Redox, potentials

The redox potentials and the strain energies at the cobalt(III) and cobalt(II) oxidation states of the most stable conformers of a number of hexaaminecobalt(III/ II) complexes are listed in Table 11.1 (selected data from[231]). The strain energy 

Compound a) Total strain energy [kJmor ]b) AH [kJmol ]c T7( bS [V]d) locale [V] 

The experimentally determined redox potentials are given as solid points while the line corresponds to the calculated potentials. Based on Eq. 11.1, where F is the Faraday constant (F = 96.5 kJ mol1) and n = 1, the slope of the line should be 96.5 kJ mol-1 V 1, if differences in AS are neglected and strain relaxation is the only contribution to the variation in redox potential. 

The mean accuracy of the computed potentials is + 0.08 V, and the reasonable linearity and the fact that the intercepts are identical within the error indicate that the neglected terms (solvation, entropy) are approximately dependent of the strain energies, as one might expect on simple qualitative considerations12311. 

From Fig. 11.1b it emerges that the simple electrochemical model (Eq. 11.3) cannot be correct, and the model of Eq. 11.4 has been proposed1343,3441. 

This is the emf of a cell made up from the standard hydrogen electrode at the left coupled with a redox electrode (see Section.9.4.4). For example the cell  

Both the thermodynamics and kinetics of electron transfer reactions (redox potentials and electron transfer rates) have steric contributions, and molecular mechanics calculations have been used to identity them. A large amount of data have been assembled on Co3+/Co2+ couples, and the majority of the molecular mechanics calculations reported so far have dealt with hexaaminecobalt (III/II) complexes. 

The determination of the structure of the encounter complex (relative orientation of the two reactants) and the ensuing information on the stereoselectivity of the electron transfer is a further possible application of molecular mechanics in this field but this has not yet been evaluated. 

Reduction potentials of hexaaminecobalt(III) complexes span a range of more than 0.9 V, with the lowest potential (-0.63 V) exhibited by rCo(tra J-diammac)]3+/2+ and the highest potential (+0.28 V) found for [Co(tmen)3]3+/2+ (for ligand structures see Table 10.1 below), /ra/is-diammac leads to relatively short metal-ligand bonds and therefore stabilizes the cobalt(III) state, tmen leads to relatively long cobalt-amine 

Redox equilibria between the various oxidation states of selenium have been little studied. This is apparently a consequence of slow reaction rates. For instance, it has been repeatedly demonstrated that the redox potential measured by a platinum electrode is not affected by the ratio of Se(VI) / Se(lV) present in solution, [87RUN/LIN]. The values of the standard electrode potentials of the important redox couples Se(VI) / Se(IV) and Se(lV) / Se(0) are essentially based on only one experimental investigation each. A more detailed discussion than would normally be required will therefore be made of the two investigations. 

Schott, Swift, and Yost [28SCH/SW1] measured the equilibrium reaetion Se(s) + 2l2(cr) + 3H20(I) H,Se03(aq) +4H -t 4P 

The data have been reconsidered by the review and the following observations were made, see also Appendix A. The solid selenium used in the experiments was obtained by reduction of a selenite solution by thiosulphate. The selenium might therefore not be in its standard state. The activity of the specimen is most likely elose enough to the standard state activity, however, since the precipitate was kept at boiling temperature for several hours. A recalculation of the side-reactions with more recent values of the auxiliary equilibrium constants made little difference to the result. The analytical data are not always consistent with the stoichiometry of Reaction (V.24) and the known initial composition of the test solution. The authors also observed this and held oxidation of iodide by initially present oxygen responsible for the discrepancies. Flowever, in some instanees the deviations from the expected concentrations are remarkably large. The deviations do not invalidate the results if equilibrium prevails, which was tested. 

