Redox potentials in natural waters

Geochemists (e.g., Thorstenson et al., 1979 Thorstenson, 1984) have long recognized that at low temperature many redox reactions are unlikely to achieve equilibrium, and that the meaning of Eh measurements is problematic. Lindberg and Runnells (1984) demonstrated the generality of the problem. They compiled from the watstore database more than 600 water analyses that provided at least two measures of oxidation state. The measures included Eh, dissolved oxygen content, concentrations of dissolved sulfate and sulfide, ferric and ferrous iron, nitrate and ammonia, and so on. 

They calculated species distributions for each sample and then computed redox potentials for the various redox couples in the analysis, using the Nernst equation, 

For example, when they found an analysis reporting concentrations of both sulfate and sulfide, they calculated the Nernst Eh for the reaction 

Given a measurement of dissolved oxygen, they similarly computed the Eh corresponding to the reaction 

In this way, they could calculate a redox potential for each redox couple reported for a sample. 

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Hostettler, J. D. (1984). Electrode reaction, aqueous electrons, and redox potentials in natural waters. Am. J. Sci. 284,734-759. 

Several of the key issues are reflected in the debate over the appropriate use of pe to describe redox conditions in natural waters (129-131). The parameter is defined in terms of the activity of solvated electrons in solution (i.e., pe = - log e ), but the species e aq does not exist under environmental conditions to any significant degree. The related concept of pe (132), referring to the activity of electrons in the electrode material, may have a more realistic physical basis with respect to electrode potentials, but it does not provide an improved basis for describing redox transformations in solution. The fundamental problem is that the mechanisms of oxidation and reduction under environmental conditions do not involve electron transfer from solution (or from electrode materials, except in a few remediation applications). Instead, these mechanisms involve reactions with specific oxidant or reductant molecules, and it is these species that define the half-reactions on which estimates of environmental redox reactions should be based. 

Investigations of redox processes in natural water systems have emphasized the disequilibrium behavior of many couples (e.g., 37). The degree of coupling of redox reactions with widely varying rates, and its effect on radionuclide transport in an NWRB needs to be considered. Because of the generally slow kinetics of autoxidation reactions, the potential surface catalyzed reduction of a radionuclide at low temperatures in the presence of trace levels of DO may explain certain sorption data (e.g., 38). 

A few elements—C, N, O, S, Fe, Mn—are predominant participants in aquatic redox processes. Tables 8.6a and 8.6b present equilibrium constants for several couples pertinent to consideration of redox relationships in natural waters and their sediments. Data are taken principally from the second edition of Stability Constants of Metal-lon Complexes and Standard Potentials in Aqueous Solution (Bard et al., 1985). A subsidiary symbol pe (W) is convenient for considering redox situations in natural waters. pe°(W) is analogous to pe except that H" and OH in the redox equilibrium equations are assigned their activities in neutral water. Values for pe°(W) for 25 °C thus apply to unit activities of oxidant and reductant at pH = 7.00. pe°(W) is defined by... 

In addition, because the net reaction at converts Fe to Fe, the measured potential exhibits a slow drift. Such mixed potentials are of little worth in determining equilibrium values. Many important redox couples in natural waters are not electroactive. No reversible electrode potentials are established for N03T-N0 -NH4-H2S or CH4-CO2 systems. Unfortunately, many measurements of Eh (or pe) in natural waters represent mixed potentials not amenable to quantitative interpretation. 

The water stability boundaries and the locus of measured Eh and pH measurements in natural waters, as reported by Baas-Becking et al. (I960), are shown in Fig. 11.3 (see also Fig. 11.4). It has been observed that frequently the Eh values measured with a Pt electrode differ significantly from values computed from Gibbs free energies or standard potentials and solution concentrations. When they exist, there are two important reasons for such differences. These include (1) misbehavior of the Pt or other indicator electrode (2) the irreversibility or slow kinetics of most redox couple reactions and resultant di.sequilibrium between and among different redox couples in the same water and (3) the common existence of mixed potentials in natural waters (see below). 

Stefansson, A., S. Arnorsson, and A. E. Sveinbjornsdottir. 2005. Redox reactions and potentials in natural waters at disequilibrium. Chem. Geol. 221 289-311. 

The solubility of Hg(II) is controlled by chemical speciation in natural waters, and the availability of ligands for complexation shifts dramatically under varying redox conditions (40). Speciation of dissolved Hg(II) in anoxic environments, such as sediments or the hypolimnion, should be strongly influenced by reactions with reduced sulfur (40, 41), whereas organic complexation is potentially important under oxic conditions (42, 43). 

Redox equilibrium is not achieved in natural waters, and no single pe can usually be derived from an analytical data set including several redox couples. The direct measurement of p thus is usually not meaningful because only certain electrochemically reversible redox couples can establish the potential at an electrode (4, 35). However, p is a useful concept that indicates the direction of redox reactions and defines the predominant redox conditions. Defining pe on the basis of the more abundant redox species like Mn(II) and Fe(II) gives the possibility of predicting the equilibrium redox state of other trace elements. The presence of suitable reductants (or oxidants) that enable an expedient electron transfer is, however, essential in establishing redox equilibria between trace elements and major redox couples. Slow reaction rates will in many cases lead to nonequilibrium situations with respect to the redox state of trace elements. 

Except for the Fea+-Fe2+ couple at concentrations greater than about 10-r>M and perhaps the Mn(1V)-Mn2+ couple (3) the over-all redox systems important in natural waters are not electroactive. No reversible electrode potentials are established for the NCV-NCV-NH., S042"-H2S, or CH4-... 

Morris, J.C. and W. Stumm. 1967. Redox equilibria and measurements of potentials in the aquatic environment. In W. Stumm, ed., Equilibrium Concepts in Natural Water Systems, pp. 270-285. American Chemical Society, Washington, DC. 

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