Reduction potential of half-reaction

So far, except for the iron(III)/iron(II) couple [reaction (6) in Table 14.2], we have considered reduction potentials of half reactions with an overall transfer of an even number of electrons (i.e., 2, 4, 6, etc.). However, in many abiotic multielectron redox processes, particularly if organic compounds are involved, the actual electron transfer frequently occurs by a sequence of one-electron transfer steps (Eberson, 1987). The resulting intermediates formed are often very reactive, and they are not stable under environmental conditions. In our benzoquinone example, BQ is first reduced to the corresponding semiquinone (SQ), which is then reduced to HQ ... 

P 14.4 Calculating Reduction Potentials of Half Reactions at Various Conditions of pH and Solution Composition... 

From the standard reduction potentials of half-reactions listed below, construct the plausible electron transfer system for the oxidation of ethanol to acetaldehyde and water. 

Table 7.1. Standard-State Reduction Potentials of Half-Reactions Involving Important Elements in Soils...

In this section, you learned that you can calculate cell potentials by using tables of half-cell potentials. The half-cell potential for a reduction half-reaction is called a reduction potential. The half-cell potential for an oxidation half-reaction is called an oxidation potential. Standard half-cell potentials are written as reduction potentials. The values of standard reduction potentials for half-reactions are relative to the reduction potential of the standard hydrogen electrode. You used standard reduction potentials to calculate standard cell potentials for galvanic cells. You learned two methods of calculating standard cell potentials. One method is to subtract the standard reduction potential of the anode from the standard reduction potential of the cathode. The other method is to add the standard reduction potential of the cathode and the standard oxidation potential of the anode. In the next section, you will learn about a different type of cell, called an electrolytic cell. 

Would the net cell voltage be different from what we calculated ft better not be, because the chemistry is still the same. Box 14-2 shows that neither E° nor E depends on how we write the reaction. Box 14-3 shows how to derive standard reduction potentials for half-reactions that are the sum of other half-reactions. 

Subtraction of reduction (or of oxidation) potentials of half-reactions having the same number of electrons gives the emf of the corresponding cell reaction, e.g. 

Calculation of standard apparent reduction potentials for half reactions... 

Note that reduction of iodine has the higher reduction potential. This half-reaction will proceed in the forward direction as a reduction. The iron half-reaction will proceed in the reverse direction as an oxidation. Rewrite the half-reactions in the correct direction. l2(s) -> 21 (aq) (reduction half-cell reaction)... 

Whereas the standard reduction potential for half-reaction 7.51 is readily measurable in aqueous solution (see Section 7.2), that for half-reaction 7.50 must be determined by a rather elaborate set of experiments involving Na amalgam electrodes (amalgams, see Box 22.3). 

The accepted convention is to give the potentials of half-reactions as reduction processes. For example ... 

The oxidation-reduction potentials for half reactions such as Fe" —> Fe + e are measured by putting a piece of platinum or other inert metal into a solution containing ferrous and ferric ions in standard concentrations, and combining this half cell with the standard hydrogen half cell. Again, the platinum serves to conduct electrons and to catalyze the equilibrium between ferrous and ferric ions. The electromotive force of the cell... 

The standard reduction potential of a reaction is the sum of the two half-reactions. For example, the standard free energy change associated with the reduction of pymvate to lactate ... 

Eq. (4.154) describes the (oxidation) reaction (4.146a) therefore, and the reverse (reduction) process (4.146b) is expressed by Eq. (4.155) when substance B is replaced by A. Because E A) = Eied( ) it also follows then that E% = Ef. This explains why only reduction standard potentials are listed in the literature, almost termed the electrode potentials of half-reactions (Table 4.7). 

The OCV of a charged battery is the electromotive force (EMF) of a cell, which comes from the difference between the reduction potentials of the reactions in the half-cells. During the discharge phase, the battery converts the heat energy that would be released during the chemical reaction between the cathode and anode terminals. This heat energy is converted into electrical energy. 

The reduction potentials in Fig. 17.2 are Latimer diagrams [3] showing the potentials of half-reactions in which the left-hand species is reduced to the right-hand species with the appropriate number of electrons, and H2G to balance the half-reaction. (Species not found in aqueous solution, but whose... 

Calculating cell potential The tendency of a substance to gain electrons is called its reduction potential. The reduction potential of a half-cell reaction is expressed in volts (V). Chemists measure the reduction potentials of half-cells against the standard hydrogen electrode, which has a reduction potential defined as 0 V at 25°C, 1.0 atm pressure, and IM hydrogen ion concentration. 

Figure 21.2a shows a sample/reference half-cell pair for measurement of the standard reduction potential of the acetaldehyde/ethanol 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 and the Fe /Fe couple both cause electrons to flow from the reference half-cell to the sample half-cell that is, reduction occurs spontaneously in each system, and the reduction potentials of both are thus positive. The standard reduction potential for the Fe /Fe half-cell is much larger than that for the fumarate/ succinate half-cell, with values of + 0.771 V and +0.031 V, respectively. For each half-cell, a half-cell reaction describes the reaction taking place. For the fumarate/succinate half-cell coupled to a H Hg reference half-cell, the reaction occurring is indeed a 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... 

In addition to defined standard conditions and a reference potential, tabulated half-reactions have a defined reference direction. As the double arrow in the previous equation indicates, E ° values for half-reactions refer to electrode equilibria. Just as the value of an equilibrium constant depends on the direction in which the equilibrium reaction is written, the values of S ° depend on whether electrons are reactants or products. For half-reactions, the conventional reference direction is reduction, with electrons always appearing as reactants. Thus, each tabulated E ° value for a half-reaction is a standard reduction potential. 

Since oxidation occurs at the anode and reduction at the cathode, the standard cell potential can be calculated from the standard reduction potentials of the two half-reactions involved in the overall reaction by using the equation ... 

As shown above, you can obtain the standard oxidation potential from a table of standard reduction potentials by reversing the reduction halfreaction, and changing the sign of the relevant potential. The reduction and oxidation half-reactions for the previous example are as follows. 

When two interval scales are used to measure the amount of change in the same property, the proportionality of differences is preserved from one scale to the other. For example. Table 1.4 shows reduction potentials of three electrochemical half-cell reactions measured in volts with reference to the standard hydrogen electrode (SHE, E°) and in millivolts with reference to the standard silver-silver chloride electrode (Ag/AgCl, ). For the SHE potentials the proportion of differences between the intervals +0.54 to +0.80 and +0.34 to +0.80 is... 

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