Thermodynamics and Electrode Potentials

This chapter is devoted to the important relationship between electrode potentials and the changes in Gibbs energy (AO ) for half-reactions and overall reactions. In discussions of the properties of ions in aqueous solution it is frequently more convenient to represent changes in Gibbs energy, quoted with units of k.I mol-1, in terms of electrode potentials, quoted with units of volts (V). The electrochemical series is introduced. The properties of the hydrated electron are described. 

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Stem layer adsorption was involved in the discussion of the effect of ions on f potentials (Section V-6), electrocapillary behavior (Section V-7), and electrode potentials (Section V-8) and enters into the effect of electrolytes on charged monolayers (Section XV-6). More speciflcally, this type of behavior occurs in the adsorption of electrolytes by ionic crystals. A large amount of wotk of this type has been done, partly because of the importance of such effects on the purity of precipitates of analytical interest and partly because of the role of such adsorption in coagulation and other colloid chemical processes. Early studies include those by Weiser [157], by Paneth, Hahn, and Fajans [158], and by Kolthoff and co-workers [159], A recent calorimetric study of proton adsorption by Lyklema and co-workers [160] supports a new thermodynamic analysis of double-layer formation. A recent example of this is found in a study... 

The thermodynamic and electrode-kinetic principles of cathodic protection have been discussed at some length in Section 10.1. It has been shown that, if electrons are supplied to the metal/electrolyte solution interface, the rate of the cathodic reaction is increased whilst the rate of the anodic reaction is decreased. Thus, corrosion is reduced. Concomitantly, the electrode potential of the metal becomes more negative. Corrosion may be prevented entirely if the rate of electron supply is such that the potential of the metal is lowered to the value where it is found that anodic dissolution does not occur. This may not necessarily be the potential at which dissolution is thermodynamically impossible. 

Some other theoretical aspects of ionic solvation have been reviewed in the last few years. The interested reader is referred to them ionic radii and enthalpies of hydration 20>, a phenomenological approach to cation-solvent interactions mainly based on thermodynamic data 21>, relationship between hydration energies and electrode potentials 22>, dynamic structure of solvation shells 23>. Brief reviews, monographs, and surveys on this subject from a more or less different point of view have also been published 24—28) ... 

As we have seen, acidity and basicity are intimately connected with electron transfer. When the electron transfer involves an integral number of electrons it is customary to refer to the process as a redox reaction. This is not the place for a thorough discussion of the thermodynamics of electrochemistry that may be found in any good textbook of physical chemistry. Rather, we shall investigate the applications of electromotive force (emf) of interest to the inorganic chemist. Nevertheless, a very brief review of the conventions and thermodynamics of electrode potentials and half-reactions will be presented. 

The properties of lithium metal were described in Chapter 4, where particular note was made of its high specific capacity and electrode potential. However, because of its highly electropositive nature, it is thermodynamically unstable in contact with a wide variety of reducible materials. In particular, lithium reacts with components of most electrolytes to form a passivating layer. Film formation of this type ensures long shelf life for primary lithium cells, but causes severe problems when the electrode is cycled in a secondary cell. 

ELECTROCHEMICAL CELLS THERMODYNAMIC PROPERTIES AND ELECTRODE POTENTIALS... 

The fundamental causative factor for the dissolution of metals and the mechanism of the process are identical to those pertinent to the establishment of an electrode potential. We may begin consideration of die dissolution process with a discussion of the thermodynamically reversible electrode potential, Eeq, of a metal, M, and proceed to show that the dissolution reaction is a departure from the equilibrium conditions. With this approach we can appreciate the role played by the electrode potential in corrosion and gain an insight into the true nature of the process. 

The chapter will assume a knowledge of basic thermodynamics, including the meaning of the terms equilibrium constant, enthalpy, entropy, free energy and electrode potential, and the relationships between them. 

The text is intended to take the reader through some of the topics covered in the first two years of an undergraduate course in chemistry. It assumes some basic knowledge of topics which should be covered in other courses at this level. These include atomic structure, simple quantum theory, simple thermodynamic relationships and electrode potentials. A knowledge of group theory is not explicitly required to follow the text, but reference is made to the symbols of group theory. The data in the text are based on published sources. However, it should not be assumed that data in the problems are based on actual measurements. Although reported data have been used where possible, some values have been calculated or invented for the purpose of the question. 

Despite the favorable properties of acetonitrile as a solvent, its use for equilibrium acidity measurements has its definite limitations. The pK range that is tolerable is limited at the high end by onset of solvent deprotonation, and at the low end by substrate autodissociation, as has been implicated for HCo(CO)4 [14a] and TpCr(CO)3H [22b]. These limitations can be overcome by the choice of a less polar solvent, e.g. 1,2-dichloroethane (DCE), dichloromethane, or THE. To make reliable, quantitative comparisons of thermodynamic data obtained in different solvents, it is necessary to link the acidity scales and electrode potential references in the different solvents. This has all too often proven to be a far from trivial task. Although, in principle, 1 1 relationship between the acidity scales in different solvents never exists, pK differences between closely related compounds are often almost constant when compared in different solvents. This is because their solvation properties are similar, because of similarities in size and charge distribution. In less... 

Numerous applications of standard electrode potentials have been made in various aspects of electrochemistry and analytical chemistry, as well as in thermodynamics. Some of these applications will be considered here, and others will be mentioned later. Just as standard potentials which cannot be determined directly can be calculated from equilibrium constant and free energy data, so the procedure can be reversed and electrode potentials used for the evaluation, for example, of equilibrium constants which do not permit of direct experimental study. Some of the results are of analjrtical interest, as may be shown by the following illustration. Stannous salts have been employed for the reduction of ferric ions to ferrous ions in acid solution, and it is of interest to know how far this process goes toward completion. Although the solutions undoubtedly contain complex ions, particularly those involving tin, the reaction may be represented, approximately, by... 

It is important to obtain experimental information on the thermodynamics of electrode processes to ascertain the tendency of a particular reaction to occur under a given set of experimental conditions namely temperature, pressure, system com H)sition and electrode potential. Such information is provided by the standard- or formal-electrode potentials for the redox couple under consideration. Appropriate combinations of these potentials enable the thermodynamics of homogeneous redox processes to be determined accurately. However, such quantities often are subject to confusion and misinterpretation. It is, therefore, worthwhile to outline their significance for simple electrochemical reactions. This discussion provides background to the sections on electrochemical kinetics which follow. The evaluation of formal potentials for various types of electrode-reaction mechanisms is dealt with in 12.3.2.2. 

Thermodynamic and over potential region For a terminal voltage smaller than the water decomposition potential Ud (Ud 2 V), no significant electrolysis happens and no current flows between the electrodes. 

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