Interpretation and measurement of redox intensity in natural waters

Abstract

Frevert deserves credit for proposing—for equilibrium systems—a distinction between a conceptually defined redox intensity, pe, and an operationally defined redox condition under stationary states, pe, as given by the response of a sensor electrode, and for pointing out that pe need not relate to pe. We would like to re-emphasize (1, 2, 3) that in defining a redox intensity, pe=−log{e}, we have treated the electron conceptually as a basic redox component which, as a species in aqueous solution, does not have an existence of its own. Morel (4) has elaborated on the use of the electron as a (phase rule) component in redox reactions. As he shows, it obviously can be treated equivalent to O2, i.e. O2=(H+)−4(e−)−4(H2O)2. We define (3) pe as “the hypothetical electron activity at equilibrium which measures the relative tendency of a solution to accept or transfer electrons”. This free energy change ΔG can be expressed as a redox potential (electrode potential) in volts (i.e., as a free energy change per mole of electrons associated with a given reduction). Electron activities may be defined in any equilibrium systems where the free activities of reductants {Red}, and oxidants {Ox}, are defined. Thus, pe (like pH) is a derivative form of free energy. Using electrons in redox reactions and as components does not at all imply that such electrons exist as species in waters. In the compilation of “Stability Constants” of the Chemical Society (London), Sillen and Martell (1964) treat the electron as an inorganic ligand and establish an electron activity scale that corresponds to the definition given.