Note 550 — Chapter XIV · OXIDATION AND REDUCTION REACTIONS. I

The constant is calculated from the data of Küster and Crotogino on the potential of solutions of iodine in potassium iodide [Z. anorg. Chem., 23, 88 (1900)]. Owing to the formation of complex ions I3−, for which due allowance has not been made in the calculation, and owing to some uncertainty as to the vague definition of the concentration of iodine used, the estimation of the constant can only be considered a rough one. The value given expresses the order of the equilibrium ratio sufficiently well for our present purposes. In a recent paper, Bray and MacKay [J. Am. Chem. Soc., 32, 914 (1910)] have determined the constant for the formation of the complex ion according to I3− ⇄ I2 + I−, which might be used to correct the data of Küster and Crotogino; but in view of other uncertainties and inaccuracies, the correction has not been considered advisable.

Several related methods may be used to calculate the equilibrium constant for [I⁻]² : [I₂] = K from the data of Küster and Crotogino. Perhaps the simplest method is the following: A solution of iodine ([I] = 1 / 32 normal, and therefore [I₂] = 1 / 64 molar) in 1/8 molar potassium iodide, in which, the degree of ionization being taken into account, [I⁻] = 0.109, was observed to show a potential ε_{I2, I⁻} = +0.860 (the convention as to signs, discussed on p. 261, is used here and the potential, observed against a so-called "calomel electrode," is reduced to the so-called "absolute potential"; cf. Le Blanc, Lehrbuch der Elektrochemie, p. 214). Now, there must be a certain concentration of iodide-ion, which we will call [C], with which iodine of the above concentration would be directly in equilibrium and would give no potential at all (cf. pp. 261 and 258 in regard to copper). With a change in the concentration of the iodide-ion, a potential would be produced according to ε_{I2, I⁻} = 0.0575 log([C] / [I⁻]). This relation is of exactly the same nature as that developed for the potential of copper plates, immersed in solutions of cupric-ion of different concentrations (but see footnote 1, p. 261, concerning the sign of the new relation). In the present case, we are dealing with univalent ions, I⁻, in place of bivalent ions Cu²⁺, and the factor 0.0575 is used instead of 0.0575 / 2 (see p. 261). If we insert the observed values, [I⁻] = 0.109 and ε = 0.860, of the experiment described above, into the equation ε_{I2, I⁻} = 0.0575 log([C] / [I⁻]) and solve the equation for [C], we find [C] = 10¹⁴. That means, 1 / 64 molar iodine would be directly in equilibrium with a concentration of iodide-ion = 10¹⁴ (if this value is inserted for [I⁻] in the logarithmic equation, the potential is found to be 0). For the condition of equilibrium for I₂ ⇄ 2 I⁻, according to [I⁻]² : [I₂] = K, we have then (10¹⁴)² : (1 / 64) = K = 6.4E29. Similarly, for [I⁻] = 0.109 and [I₂] = 1 / 512 the potential ε = 0.831 is observed, and the equilibrium constant is found to be 5.1E29. When [I⁻] = 0.109 and [I₂] = 1 / 128, the potential is 0.850 and the constant is calculated to be 5.3E29. The mean value for K is 5.6E29. In these calculations, the formation of ions I₃⁻, affecting the values for [I⁻] and [I₂], has not been considered, and there is some doubt whether the concentrations of iodine, given by Küster and Crotogino, do not represent [I₂] rather than [I], as assumed in the calculations. If the former be the case, the mean value of the above experiments would be 2.8E29. The value, used in the text, is considered sufficiently accurate for the purposes of this book.