Journal of the Chemical Society, Faraday Transactions
Journal of the Chemical Society, Faraday Transactions; was published from 1990 - 1998. In 1999 it merged with a number of European chemical society physical Chemistry journals to become Physical Chemistry Chemical Physics.
Paper
J. Chem. Soc., Faraday Trans., 1998, 94, 2763 - 2767, DOI: 10.1039/a804551h
The equilibrium between the oxidation of hydrogen peroxide by oxygen and the dismutation of peroxyl or superoxide radicals in aqueous solutions in contact with oxygen
Jerzy Petlicki and Theo G. M. van de Ven
The equilibrium O2+HO2-+OH-=2O2-+H2O, which plays an important role in the chemistry of H2O2 in aqueous media, including H2O2 bleaching of wood pulps and other materials, has been overlooked in the literature. Usually only the backward reaction is considered. We have calculated the equilibrium constant of this reaction from literature data on reversible electrochemical cell measurements without a liquid junction. Re-analysis of these data reveals the reason for a long-standing discrepancy between experiments and predictions based on thermal and equilibria data of the National Bureau of Standards: since HO2- is in equilibrium with O2-, the actual HO2- concentration is less than that assumed. We find that the standard redox potential of the oxygen/superoxide couple is -0.137 V, different from the widely quoted literature value of -0.33 V in irradiated solutions. Re-evaluation of pulse radiolysis equilibria on duroquinone-superoxide yields a redox potential for the O2(g)/O2(aq)- couple of ca. -0.13 V. The standard Gibbs energy of formation of the superoxide radical is found to be 13.18 kJ mol-1 (3.15 kcal mol-1) and that of the peroxyl radical -14.69 kJ mol-1 (-3.51 kcal mol-1). Hydrogen peroxide, H2O2, is one of the products formed during any oxidation by oxygen in the presence of water. The superoxide radical, O2-, and the peroxyl radical, HO2, are recognized intermediates in the oxygen/hydrogen peroxide couple–4 The number of mechanisms, explanations and speculations in wet oxygen and hydrogen peroxide chemistry since Priestley (1774) and Thenard (1818) discovered these compounds are simply uncountable. A large number of reactions involving H2O2 remain poorly understood or unexplained. Some examples are the stability of pure aqueous solutions of H2O2 and its catalytic decomposition, the ferryl and hydroxyl pathway in the Fenton reactio–8 and in enzymatic reactions the analytical metho for determining H2O2 concentrations with Ce4+/Ce3+, the stability of the Fe2+phenanthroline comple used as redox indicator in this method, the stability of ortho- and para-hydroquinones formed in high yields in the Dakin reactio in alkaline peroxide solutions, and H2O2 bleaching of wood pulps and other materials.In reanalysing literature data on H2O2 reactions it became clear to us that an important equilibrium was overlooked in the literature. In aqueous alkaline solutions of H2O2 in contact with oxygen or air, the following equilibrium may occur: provided that O2(g) is in equilibrium with dissolved oxygen.Since, in all reactions involving aqueous H2O2 solutions studied in the literature, the solutions contained dissolved oxygen, either from its decomposition or from contact with air, equilibrium (1) plays an important role in the chemistry of H2O2 and O2 in aqueous solutions.Before exploring the consequences of equilibrium (1), one fundamental question must be answered: What is the value of the equilibrium constant, K1 of the above reaction? The key to answering this question is the value of the Gibbs energy of formation,
fG° of the superoxide radical or, equivalently, the standard redox potential E° of the oxygen/superoxide couple which plays an important role in the wet chemistry and biochemistry of oxygen. Obviously in determining these values one must use well documented standards and data consistent with chemical evidence.From the equilibrium constant one can calculate the maximum possible concentrations of superoxide and peroxyl radicals as a function of pH. In the pas only the backward reaction, i.e. the disproportionation of two free radicals, was considered, while the forward reaction, i.e. the oxidation of H2O2 by O2, was neglected. As we will show here, with this equilibrium we can explain a long-standing discrepancy between electrochemical cell measurements on H2O2 alkaline solutions and predictions based on National Bureau of Standards (NBS) data. Standard redox potential of the O2/O2- couple from cell measurements
