Oxidation-reduction electrodes equilibria

The complete reaction may be regarded as composed of two oxidation-reduction electrodes, a Ox, a Red, and frOx , b Red, combined together into a cell at equilibrium, the potentials of both electrodes are the same ... 

Oxidation-reduction electrodes. An inert metal (usually Pt, Au, or Hg) is immersed in a solution of two soluble oxidation forms of a substance. Equilibrium is established through electrons, whose concentration in solution is only hypothetical and whose electrochemical potential in solution is expressed in terms of the appropriate combination of the electrochemical potentials of the reduced and oxidized forms, which then correspond to a given energy level of the electrons in solution (cf. page 151). This type of electrode differs from electrodes of the first kind only in that both oxidation states can be present in variable concentrations, while, in electrodes of the first kind, one of the oxidation states is the electrode material (cf. Eqs 3.1.19 and 3.1.21). 

All electrodes depend on oxidation and reduction, but the term oxidation-reduction electrode, or redox electrode, is usually reserved for the case in which a species exists in solution in two oxidation stages. This electrode is denoted M(s) Ox, Red, where M is an inert metal (usually platinum) serving as an electron carrier and making electrical contact with the solution. The half-cell equilibrium can either be simple (e.g., Fe + + e = Fe +) or be affected by other... 

The basic electrode is an oxidation-reduction electrode operating under equilibrium conditions between electrons in a noble metal, hydrogen ions in solution, and dissolved molecular hydrogen. The activity of dissolved hydrogen, an+, is taken as the independent variable and is fixed by maintain-... 

At this stage reference may be made to potential mediators, i.e. substances which undergo reversible oxidation-reduction and reach equilibrium rapidly. If we have a mixture of two ions, say M2+ and M +, which reaches equilibrium slowly with an inert electrode, and a very small quantity of cerium(IV) salt is added, then the reaction ... 

Analytical methods based upon oxidation/reduction reactions include oxidation/reduction titrimetry, potentiometry, coulometry, electrogravimetry and voltammetry. Faradaic oxidation/reduction equilibria are conveniently studied by measuring the potentials of electrochemical cells in which the two half-reactions making up the equilibrium are participants. Electrochemical cells, which are galvanic or electrolytic, reversible or irreversible, consist of two conductors called electrodes, each of which is immersed in an electrolyte solution. In most of the cells, the two electrodes are different and must be separated (by a salt bridge) to avoid direct reaction between the reactants. 

Fig. 8-2. Electron state draisity in a metal electrode and in hydrated redox particles on both sides of an electrode interface in equilibrium with redox electron transfer I>m = state density of electrons in metal electrode Oo Dhbd)=state density of redox electrons in hydrated oxidant (reductant) particle cfcredox) = Fermi level of redox electrons ...

Fig. 11-6. Polarization curves of anodic metal dissolution and of cathodic oxidant reduction at a corroding metallic electrode (mixed electrode) s equilibrium...

There existed oxidation-reduction reactions with the same reaction speed on the sulphide mineral surface in water. One is the self-corrosion of sulphide mineral. Another is the reduction of oxygen. If the equilibrium potential for the anodic reaction and the cathodic reaction are, respectively, E and, and the mineral electrode potential is E, the relationship among them is as follows ... 

Since natural waters are generally in a dynamic rather than an equilibrium condition, even the concept of a single oxidation-reduction potential characteristic of the aqueous system cannot be maintained. At best, measurement can reveal an Eh value applicable to a particular system or systems in partial chemical equilibrium and then only if the systems are electrochemically reversible at the electrode surface at a rate that is rapid compared with the electron drain or supply by way of the measuring electrode. Electrochemical reversibility can be characterized... 

The electron activity (or intensity) at redox equilibrium may be measured by a potentiometer. A pH meter or a millivolt meter may be used for measuring the potential difference between a reference electrode (such as a calomel electrode) and an oxidation-reduction indicator electrode (such as platinum, gold, or a wax-impregnated graphite electrode). 

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. 

Redox titrations involve determination of equilibrium between the enzyme and a redox agent of known redox potential. The method requires a redox agent with redox potential close to the protein of interest, to ensure reversibility. The protein is exposed to different concentrations of the redox agent, and once equilibrium is attained, the half cell potential is measured with electrodes and the oxidation-reduction state of the proteins is measured by some physical technique, usually UV-Vis spectrophotometry. The concentration of the oxidized and reduced forms is determined at isosbestic points, and thus spectral characterization of redox species (ferric enzyme,... 

While the redox titration method is potentiometric, the spectroelectrochemistry method is potentiostatic [99]. In this method, the protein solution is introduced into an optically transparent thin layer electrochemical cell. The potential of the transparent electrode is held constant until the ratio of the oxidized to reduced forms of the protein attains equilibrium, according to the Nemst equation. The oxidation-reduction state of the protein is determined by directly measuring the spectra through the tranparent electrode. In this method, as in the redox titration method, the spectral characterization of redox species is required. A series of potentials are sequentially potentiostated so that different oxidized/reduced ratios are obtained. The data is then adjusted to the Nemst equation in order to calculate the standard redox potential of the proteic species. Errors in redox potentials estimated with this method may be in the order of 3 mV. 

