Indicator redox couple, half-reaction

The general half-reaction applicable to an indicator redox couple is... 

FIGURE 20.3 Expenmental apparatus used to measure the standard reduction potential of the indicated redox couples (a) the ethanol/acetaldehyde couple, (b) the fumarate/ succinate couple. Part (a) shows a sample/reference half-cell pair for measurement of the standard reduction potential of the ethanol/acetaldehyde couple. Because electrons flow toward the reference half-cell and away from the sample half-cell, the standard reduction potential is negative, specifically -0.197 V. In contrast, the fumarate/succinate couple (b) accepts electrons from the reference half-cell that is, reduction occurs spontaneously in the system, and the reduction potential is thus positive. For each halfcell, a half-cell reaction describes the reaction taking place. For the fumarate/succinate half-cell coupled to a H /H2 reference half-cell (b), the reaction taking place is indeed the reduction of fumarate. 

Some typical half-cell reactions and their respective standard reduction potentials are listed in Table 21.1. Whenever reactions of this type are tabulated, they are uniformly written as reduction reactions, regardless of what occurs in the given half-cell. The sign of the standard reduction potential indicates which reaction really occurs when the given half-cell is combined with the reference hydrogen half-cell. Redox couples that have large positive reduction potentials... 

Calculate the half reaction reduction potentials of the following redox couples in aqueous solution at 25°C under the conditions indicated using (i) and/or (ii) Eh(W) as starting point (see Tables 14.2 and 14.3). Compare the calculated Eu values with the corresponding (W) values. 

A ubiquitous characteristic of vanadium chemistry is the fact that vanadium and many of its complexes readily enter into redox reactions. Adjustment of pH, concentration, and even temperature have often been employed in order to extend or maintain system integrity of a specific oxidation state. On the other hand, deliberate attempts to use redox properties, particularly in catalytic reactions, have been highly successful. Vanadium redox has also been successfully utilized in development of a redox battery. This battery employs the V(V)/V(IV) and V(III)AT(II) redox couples in 2.5 M sulfuric acid as the positive and negative half-cell electrolytes, respectively. Scheme 12.2 gives a representation of the battery. The vanadium components in both redox cells are prepared from vanadium pentoxide. There are two charge-discharge reactions occurring in the vanadium redox cells, as indicated in Equation 12.1 and Equation 12.2. The thermodynamics of the redox reactions involved have been extensively studied [8],... 

It is very important to understand that this kind of speciation calculation indicates that certain redox reactions can occur in soils, but not that they will occur a chemical reaction that is favored by a large value of log K is not necessarily favored kinetically. This fact is especially applicable to redox reactions because they are often extremely slow, and because reduction and oxidation half-reactions often do not couple well to each other. For example, the coupling of the half-reaction for 02(g) reduction with that for N2(g) oxidation leads to log K = -0.3 for the overall redox reaction ... 

By convention, the SHE is defined as having a standard electrode potential, EjJ, of zero volts (V) at standard conditions (25 °C and 1 atm). The standard electrode potentials of other couples are similarly determined as reduction half-reactions at unit activity versus the SHE. If the EjJ for a given half-reaction is > 0, that couple has the potential to oxidize the SHE. A negative EjJ indicates a couple that can reduce the SHE. Tables of redox half-reactions and the corresponding EjJ values can be found in Stumm and Morgan (1996). Table 3.7 gives EjJ values and related parameters from these sources for a dozen environmentally important redox reactions. 

C. Region containing an important potential parameter called the potential at half-maximum current ( 1/2). The reaction rate is fastest at j/2, which can be understood by noting that the slope of the voltammogram is at its maximum [49]. 1/2 is indicative of the formal potential of the dominating redox couple. As there may be a distribution of redox couples in the biofilm, it is appropriate to refer to the distribution as an apparent redox couple. We note that this j/2 is different from the definition presented in Table 5.2 because that 1/2 was derived for diffusion-based electron transfer such as that of the ferricyanide in Case study 5.2. 

Figure 3. Half-reaction reduction potentials of selected organic redox couples (left side), iron(III)/iron(II) couples (middle), and of some biogeochemically important redox couples (right side). Indicated are standard reduction potentials, EJ(w), at environmentally relevant conditions, i.e., T = 25°C, pH = 7.0, [Cl ] = [HCO3] = lO M, [Br ] = 10 M. Note that the standard free-energy change, AG°(w), for a given redox reaction is obtained from the difference between the EJ(w) values of the corresponding half-reactions (see also example given by Eqs. 3-1 and 3-2) AG°(w) = - n F AEj(w) where n is the number of electrons transferred, and F = 96.48 kJ mol V is the Faraday s constant. Data from Stumm (72), Schwarzenbach et al. (63), and references cited therein (am = amorphous aq = aqueous phen = phenanthroline sal = salicylate s = solid porph = porphyrin).

Oxidizing and reducing power is indicated quantitatively by the redox potential or standard electrode potential, E. Redox potentials are normally expressed as reduction potentials. They are obtained by electrochemical measurements and the values are referred to the H /H2 couple for which E is set equal to zero. Thus increasingly negative potentials indicate increasing ease of oxidation or difficulty of reduction. Thus in a redox reaction the half reaction with the most positive value of E is the reduction half and the half reaction with the least value of E (or most highly negative) becomes the oxidation half. 

---