Thermodynamics, redox reaction

Unlike the reactions that we have already considered, the equilibrium position of a redox reaction is rarely expressed by an equilibrium constant. Since redox reactions involve the transfer of electrons from a reducing agent to an oxidizing agent, it is convenient to consider the thermodynamics of the reaction in terms of the electron. 

The holistic thermodynamic approach based on material (charge, concentration and electron) balances is a firm and valuable tool for a choice of the best a priori conditions of chemical analyses performed in electrolytic systems. Such an approach has been already presented in a series of papers issued in recent years, see [1-4] and references cited therein. In this communication, the approach will be exemplified with electrolytic systems, with special emphasis put on the complex systems where all particular types (acid-base, redox, complexation and precipitation) of chemical equilibria occur in parallel and/or sequentially. All attainable physicochemical knowledge can be involved in calculations and none simplifying assumptions are needed. All analytical prescriptions can be followed. The approach enables all possible (from thermodynamic viewpoint) reactions to be included and all effects resulting from activation barrier(s) and incomplete set of equilibrium data presumed can be tested. The problems involved are presented on some examples of analytical systems considered lately, concerning potentiometric titrations in complex titrand + titrant systems. All calculations were done with use of iterative computer programs MATLAB and DELPHI. 

Redox reactions are particularly instructive. If all thermodynamically allowed reactions in liquid NH3 were kinetically rapid, then no oxidizing agent more powerful than N2 and no reducing agent more powerful than H2 could exist in this solvent. Using data for solutions at 25° ... 

S.3.3 Electrocatalytic Modified Electrodes Often the desired redox reaction at the bare electrode involves slow electron-transfer kinetics and therefore occurs at an appreciable rate only at potentials substantially higher than its thermodynamic redox potential. Such reactions can be catalyzed by attaching to the surface a suitable electron transfer mediator (45,46). Knowledge of homogeneous solution kinetics is often used to select the surface-bound catalyst. The function of the mediator is to facilitate the charge transfer between the analyte and the electrode. In most cases the mediated reaction sequence (e.g., for a reduction process) can be described by... 

What Do We Need to Know Already This chapter extends the thermodynamic discussion presented in Chapter 7. In particular, it builds on the concept of Gibbs free energy (Section 7.12), its relation to maximum nonexpansion work (Section 7.14), and the dependence of the reaction Gibbs free energy on the reaction quotient (Section 9.3). For a review of redox reactions, see Section K. To prepare for the quantitative treatment of electrolysis, review stoichiometry in Section L. 

An irreversible reaction of the intermediate of a redox reaction will greatly facilitate redox catalysis by thermodynamic control. A good example is the reduction of the carbon halogen bond where the irreversible reaction is the cleavage of the carbon halogen bond associated, or concerted, with the first electron transfer -pEe... 

The last chapter in this introductory part covers the basic physical chemistry that is required for using the rest of the book. The main ideas of this chapter relate to basic thermodynamics and kinetics. The thermodynamic conditions determine whether a reaction will occur spontaneously, and if so whether the reaction releases energy and how much of the products are produced compared to the amount of reactants once the system reaches thermodynamic equilibrium. Kinetics, on the other hand, determine how fast a reaction occurs if it is thermodynamically favorable. In the natural environment, we have systems for which reactions would be thermodynamically favorable, but the kinetics are so slow that the system remains in a state of perpetual disequilibrium. A good example of one such system is our atmosphere, as is also covered later in Chapter 7. As part of the presentation of thermodynamics, a section on oxidation-reduction (redox) is included in this chapter. This is meant primarily as preparation for Chapter 16, but it is important to keep this material in mind for the rest of the book as well, since redox reactions are responsible for many of the elemental transitions in biogeochemical cycles. 

Thermodynamically, virtually all metals in the elemental form are unstable with respect to redox reactions in environments where they are exposed to air and water, i.e., virtually all environments where they are used. Those metals least likely to oxidize (corrode) were long ago given the distinguished title "noble metals." Efforts to prevent metals from corroding, and the cost of repairing and replacing metal structures that have done so, runs into the billions of dollars annually. Thus, one characteristic feature of the society s use of metals is that the metals are continuously, albeit slowly, "degrading" to a less useful form from the moment they are put into use. 

First, the simple thermodynamic description of pe (or Eh) and pH are both most directly applicable to the liquid aqueous phase. Redox reactions can and do occur in the gas phase, but the rates of such processes are described by chemical kinetics and not by equilibrium concepts of thermodynamics. For example, the acid-base reaction... 

A detailed discussion of redox reactions must wait until Chapter 19, after we explore the nature of the atom, periodic properties of the elements, and thermodynamics. For now, we focus on only a few types of redox reactions that are common and relatively simple. 

C19-0059. Use standard thermodynamic values to determine whether or not each of the following redox reactions is spontaneous under standard conditions ... 

