Solvated electron standard oxidation potential

From measured forward and reverse rate constants of the reaction of a solvated electron with water Baxendale (I) estimated the standard potential for the solvated electron. Use of a more recent rate constant (3) and correction (8) for a certain omitted entropy change yields a value, E° e = + 2.7 volts, for the standard oxidation potential of the solvated electron. To calculate AF0,int for a reaction from a difference of the standard oxidation potentials of the two reactants, E° e — E°, the AF°f must be corrected for the AF0. of about 5 kcal./mole. This correction can be made (8) by taking the effective E° e, E° eff, for a solvated electron to be 2.9 volts. 

The E value reflects the stabilization energy of the negative charge by surrounding solvent molecules. Thus its variation by solvation can be used as one of the solvent parameters. Since reaction 1 is nothing but a one-electron oxidation reaction, E can also be called optical oxidation potential The standard oxidation potential value is often diffrcult to be determined because many redox reactions are not reversible. Therefore, the E value should be a good alternative as a measure of redox reactivity for which the equilibrium redox potential is not known. 

An example of the use of these standard states for working with solvation effects on one-electron oxidation potentials is provided elsewhere.11... 

Studies on the electrochemical behavior of ferrocene encapsulated in the hemi-carcerands 61 and 62, indicated that encapsulation induces substantial changes in the oxidation behavior of the ferrocene subunit [98]. In particular, encapsulated ferrocene exhibits a positive shift of the oxidation potential of c. 120 mV, probably because of the poor solvation of ferrocenium inside the apolar guest cavity. Lower apparent standard rate constants were found for the heterogeneous electron transfer reactions, compared to those found in the uncomplexed ferrocene under identical experimental conditions. This effect may be due to two main contributions (i) the increased effective molecular mass of the electroactive species and (ii) the increased distance of maximum approach of the redox active center to the electrode surface. 

While the above description has qualitative merit at an introductory level, it is important to recognize that the ground state electronic configurations apply to unassociated atoms. Ionization (oxidation) of a metal phase in a solvent produces solvated ions, and the stability of these ions is an important influence on the value of the standard potential. Descriptions of alkali metal electrochemistry that are more informative than descriptions based on ground state electronic configurations and periodicity can be obtained by inspection of simple thermodynamic balance sheets. One scheme for assigning the contributions to the oxidation process in solutions is shown in Fig. 1. 

It has long been realized that excellent correlations exist between 1/2 or and molecular orbital parameters [172,178-181]. Correlations between voltammetric data and experimental observables such as ionization potentials (IP) [172,178,182,183], charge-transfer transition energies [172], and positions of p-bands in ultraviolet (UV) absorption spectra of the hydrocarbons [4,172] have been reported as well. The basis and limitations of such correlations have been examined critically [147,175,184]. For alkyl aromatic hydrocarbons (AAHs), the slope, a, of the correlation line [Eq. (56)], where E° is the standard potential for the reversible one-electron oxidation, is close to unity and has been used to suggest that in MeCN the solvation energies of the hydrocarbon radical cations are constant throughout the series [175]. 

Thermodynamically, dissolution should occur whenever the fiee energy of the solvated metal ion at a givai concentration is lower than the free energy of the atom in the metal, plus the thermodynamic potential of the electrons exchanged during the reaction via the standard hydrogen electrode. Such thermodynamic conditions arc summarized in the series of potential- H phase diagrams that have been extensively collated by Poutbaix [69]. Even in cases where a metal is coveted by a thin protective oxide, deleterious corrosion effects can... 

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