Oxidation-reduction reactions potentials, standard

Determining the direction of spontaneity from electrode potentials Given standard electrode potentials, decide the direction of spontaneity for an oxidation-reduction reaction under standard conditions. (EXAMPLE 20.7)... 

In addition to simple dissolution, ionic dissociation and solvolysis, two further classes of reaction are of pre-eminent importance in aqueous solution chemistry, namely acid-base reactions (p. 48) and oxidation-reduction reactions. In water, the oxygen atom is in its lowest oxidation state (—2). Standard reduction potentials (p. 435) of oxygen in acid and alkaline solution are listed in Table 14.10- and shown diagramatically in the scheme opposite. It is important to remember that if or OH appear in the electrode half-reaction, then the electrode potential will change markedly with the pH. Thus for the first reaction in Table 14.10 O2 -I-4H+ -I- 4e 2H2O, although E° = 1.229 V,... 

The standard potentials of practically all oxidation and reduction reactions, especially those common in the environment and soil, are known or can easily be determined. Because of the specificity and relative ease of conducting voltammetric measurements, they might seem well suited to soil analysis. There is only one major flaw in the determination of soil constituents by voltammetric analysis and that is that in any soil or soil extract, there is a vast array of different oxidation-reduction reactions possible, and separating them is difficult. Also, it is not possible to start an investigation with the assumption or knowledge that all of the species of interest will be either oxidized or reduced. 

Density, AGf, solubility, dissociation const, and standard potentials of oxidation-reduction reactions for aqueous species... 

The reduction-oxidation potential (typically expressed in volts) of a compound or molecular entity measured with an inert metallic electrode under standard conditions against a standard reference half-cell. Any oxidation-reduction reaction, or redox reaction, can be divided into two half-reactions, one in which a chemical species undergoes oxidation and one in which another chemical species undergoes reduction. In biological systems the standard redox potential is defined at pH 7.0 versus the hydrogen electrode and partial pressure of dihydrogen of 1 bar. 

Practically in every general chemistry textbook, one can find a table presenting the Standard (Reduction) Potentials in aqueous solution at 25 °C, sometimes in two parts, indicating the reaction condition acidic solution and basic solution. In most cases, there is another table titled Standard Chemical Thermodynamic Properties (or Selected Thermodynamic Values). The former table is referred to in a chapter devoted to Electrochemistry (or Oxidation - Reduction Reactions), while a reference to the latter one can be found in a chapter dealing with Chemical Thermodynamics (or Chemical Equilibria). It is seldom indicated that the two types of tables contain redundant information since the standard potential values of a cell reaction ( n) can be calculated from the standard molar free (Gibbs) energy change (AG" for the same reaction with a simple relationship... 

Many half-reactions of interest to biochemists involve protons. As in the definition of AG °, biochemists define the standard state for oxidation-reduction reactions as pH 7 and express reduction potential as E °, the standard reduction potential at pH 7. The standard reduction potentials given in Table 13-7 and used throughout this book are values for E ° and are therefore valid only for systems at neutral pH Each value represents the potential difference when the conjugate redox pair, at 1 m concentrations and pH 7, is connected with the standard (pH 0) hydrogen electrode. Notice in Table 13-7 that when the conjugate pair 2ET/H2 at pH 7 is connected with the standard hydrogen electrode (pH 0), electrons tend to flow from the pH 7 cell to the standard (pH 0) cell the measured E ° for the 2ET/H2 pair is -0.414 V... 

Biological oxidation-reduction reactions can be described in terms of two half-reactions, each with a characteristic standard reduction potential, E °. 

The standard free-eneigy change for an oxidation-reduction reaction is directly proportional to the difference in standard reduction potentials of the two half-cells ... 

Redox pairs Oxidation (loss of electrons) of one compound is always accompanied by reduction (gain of electrons) of a second substance. For example, Figure 6.11 shows the oxidation of NADH to NAD+ accompanied by the reduction of FAD to FADH2. Such oxidation-reduction reactions can be written as the sum of two halfreactions an isolated oxidation reaction and a separate reduction reaction (see Figure 6.11). NAD+ and NADH form a redox pair, as do FAD and FADH2. Redox pairs differ in their tendency to lose electrons. This tendency is a characteristic of a particular redox pair, and can be quantitatively specified by a constant, E (the standard reduction potential), with units in volts. 

Standard half-cell potentials can be used to compute standard cell potentials, standard Gibbs free energy changes, and equilibrium constants for oxidation-reduction reactions. 

