Relative standard electrode potential

Figure 5.9. Relative standard electrode potential of a Cu/Cu electrode.

The thermodynamics of electrochemical reactions can be understood by considering the standard electrode potential, the potential of a reaction under standard conditions of temperature and pressure where all reactants and products are at unit activity. Table 1 Hsts a variety of standard electrode potentials. The standard potential is expressed relative to the standard hydrogen reference electrode potential in units of volts. A given reaction tends to proceed in the anodic direction, ie, toward the oxidation reaction, if the potential of the reaction is positive with respect to the standard potential. Conversely, a movement of the potential in the negative direction away from the standard potential encourages a cathodic or reduction reaction. 

In the introductory chapter we stated that the formation of chemical compounds with the metal ion in a variety of formal oxidation states is a characteristic of transition metals. We also saw in Chapter 8 how we may quantify the thermodynamic stability of a coordination compound in terms of the stability constant K. It is convenient to be able to assess the relative ease by which a metal is transformed from one oxidation state to another, and you will recall that the standard electrode potential, E , is a convenient measure of this. Remember that the standard free energy change for a reaction, AG , is related both to the equilibrium constant (Eq. 9.1)... 

Knowledge of the value of ij (abs) makes it possible to convert all relative values of electrode potential to an absolute scale. For instance, the standard electrode potentials of the oxygen electrode, the zero charge of mercury, and the hydrated electron, in the absolute scale are equal to -5.67,. 25, and 1.57 V, recpectively. ... 

Then, knowing F H2, it was relatively easy to determine values of electrode potentials for any other couple. With this methodology, they devised the standard electrode potentials ° scale (often called the E nought scale , or the hydrogen scale ). 

Figure 2.1 Simplified schematic plots showing the exponential relationship between the current density i and the potential of the electrode, E. (The latter is represented here as being relative to the standard electrode potential of the couple undergoing electromodification for now, the abscissa ( — ) can be thought of as deviation from equilibrium.) Three examples of electron-transfer rate (/feei) are shown (a) (coincident with the y-axis) representing a very fast rate of electron transfer of 10 A cm" (b) representing an average rate of electron transfer of 10 A cm (c) representing a slow rate of electron transfer of 10 A cm . For each trace, T = 298 K and the reaction was symmetrical , i.e. a = 0.5, as defined later in Section 7.5.

Symbol Cd atomic number 48 atomic weight 112.41 a Group IIB (Group 12) metallic element ionization potential 8.994eV electron configuration [Kr]4di°5s2 valence state +2 standard electrode potential, E° -0.40V. The isotopes and their natural relative abundance are ... 

A model has been considered for Sn2 reactions, based on two interacting states. Relevant bond energies, standard electrode potentials, solvent contribntions (nonequi-librinm polarization), and steric effects are included. Applications of the theory are made to the cross-relation between rate constants of cross- and identity reactions, experimental entropies and energies of activation, the relative rates of Sn2 and ET reactions, and the possible expediting of an outer sphere ET reaction by an incipient SN2-type interaction (Marcus, 1997). 

It has been stressed that the reaction order must be measured at a constant potential difference across the electrode/solution interface at which the reaction occurs. It will be recalled that the condition At > = constant is tantamount to the condition that the potential of the electrode relative to a standard reference electrode is constant (Section 7.5.7.3). Thus, in order to obtain in practice the reaction order of, say, species A, one would measure current densities obtained at a certain potential E referred to a standard electrode potential, in solutions containing various concentrations of A and constant concentrations of all other reactants (Fig. 7.73). The potential E must be chosen sufficiently far from the reversible potential E, so that the exponential law (Section 7.2.3b.2) applies even at the highest concentrations of the given species in deelectronation reactions and at the lowest concentrations in electronation reactions. 

Standard free energy of hydration of the gaseous M ions. d Standard electrode potential relative to NHE. 

Lacking a table of standard electrode potentials, or one that is adequate, what guidelines can be used to identify oxidizing and reducing agents, and to estimate their relative strengths Here are a few. 

It is a relatively simple process to set up a scale of redox potentials in a non-aqueous medium using the standard hydrogen electrode in that medium as the fundamental reference electrode. Thus in liquid ammonia, which is a well studied non-aqueous solvent and for which there exists a considerable amount of thermodynamic information,31 the scale of standard electrode potentials is referred to the standard hydrogen electrode in liquid ammonia (equation 25), which is assigned the value of zero volts, and in which the H+ exists as a solvated species, i.e. NH4+. 