Applying the common equations for the thermodynamics of reversible cells, it is possible to extract energetic parameters for the adatom redox reaction. This approach requires the measurement of voltammograms at different temperatures. If we consider that the adatom oxidation reaction involves the formation of the hydroxide, we can write the following equation for the overall cell reaction ... 

The above important relationship now allows evaluation of the thermodynamic driving force of a redox reaction in terms of a measurable cell emf. Moreover, it is possible to utilize the relationship between the standard state potential and the standard state free energy to arrive at an expression for the equilibrium constant of a redox reaction in terms of the emf. Thus... 

In redox mediation, to have an effective electron exchange, the thermodynamic redox potentials of the enzyme and the mediator have to be accurately matched. For biocatalytic electrodes, efficient mediators must have redox potentials downhill from the redox potential of the enzyme a 50 mV difference is proposed to be optimal [1, 18]. The tuning of these potentials is a compromise between the need to have a high cell voltage and a high catalytic current. Furthermore, an obvious requirement is that the mediator must be stable in the reduced and oxidized states. Finally, for operation of a membraneless miniaturized biocatalytic fuel cell, the mediators for both the anode and the cathode must be immobilized to prevent power dissipation by solution redox reactions between them. 

In conclusion, the unique properties of Prussian blue and other transition metal hexa-cyanoferrates, which are advantageous over existing materials concerning their analytical applications, should be mentioned. First, metal hexacyanoferrates provide the possibility to develop amperometric sensors for non-electroactive cations. In contrast to common smart materials , the sensitivity and selectivity of metal hexacyanoferrates to such ions is provided by thermodynamic background non-electroactive cations are entrapped in the films for charge compensation upon redox reactions. 

The charge carriers may reduce or oxidize the semiconductor itself leading to decomposition. This poses a serious problem for practical photoelectrochemical devices. Absolute thermodynamic stability can be achieved if the redox potential of oxidative decomposition reaction lies below the valence band and the redox potential of the reductive decomposition reaction lies above the conduction band. In most cases, usually one or both redox potentials lie within the bandgap. Then the stability depends on the competition between thermodynamically possible reactions. When the redox potentials of electrode decomposition reactions are thermodynamically more favored than electrolyte redox reactions, the result is electrode instability, for example, ZnO, Cu20, and CdS in an aqueous electrolyte. 

The reactions that are more favored thermodynamically tend to be also favored kineti-cally. Semiconductor electrodes can be stabilized by using this effect. For this purpose, redox couples in the electrolyte are established with the redox potential more negative than the oxidative decomposition potential, or more positive than reductive decomposition potential in such a manner that the electrolyte redox reaction occurs preferentially compared to the electrode decomposition reaction. 

The equilibrium model, despite its limitations, in many ways provides a useful if occasionally abstract description of the chemical states of natural waters. However, if used to predict the state of redox reactions, especially at low temperature, the model is likely to fail. This shortcoming does not result from any error in formulating the thermodynamic model. Instead, it arises from the fact that redox reactions in natural waters proceed at such slow rates that they commonly remain far from equilibrium. 

Fig. 22.6. Redox potentials (mV) of various half-cell reactions during mixing of fluid from a subsea hydrothermal vent with seawater, as a function of the temperature of the mixture. Since the model is calculated assuming 02(aq) and H2(aq) remain in equilibrium, the potential for electron acceptance by dioxygen is the same as that for donation by dihydrogen. Dotted line shows currently recognized upper temperature limit (121 °C) for microbial life in hydrothermal systems. A redox reaction is favored thermodynamically when the redox potential for the electron-donating half-cell reaction falls below that of the accepting half-reaction.

Reduction potentials are thermodynamic quantities and cannot be used to predict the rate at which a redox reaction will occur. 

The general features discussed so far can explain the complexity of these reactions alone. However, thermodynamic and kinetic couplings between the redox steps, the complex equilibria of the metal ion and/or the proton transfer reactions of the substrate(s) lead to further complications and composite concentration dependencies of the reaction rate. The speciation in these systems is determined by the absolute concentrations and the concentration ratios of the reactants as well as by the pH which is often controlled separately using appropriately selected buffers. Perhaps, the most intriguing task is to identify the active form of the catalyst which can be a minor, undetectable species. When the protolytic and complex-formation reactions are relatively fast, they can be handled as rapidly established pre-equilibria (thermodynamic coupling), but in any other case kinetic coupling between the redox reactions and other steps needs to be considered in the interpretation of the kinetics and mechanism of the autoxidation process. This may require the use of comprehensive evaluation techniques. 

The thermodynamic feasibility of redox reactions at the semiconductor-electrolyte interface can be assessed from thermodynamic considerations. Since typical redox potentials for many redox couples encountered in electrolytes of natural or technical systems often lie between the band potentials of typical semiconductors, many electron transfer reactions are (thermodynamically) feasible (Pichat and Fox, 1988). With the right choice of semiconductor material and pH the redox potential of the cb can be varied from 0.5 to 1.5 V and that of the vb from 1 to more than 3.5 V (see Fig. 10.4). 

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