When a biochemical half-reaction involves the production or consumption of hydrogen ions, the electrode potential depends on the pH. When reactants are weak acids or bases, the pH dependence may be complicated, but this dependence can be calculated if the pKs of both the oxidized and reduced reactants are known. Standard apparent reduction potentials E ° have been determined for a number of oxidation-reduction reactions of biochemical interest at various pH values, but the E ° values for many more biochemical reactions can be calculated from ArG ° values of reactants from the measured apparent equilibrium constants K. Some biochemical redox reactions can be studied potentiometrically, but often reversibility cannot be obtained. Therefore a great deal of the information on reduction potentials in this chapter has come from measurements of apparent equilibrium constants. 

For biological systems, the standard redox potential for a substance (E0 ) is measured under standard conditions, at pH 7, and is expressed in volts. In an oxidation-reduction reaction, where electron transfer is occurring, the total... 

Because any two oxidation-reduction reactions can be combined to make a cell, the tabulation of standard electrode potentials becomes a very efficient way of calculating cell potentials under standard conditions. As indicated by Eq. (54), if the electrode reactions involve the metals of the cell terminals, the metal-metal potential due to the cell terminals is automatically included in the result. A short table of standard electrode potentials is given in Table 2. 

Ferroin With the introduction of Ce(IV) as an oxidant and the evaluation of the formal potential of the Ce(rV)-Ce(III) couple, the need for indicators with higher electrode potentials became evident. The indicator ferroin, tris(l,10-phenanthroline)-iron(II), was discovered by Walden, Hammett, and Chapman, and its standard potential was evaluated at 1.14 V. Hume and KolthofiF found that the formal potential was 1.06 V in 1 M hydrochloric or sulfuric acid. The color change, however, occurs at about 1.12 V, because the color of the reduced form (orange-red) is so much more intense than that of the oxidized form (pale blue). From Figure 15-1 it can be seen that ferroin should be ideally suited to titrations of Fe(II) and other reductants with Ce(lV), particularly when sulfuric acid is the titration medium. It has the further advantages of undergoing a reversible oxidation-reduction reaction and of being relatively stable even in the presence of oxidant. 

Many reactions that occur in living cells are oxidation-reduction reactions. Appendix IX lists several compounds of biological importance and shows their relative tendencies to gain electrons ai 25°C and pH 7 under standard conditions. The numerical values of Ho reflect the reduction potentials relative to the 2H + 2e" H2 half-reaction which is taken as — 0.414 volt at pH 7. The value for the hydrogen half-reaction at pH 7 was calculated from the arbitrarily assigned value (Ho) of 0.00 volt under true standard-state conditions (1 M H and 1 atm Hs). For those few halfreactions of biological importance that do not involve as a reactant, the Ho and Ho values are essentially identical. 

Most analytical oxidation/reduction reactions are carried out in solutions that have such high ionic strengths that activity coefficients cannot be obtained via the Debye-Hiickel equation (see Equation 10-1, Section lOB-2). Significant errors may result, however, if concentrations are used in the Nernst equation rather than activities. For example, the standard potential for the half-reaction... 

Standard reduction potentials provide a basis for comparison among oxidation-reduction reactions. 

Reflect and Apply Why do the electron-transfer reactions of the cytochromes differ in standard reduction potential, even though all the reactions involve the same oxidation-reduction reaction of iron ... 

Redox potential (Eh). The potential that is generated between an oxidation or reduction halfreaction and the standard hydrogen electrode (SHE) (0.0 V at pH = 0). In soils, it is the potential created by oxidation-reduction reactions that take place on the surface of a platinum electrode measured against a reference electrode minus the Eh of the reference electrode. This is a measure of oxidation-reduction potential of redox active components in the soil (see Chapter 4). 

These equations suggest that in oxidation-reduction reactions, the relationship of chemical reactions can be expressed as standard electrode potentials or equilibrium constants. 

Skill 27.1 Demonstrate an understanding of oxidation/reduction reactions and their relationship to standard reduction potentials. 

Compare the potentials, as measured in this experiment, for the oxidation-reduction reactions which occurred in the three galvanic cells with the standard voltages calculated from standard reduction potentials. Suggest possible reasons for any differences. 

The magnitude of the net cell potential AV° will signify the spontaneity of the oxidation-reduction reaction. However, it does not indicate the rate at which corrosion will occur. As noted before, we apply the superscript 0 to denote that we are considering the Standard Electrode Potentials. Engineers may be required to calculate the potential of a particular half-cell at concentrations and temperatures other than the standard conditions. For this purpose, we shall use the Nernst equation, which allows us to account for non-standard temperatures and solution concentrations. 

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