when metals producing ions of different valence are used.
[541] The value of the constant is calculated from the data given by Peters, Z. phys. Chem., 26, 193 (1898).
[542] The fact that this equilibrium relation has been proved to hold for the action Fe2+ ⇄ Fe3+ and that it must be taken into account in all oxidation-reduction reactions involving these ions, in no wise excludes the possibility that other equilibrium relations can also exist between ferrous and ferric compounds. For instance, ferrous hydroxide Fe(OH)2 may well have a characteristic tendency of its own to assume a further positive charge (lose an electron) according to Fe(OH)2 ⇄ Fe(OH)2+, the potential of which action may, under given conditions, be a main determining factor in the course of an action, e.g. in alkaline mixtures. It is not impossible, even, that we also must consider negative ions FeO22− and their tendency to be oxidized. Evidence would suggest that ferrous hydroxide, or its negative ion FeO22−, may have, indeed, a very great tendency to be oxidized, possibly much greater than the tendency of Fe2+ to form Fe3+. (Cf. Manchot, Z. anorg. Chem., 27, 419 (1901), and McCoy and Bunzel, J. Am. Chem. Soc., 31, 370 (1909)). Closer investigations of these relations, from a quantitative viewpoint, would probably determine this question and bring exceedingly important relations to light.