Redox reaction thermodynamic considerations

It follows from the electrochemical mechanism of corrosion that the rates of the anodic and cathodic reactions are interdependent, and that either or both may control the rate of the corrosion reaction. It is also evident from thermodynamic considerations (Tables 1.9 and 1.10) that for a species in solution to act as an electron acceptor its redox potential must be more positive than that of the M /M equilibrium or of any other equilibrium involving an oxidised form of the metal. 

In the presence of trace amounts of iron, superoxide can then reduce Fe3+ to molecular oxygen and Fe2+. The sum of this reaction (13.2) plus the Fenton reaction (13.1) produces molecular oxygen, hydroxyl radical and hydroxyl anion from superoxide and hydrogen peroxide, in the presence of catalytic amounts of iron. This is the Haber-Weiss reaction (13.3), originally described by Haber and Weiss (1934), but manifestly impossible from thermodynamical considerations in the absence of catalytic amounts of redox metals such as iron and copper ... 

In both reactions, electron transfer induces the dissolution of the solid phase i.e., reductive and oxidative dissolution, respectively. Although no kinetic implications follow directly from the thermodynamic considerations, there are cases where the redox rate is related to the redox equilibrium (see e.g., Eq. 9.12). 

The thermodynamic feasibility of redox reactions at the semiconductor-electrolyte interface can be assessed from thermodynamic considerations. Since typical redox potentials for many redox couples encountered in electrolytes of natural or technical systems often lie between the band potentials of typical semiconductors, many electron transfer reactions are (thermodynamically) feasible (Pichat and Fox, 1988). With the right choice of semiconductor material and pH the redox potential of the cb can be varied from 0.5 to 1.5 V and that of the vb from 1 to more than 3.5 V (see Fig. 10.4). 

In many cases, redox reactions that are favorable from a thermodynamic point of view may not actually take place sometimes, the activation energy barriers for such reactions are too high to allow fast transformation, according to the preferred thermodynamic considerations. For example, the complete oxidation of any organic molecule to carbon dioxide and water is thermodynamically favorable. However, such oxidation is not favorable kinetically, which implies that organic molecules— including all forms of living species—are not oxidized immediately this fact explains the ability to sustain life. The reason for this difference between kinetic and thermodynamic considerations, for redox reactions, is partly because redox reactions are relatively slow compared to other reactions and partly due to the fact that, in many cases, reactions are poorly coupled because of slow species diffusion... 

From thermodynamic considerations, the course of such redox reaction > can be expressed in terms of the redox potentials of the two couples Ed and Er... 

Let us now turn to some kinetic considerations of NAC reduction. As an example, consider the time courses of nitrobenzene (NB) concentration in 5 mM aqueous hydrogen sulfide (H2S) solution in the absence and presence of natural organic matter (Fig. 14.7). As is evident, although reduction of NB by H2S to nitrosobenzene and further to aniline (Eq. 14-31) is very favorable from a thermodynamic point of view (see Fig. 14.4), it seems to be an extremely slow process. However, when DOM is added to the solution, reduction occurs at an appreciable rate (Fig. 14.7). In order to understand these findings, some general kinetic aspects of redox reactions involving NACs should be recognized. 

A consideration of these relationships reveals8 that because E° is a thermodynamic parameter and represents an energy difference between two oxidation states and in many cases the spectroscopic or other parameter refers to only one half of the couple, then some special conditions must exist in order for these relationships to work. The special conditions under which these relationships work are that (a) steric effects are either unimportant or approximately the same in both halves of the redox couple and (b) changes in E° and the spectroscopic or other parameters arise mainly through electronic effects. The existence of many examples of these relationships for series of closely related complexes is perhaps not too unexpected as it is likely that, for such a series, the solvational contribution to the enthalpy change, and the total entropy change, for the redox reaction will remain constant, thus giving rise to the above necessary conditions. 

Upon electronic excitation the redox properties of either the electron donor (D) or the acceptor (A) are enhanced. The feasibility of an electron transfer can be estimated from a simple free reaction energy consideration as customary in the frame of the Rehm-Weller approach (Eq. (1)) [11], where Efy2 (P) and 4) represent the oxidation and reduction potential of the donor or the acceptor, respectively. AEexcit stands for the electronic excitation energy, whereas Aiscoui indicates the coulombic interaction energy of the products formed (most commonly radical ions). This simplified approach allows a first approximation on the feasibility of a PET process without considering the more complex kinetics as controlled by the Marcus theory [6c]. For exergonic processes (AG<0) a PET process becomes thermodynamically favorable. 

In several cases it has been found that the oxidation of the redox system occurs entirely via hole transfer directly from the valence band to the reduced form of the couple. Then both processes, oxidation of the redox system and corrosion, proceed independently. This is usually not visible from measurements with an n-type electrode, because the photocurrent is entirely determined by the light intensity. As already mentioned above, p-type electrodes are more suitable, because the current is determined by majority carrier transfer (reaction rate Vf, in Fig. 21). From the thermodynamic point of view, the oxidation of Cu at GaAs is an interesting case. The corresponding current-potential curves are given in Fig. 22 [93]. The corrosion current is not changed upon addition of Cu, i.e. corrosion and redox process are completely independent. In this case, the kinetics of the direct hole transfer is obviously very fast, i.e. the redox current is considerably larger than the corrosion current. Both processes occur indepen-... 

It is important to obtain experimental information on the thermodynamics of electrode processes to ascertain the tendency of a particular reaction to occur under a given set of experimental conditions namely temperature, pressure, system com H)sition and electrode potential. Such information is provided by the standard- or formal-electrode potentials for the redox couple under consideration. Appropriate combinations of these potentials enable the thermodynamics of homogeneous redox processes to be determined accurately. However, such quantities often are subject to confusion and misinterpretation. It is, therefore, worthwhile to outline their significance for simple electrochemical reactions. This discussion provides background to the sections on electrochemical kinetics which follow. The evaluation of formal potentials for various types of electrode-reaction mechanisms is dealt with in 12.3.2.2. 

The lack of photostability of M-type semiconductor electrodes is a severe problem when using aqueous solutions. Holes excited by light excitation and then driven toward the surface can be either used for the decomposition or are transferred to a redox system. In the case of large band-gap oxides, the stability is thermodynamically controlled i.e. the anodic decomposition occurs only at potentials that are considerably more positive than typical redox reactions. In all other cases, the competition between these two processes is kinetically controlled (Memming, 1990, 1994). In several cases, a decrease of the stability of n-type semiconductors with increasing... 

The redox chemistry of nitrogen compounds in aqueous solution is illustrated in the Frost diagram in Fig. 1 (see Topic E5 for construction and use). All oxides and oxoacids are strong oxidizing agents, and all oxidation states except -3, 0 and +5 are susceptible to disproportionation. The detailed reactions are, however, mostly controlled by kinetic rather than thermodynamic considerations. In conjunction with oxidizable groups, as in ammonium nitrate NH4N03 or in... 

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