The Stabilization of Oxidation States, and Reduction Potentials

The stability of oxidation states of the heaviest elements, and the influence of relativistic effects, can be investigated by reduction experiments. For that purpose, knowledge of relative stabilities of oxidation states, redox potentials E°, is of crucial importance E° is needed to decide which reducing or oxidizing agent should be chosen to reach the desired state. 

Excluded here are those techniques relating to tracer-level work (below weighable quantity of sample measurement by radioassay only), for the concern here will be with the determination of bulk properties of these elements. Tracer-scale studies usually reveal directly only one property of the element under investigation, that is, its relative preference for one environment over another, or more simply, its phase distribution. Nevertheless, as each new transuranium element was discovered and was available only in trace quantities, a great deal of chemistry was learned by inference from tracer-scale studies, including the identity of oxidation states, approximate values of oxidation or reduction potentials, the composition and stability of complex ions, and relative volatilities. 

Liquid chemistry experiments. Relative stabilities of oxidation states, i.e. standard redox potentials E°, and influence of relativistic effects on them, can be established by reduction experiments. The E° can then be used to define heats of formation in the liquid phase by electrochemical methods through temperature variation of the standard reduction potential. 

Halide chemistry serves to reemphasize both the uniqueness of carbon as well as the dominance of the inert-pair efect in tin and lead. Tin(II) and lead(II) compounds are more ionic than their analogous compounds with + 4 oxidation states. Standard reduction potentials further document the increased stability of the +2 state in tin and lead. 

A major consideration before the ligand exchange equilibria can be considered with reference to biological systems is the stability of a particular oxidation state in the biological medium. Low-spin complexes undergo rapid one-electron oxidation and reduction. As a biological system operates at a low redox potential, say —0.5 to 0.0 volts, reduced, i.e. low valence, states of the metals are to be expected. The metal complexes, Ru, Os, Rh, Ir, Pd, Pt and Au should be reduced to the metallic state in fact but for the slow speed of this reduction. The metals of Fig. 6 will tend to go to the following redox states ... 

Simple thermodynamic considerations state that the reduction process is favoured (i.e. more positive cu(ii)/cu(p potential values are obtained) if the electron transfer is exothermic (AH° negative) and if the molecular disorder increases (AS° positive). It is therefore evident that the positive potential value for the reduction of azurin (as well as that of the most blue copper proteins) is favoured by the enthalpic factor. This means that the metal-to-ligand interactions inside the first coordination sphere (which favour the stability of the reduced form over the oxidized form) prevail over the metal complex-to-solvent interactions inside the second... 

Perhaps the most fundamental fimctional property of a heme prosthetic group at the active site of a heme protein is the relative stability of the reduced and oxidized states of the heme iron. A number of structural characteristics of the heme binding environment provided by the apo-protein have been identified as contributing to the regulation of this equilibrium and have been reviewed elsewhere 82-84). Although a comprehensive discussion of these factors is not possible in the space available here, they can be summarized briefly. The two most significant influences of the reduction potential of the heme iron appear to be the dielectric constant of the heme environment 81, 83) and the chemical... 

The introduction of -alkyl substituents to the secondary amine donors of the macrocycle results in anodic shifts in both oxidation and reduction potentials of the complexes relative to the parent ligand systems (Table II). The extent of anodic shifts depends on the number of alkyl groups introduced to the ligand (47,55a). That is, -alkylation makes the attainment of the Ni(I) state easier and the Ni(III) state more difficult. The stabilization of Ni(I) species by -alkylation is ascribed to solvation and stereochemical effects (55b, 60). -Ethyl groups have greater inductive effects than -methyl groups and yield less anodic shift in both oxidation and reduction potentials (47). This anodic shift of redox potentials may be attributed to weaker Ni-N interactions in the -alkylated complexes. The weaker Ni-N interaction for the tertiary amine results in the stabilization of antibonding o--orbitals of the Ni(II) complex, which makes it more favorable to add an electron, but less favorable to remove an electron. 

In contrast to the lanthanide 4f transition series, for which the normal oxidation state is +3 in aqueous solution and in solid compounds, the actinide elements up to, and including, americium exhibit oxidation states from +3 to +7 (Table 1), although the common oxidation state of americium and the following elements is +3, as in the lanthanides, apart from nobelium (Z = 102), for which the +2 state appears to be very stable with respect to oxidation in aqueous solution, presumably because of a high ionization potential for the 5/14 No2+ ion. Discussions of the thermodynamic factors responsible for the stability of the tripositive actinide ions with respect to oxidation or reduction are available.1,2... 

The Ncrnst equation was given before (Eq. 10.115), and in this chapter the effect of pH on the reduction potential of the hydrogen ion has been mentioned, but the effect in general should be emphasized. There are several types of reactions in which concentrations of the reactants and products affect the stability of various oxidation states. This can be understood through application of the Nemst equation. The reduction potential of hydrogen will vary with the concentration of the hydrogen ion hence the commonly known fact that many reasonably active metals dissolve in acid but rot in base. 

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