The electrochemical redox potential of several possible decomposition reactions at pH = 0 (relative to the potential of the saturated calomel electrode), which have been estimated from thermodynamic parameters (6,17-21), are shown schematically in Figure A. The band levels are shown for open-circuit conditions. The standard electrode potentials were calculated from the free energies of formation, which are summarized below in Table III. 

The standard electrode potential and its temperature coefficient are found in the literature.36 Kinetic parameter values were measured in-house for HOR,33 ORR,34 OER,35 and COR.12 22 Table 2 gives cell component materials and transport properties. The membrane and electrode proton conductivity in Table 2 are based on the measured membrane and electrode resistance,42,43 which is a strong function of relative humidity (RH). In what follows next, we will describe the... 

Considerable insight into the chemistry of a single element can be had by comparing the standard electrode potentials (and thus the relative free energies) of the various oxidation states of the element. The most convenient means of doing this is the Latimer diagram. 

The Nernst equation relates the activities of the species involved with the electrode potential, E, of the half-reaction and its standard electrode potential, which is the value of the potential relative to the standard hydrogen electrode when the activities of all species are unity. For the generic half-reaction... 

The following is a list of standard electrode potentials of common half-reactions in aqueous solution, that is measured relative to the standard hydrogen electrode at 25°C (298.15 K) with all species at unit activity. Most of these values were taken from Standard potentials in aqueous solution, ed. A. J. Bard, R. Parsons, and J. Jordan, Dekker, New York, 1985, in which values for many other half-reactions may also be found. 

Standard (electrode) potential — (E ) represents the equilibrium potential of an electrode under standard-state conditions, i.e., in solutions with the relative activities of all components being unity and a pressure being 1 atm (ignoring the deviations of fugacity and activity from pressure and concentration, respectively) at temperature T. A pressure of 1 bar = 105 Pa was recommended as the standard value to be used in place of 1 atm = 101,325 Pa (the difference corresponds to 0.34 mV shift of potential). If a component of the gas phase participates in the equilibrium, its partial pressure is taken as... 

The common oxidation states of iron are + 2 and + 3. The relative stability of the two oxidation states in acid aqueous solution is defined by the standard electrode potential of + 0.77 V for the Fe3+/Fe2+ couple.1 This potential is such that the hydrated Fe11 cation is thermodynamically unstable with respect to atmospheric oxidation (equation 1). The oxidation is even more favourable in basic solution (equation 2). It is apparent, therefore, that the chemistry of iron, including its... 

One of the first questions one might ask about forming a metal complex is how strong is the metal ion to ligand binding In other words, what is the equilibrium constant for complex formation A consideration of thermodynamics allows us to quantify this aspect of complex formation and relate it to the electrode potential at which the complex reduces or oxidizes. This will not be the same as the electrode potential of the simple solvated metal ion and will depend on the relative values of the equilibrium constants for forming the oxidized and reduced forms of the complex. The basic thermodynamic equations which are needed here show the relationships between the standard free energy (AG ) of the reaction and the equilibrium constant (K), the heat of reaction, or standard enthalpy (A// ), the standard entropy (AS ) and the standard electrode potential (E for standard reduction of the complex (equations 5.1-5.3). 

Equation (3) can be used to calculate the standard electrode potentials. Calculations based on the Bom-Haber cycle to obtain the relative stabilities of oxidation states are known as Oxidation State Diagrams . These diagrams have been found useful in clarifying inorganic chemistry (69), even though their accuracy is sometimes low. 

The most important property we need to know about an element is the stable oxidation states it can assume, because so many other chemical and physical properties depend upon the oxidation state. The second most important property to know concerns the relative stabilities of these oxidation states that is to say, we need to know the standard electrode potential. As will be clear from the discussions in Section III.2b on this subject, the understanding of these two outstanding characteristics of the elements involves at least a knowledge of heats of sublimation, ionization potentials, ionic and atomic radii, and electronic energy levels. In this paragraph we try to focus on these properties, but we also summarize all the other properties so far predicted for the superheavy elements. 

By convention, the electrode potential of any half-reaction is expressed relative to that of a standard hydrogen electrode (half-reaction 2H+ -p 2e -H2) and is called the standard electrode potential, E . Table 34.1 shows the values of E" for selected half-reactions. With any pair of half-reactions from this series, electrons will flow from that having the lowest electrode potential to that of the highest. " is determined at pH = 0. It is often more appropriate to express standard electrode potentials at pH 7 for biological systems, and the symbol is used in all circumstances, it is important that the pH is clearly stated. 

---