Although the law of mass action is equally valid for oxidation-reduction processes, and therefore conclusions as to the direction of reactions may be drawn from the knowledge of equilibrium constants, traditionally a different approach is used for such processes. This has both historical and practical reasons. As pointed out in the previous sections, in oxidation-reduction processes electrons are transferred from one species to another. This transfer may occur directly, i.e. one ion collides with another and during this the electron is passed on from one ion to the other. It is possible, however, to pass these electrons through electrodes and leads from one ion to the other. A suitable device in which this can be achieved is a galvanic cell, one of which is shown in Fig. 1.14. A galvanic cell consists of two half-cells, each made up of an electrode and an electrolyte. The two electrolytes are connected with a salt bridge and, if... 

Ml OXIDATION-REDUCTION POTENTIALS In the previous section metal electrodes were dealt with, and it was shown that the equilibrium between the metal ion and the metal... 

Approximate Determination of Standard Potentials.—Many studies have been made of oxidation-reduction systems with which, for one reason or another, it is not possible to obtain accurate results this may be due to the difficulty of applying activity corrections, uncertainty as to the exact concentrations of the substances involved, or to the slowness of the establishment of equilibrium with the inert metal of the electrode. It is probable that whenever the difference in the number of electrons between the oxidized and reduced states, i.e., the value of n for the oxidation-reduction system, is relatively large the processes of oxidation and reduction occur in stages, one or more of which may be slow. In that event equilibrium between the system in the solution and the electrode will be established slowly, and the measured potential may be in error. To expedite the attainment of the equilibrium a potential mediator may be emploj cd this is a substance that undergoes reversible oxidation-reduction and rapidly reaches equilibrium with the electrode. 

There is nothing in the foregoing discussion that restricts it to reactions at the cathode or to ions it holds, in fact, for any electrode process, either anodic, i.e., oxidation, or cathodic, i.e., reduction, using the terms oxidation and reduction in their most general sense, in which the concentration of the reactant is decreased by the electrode process, provided the potential-determining equilibrium is attained rapidly. The fundamental equation (10) is applicable, for example, to cases of reversible oxidation of ions, e.g., ferrous to ferric, ferrocyanide to ferricyanide, iodide to iodine, as well as to their reduction, and also to the oxidation and reduction of non-ionized substances, such as hydroquinone and qui-none, respectively, that give definite oxidation-reduction potentials. 

At the electrode equilibrium potential Feq e the cathodic current, ic, and the anodic one, 4, which represent the reduction and oxidation reaction rates at the electrode-solution interphase, respectively, are equal and the net current i = jid - /a is zero the ic = 4 value is called the exchange current, iq. The passage of net current i 0) through the cell causes some changes with respect to equilibrium, and these are generically indicated by the term polarizations . The differenee between the value of the electrode potential under flowing current, Fi g, and that of the equilib-... 

FIGURE 14-3 Change in potential energy for an electrode reaction in which reactants go through an activated complex to products A, for reductant at equilibrium potential B, for oxidant at equilibrium potential C, for oxidant at a potential more negative than the equilibrium potential. 

The exchange current density, which governs the rate of attainment of electrode equilibrium, varies enormously from one potential-determining redox couple to another. It varies not only with the initial concentration but also with the ratio of oxidant to reductant [Equation (14-14)]. In titrations performed at great dilution, the equilibrium near the end point may be reached slowly. Therefore, it may be advantageous to select a method of end-point detection that is not dependent on equilibrium near the end point. 

Reactions that take place consecutive to the electrode process can be studied polarographioally only in those cases in which the electrode process is reversible. In these cases the wave-heights and the wave-shape remain unaffected by the chemical processes. However, the half-wave potentials are shifted relative to the equilibrium oxidation-reduction potential, determined e.g. potentiometrically. Hence, whereas in all above examples, limiting currents were measured to determine the rate constant, it is the shifts of half-wave potentials which are measured here. First- and second-order chemical reactions will be discussed in the following. 

The silver-silver chloride electrode is an example of a metal electrode that participates as a member of a redox couple. The silver-silver chloride electrode consists of a silver wire or rod coated with AgCl(s) that is immersed in a chloride solution of constant activity this sets the half-cell potential. The Ag/AgCl electrode is itself considered a potentiometric electrode, as its phase boundary potential is governed by an oxidation-reduction electron transfer equilibrium reaction that occurs at the surface of the silver ... 

Standard electrode potential data are available for an enormous number of halfreactions. Many have been determined directly from electrochemical measurements. Others have been computed from equilibrium studies of oxidation/reduction systems and from thermochemical data associated with such reactions. Table 18-1 contains standard electrode potential data for several half-reactions that we will be considering in the pages that follow. A more extensive listing is found in Appendix 5. ... 

We can generalize Equation 19-6 by stating that at equilibrium, the electrode potentials for all half-reactions in an oxidation/reduction system are equal. This generalization applies regardless of the number of half-reactions present in the system because interactions among all must take place until the electrode potentials are identical. For example, if we have four oxidation/reduction systems in a solution, interaction among all four takes place until the potentials of all four redox couples are equal. 

Oxidation-reduction equilibrium also implies that the electrode potentials of all redox couples in the system are equal. Because of irreversibili ty, this condition is rare in mixtures of redox couples, especially in mixtures containing organic and nitrogen compounds such as the soil solution. 

Applying models of equilibrium oxidation-reduction, such as Figs. 4.2,4.4, and 4.6, quantitatively to soils requires that the electrode potential be known. From the electrode potential one could then calculate the soil solution concentrations of Fe2+, Mn2+, and NO and the sulfate/sulfide ratio from Eq. 4.20. Ideally, the potential of an inert electrode in the system should equal the electrode potential, because the electrode should take on a potential corresponding to the electron availability. This measurement is called the redox potential. 

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