Standard apparent reduction potential tables

Since tables of standard apparent reduction potentials and standard transformed Gibbs energies of formation contain the same basic information, there is a question as to whether this chapter is really needed. However, the consideration of standard apparent reduction potentials provides a more global view of the driving forces in redox reactions. There are two contributions to the apparent equilibrium constant for a biochemical redox reaction, namely the standard apparent reduction potentials of the two half-reactions. Therefore it is of interest to compare the standard apparent reduction potentials of various half reactions. 

Table 9.2 Standard Apparent Reduction Potentials E ° in Volts at 298.15 K and 1 bar as a Function of pH and Ionic Strength...

Table 9.3 Standard Transformed Gibbs Energies (in kJ moE ) of Reactions and Standard Apparent Reduction Potentials (in volts) at 289.15 K, 1 bar, pH 7, and Ionic Strength 0.25 M for Reactions Involved in the Methane Monooxygenase Reaction...

Table 9.4 Standard Apparent Reduction Potentials E ° (in volts) at 298.15 K of Half-reactions Involving Reactants with Multiple Species...

The effects of pH on the standard apparent reduction potentials of the half reactions involved in the nitrogenase reaction are shown in Table 9.5. The effects of pH on the apparent equilibrium constants of the reactions involved in the nitrogenase reaction as shown in Table 9.6. 

These tables can be used to calculate ArG ° and ATH ° at pH 7 and ionic strengths of 0, 0.10, and 0.25 M or at ionic strength 0.25 M and pHs of 5, 6, 7, 8, and 9 for any reaction for which all the reactants are in these tables. They can also be used to calculate standard apparent reduction potentials. The species data can be used to calculate average bindings of hydrogen ions by reactants. Mathematica programs for carrying out these calculations are provided. 

Tables of Standard Apparent Reduction Potentials of Half Reactions... 

When the pH is specified, each biochemical half reaction makes an independent contribution to the apparent equilibrium constant K for the reaction written in terms of reactants rather than species. The studies of electochemical cells have played an important role in the development of biochemical thermodynamics, as indicated by the outstanding studies by W. Mansfield Clarke (1). The main source of tables of ° values for biochemical half reactions has been those of Segel (2). Although standard apparent reduction potentials ° can be measured for some half reactions of biochemical interest, their direct determination is usually not feasible because of the lack of reversibility of the electrode reactions. However, standard apparent reduction potentials can be calculated from for oxidoreductase reactions. Goldberg and coworkers (3) have compiled and evaluated the experimental determinations of apparent equilibrium constants and standard transformed enthalpies of oxidoreductase reactions, and their tables have made it possible to calculate ° values for about 60 half reactions as functions of pH and ionic strength at 298.15 K (4-8). 

Table 8.1 Standard apparent reduction potentials in volts at 298.15 K, pH 6, and 0.25 M ionic strength...

The tables of standard apparent reduction potentials produced here are different from classical tables (2) of ° in that they can be reproduced at other pHs in the range 5 to 9 and other ionic strengths in the range 0 to 0.25 M. For some half reactions, standard transformed reduction potenials can be calculated at temperatures other than 298.15 K. These values have all been calculated from the species database BasicBiochemDataS (9) that has been calculated from experimental measure-... 

Electron donors and acceptors differ in the efficiency with which they donate or accept electrons. Their ability to transfer electrons is expressed as the standard oxidation-reduction potential (or standard redox potential) denoted by which is a constant for a redox couple dependent upon temperature, pH and the concentration of the oxidized and reduced species. The measurement of the standard redox potential of redox couples has been by three methods a spectrophoto-metric method, a potentiometric method and electron spin resonance. By convention, standard redox potentials (Table 13.2) refer to reactions recorded as oxidant+ electron(s)- reductant. Electrons flow from couples of higher potential to those of lower potential in an attempt to equalize the two potentials, a phenomenon termed the electron motive force which is measured in volts (or millivolts). These data are not absolute values since measurements of free carriers differ from that of bound carriers, e.g. FeSg., exhibits an apparent E ... 

Voltammetric simulations for ET reactions of various fe and X at electrodes of various sizes have shown that, for ET reactions with fe near 0.1 cm/s, the BV theory could predict voltammetric responses visibly deviating from that expected by the MHC model as the electrode radii are smaller than 50 mn, while this occurs as Tq approaches 10 nm for ET reactions with fe around 1.0 cm/s. According to the half-wave-potential difference in the polarization curves predicted by the BV and MHC models, the BV-based voltammetric analysis would give standard rate constants of ca. 0.6 and 0.5fe , respectively, for a reaction of 0.1 cm/s kP and 100 kJ/mol X at electrodes with radii of 20 and 10 nm. For the ET reaction with k° of 1.0 cm/s, apparent standard rate constants of 0.8fe and 0.6fe will be obtained by BV-based analysis at an electrode with radii of 10 and 5 nm. Considering that the EDL effect would result in enhanced apparent ET kinetics for cation reduction or anion oxidation (Table 2.1), the measured polarization curves at nanoelectrodes would be closer to that predicted by the BV formalism without including the EDL effect. For anion reduction or cation oxidation, the EDL effect and the MHC formalism both predict inhibited ET kinetics as compared with the conventional BV model combined with the diffusion-based MT theory. In this case, the measured polarization would significantly deviate from the prediction of conventional voltammetric theory. Therefore, BV-based voltammetric analysis would result in apparent rate constants that are significantly lower than the real fc . 

In their first study [78], Fawcett et al. determined apparent rate constants at the formal redox potentials for the first reduction of Cgo and C70 in benzonitrile and dichloromethane. The rate constants obtained for both fullerenes were rather similar. Higher rates were observed for dichloromethane (Table 7.5). Fawcett et al. [78] pointed out that the rate constants at the formal potential are underestimated with respect to the rate constants at the standard potential, since the ion association, which follows the electron transfer reaction, shifts the equilibrium potential in a positive direction. This effect was only estimated in the first study [78], but in a subsequent report [39] on the kinetics of the C6o/C6o and systems, a... 

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