[525] The values of this and similar equilibrium constants are derived by means of Nernst's formula (see below) for the potential difference between an element and solutions of its ions. The derivation involves the assumption that this formula expresses correctly the relation between the potential change and the concentration change at all concentrations. This assumption appears to be justified by all experimental indications thus far observed. The constants are of importance, primarily, for the calculations which can be made with their aid (see below), and may, conservatively, be considered to be essentially "calculation factors" ("Rechengrössen," according to Haber. See pp. 232–7, Chapter XII). The constants may be expressed, as in the text, in terms of (molar) ‹concentrations› of the ions, or in terms of the ‹osmotic pressures› of the ions, a molar solution at 0° producing an osmotic pressure of 22.4 atmospheres. Where osmotic pressure and concentration are not strictly proportional (‹e.g.› for concentrated solutions), the osmotic pressure, rather than the concentration, is the determining factor and, when known, is used in exact calculations. The plan, pursued in the text, is adopted in order to express these constants in the terms used for all the other equilibrium constants. It should be recalled (‹e.g.› p. 30) that in calculations, in general, where pressure and concentration are not strictly proportional, the pressure is the determining factor. A third method of expressing the solution-tension relations consists in giving the ‹potential differences›, which exist ‹between elements› and solutions of their ‹ions, in which the ions have unit (molar) concentration›. These potential differences are ‹functions› of the solution-tension constants, as will be discussed below, and the constants, in terms of concentrations or osmotic pressures, may be easily calculated, from the potential differences, with the aid of this function (see below, and see the table at the end of Chapter XV). [526] According to Wilsmore's tabulation (‹Z. phys. Chem.›, «36», 92 (1901)), the potential difference ε_{Cu, Cu^{2+}} of copper against a 0.5 molar solution of cupric sulphate, in which [Cu^{2+}] = 0.11, is +0.584 volt. Inserting these values for [Cu^{2+}] and ε_{Cu, Cu^{2+}} in the equation ε_{Cu, Cu^{2+}} = (0.0575 / 2) log([Cu^{2+}] / K) (see below) and solving the equation for K, we find K = 8E−22. For [Cu^{2+}] = 0.24, ε_{Cu, Cu^{2+}} is +0.594 volt and K = 8E−22. In regard to the convention determining the signs used (in the present case ε_{Cu, Cu^{2+}} is ‹positive›), see the footnote below, p. 262, and in regard to the definition of zero potential, to which the potential differences used in this book refer, see the table and summary at the end of Chapter XV. [527] Nernst, ‹loc. cit.›, p. 151.