A major uncertainty in log, K° (V.24) arises from the activity coefficient correction. This correction has been applied under the initial assumption, but with more recent data from Harned and Robinson [41HAR/ROB]. An estimate of the activity coefficient was also made with the SIT approach using s(H, ClOj) = 0.12 kg-moP from [92GRE/FUG], and e(H, 1) = 0.18 and s(K, I ) = 0.014 kg-mol calculated from data in [59ROB/STO]. The results of the calculations were  

With the Gibbs energy of formation of I from [92GRE/FUG] the Gibbs energy change of the reaction  

n represents number of electrons either released or absorbed in an electrochemical reaction. If n = 0, it is a simple chemical reaction of the type discussed in previous chapters in which v stands for stoichiometric coefficients and M for specific chemical species that are either reactants or the reaction products. For example, and V2 in Eq. 7.1 would be + 1 and — 1, Mi and M2 would be Fe, and Fe (aq), while n will be equal to — 2. Similarly, in Eq. 7.4 would be 1, 1, — 2, and — 2, while M1-M4 would be Ti203, H2O, Ti02, and 2H, and n = —2. 

ffie electrons will be liberated. Such a reaction is an oxidation reaction because one of the components involved in such a reaction will lose an electron and reach into a higher oxidation state. Examples of reactions with n 0 are those represented by Eqs. 7.1, 7.3 and 7.4. On the other hand, if 0, the reaction is a reduction reaction in which the electrons are absorbed by one of the components and this component will reach into a lower oxidation state. Reaction given by Eq. 7.2 is an example of this reduction. 

If electrons are liberated, they will not remain as free charges. They will be absorbed somewhere else in a complete reaction. Thus in a complete chemical reaction, oxidation and reduction reactions are coupled as in the case of a galvanic cell where electrons are liberated by an anode and are subsequently absorbed by a cathode. Thus, these coupled reactions are redox reactions. 

A good example of such redox reactions is a complete reaction of a mixture of iron (Fe) and hematite (Fe203). When this mixture is mixed with an acidic solution, Fe is ionized by the oxidation reaction given by Eq. 7.1. The liberated electrons are now captured by hematite. The reactions that represent capture of electrons are given by 

In reactions 7.6 and 7.7, the protons are obtained from the dissolution of the acid in the aqueous solution. In highly acidic solutions, reaction 7.6 occurs while reaction 7.7 is more probable in less acidic solutions. Reaction 7.1 and either reaction 7.6 or 7.7 represent a complete redox reaction. 

The standard electrode potential, E, in which the suffix H indicates that the potential is on the H2-H+ scale, is derived as follows. We have 

If one redox couple in a redox reaction is present at a much greater concentration than the other, then the concentration of the reduced and oxidized species in this 

For example, in oxic natural waters the principal oxidant is O2 and in agreement with expectations the pe of such waters is generally poised in the range expected for the O2-H2O couple (Morel and Herring, 1993). Thns for water at pH = 7 in eqnilibrinm with atmospheric Pq2 (= 10 atm), the half reaction is 

A further informative example is the organic matter—CO2 couple, which is the principal reductant in natural systems. Consider a solution in equilibrium with atmospheric CO2 at neutral pH and containing 10p,M CH2O , where CH2O represents average organic C in natural systems, whose composition is similar stoichiometrically to that of carbohydrates. The half reaction is 

Notice that, as in Equation (4.17), pe in Equation (4.18) is sensitive to pH but not to the concentrations of the redox species. The sensitivity to the concentration of the redox species depends on the reaction stoichiometry. For the Ee(OH)3-Fe + couple, for example, the half reaction is 

The slope of50kj mol indicates that the variation of redox potentials of over 

4 V is due to a large extentto strain relaxation. Other possible contributions to redox potentials are electronic effects [189], specific hydrogen bonding [435], ion-pairing 

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The diagram gives regions of existence, i.e. for a particular combination of pH and redox potential it can be predicted whether it is thennodynamically favourable for iron to be inert (stable) (region A), to actively dissolve (region B) or to fonn an oxide layer (region C). 

Rossetti R, Nakahara S and Brus L E 1983 Quantum size effects In the redox potentials, resonance Raman spectra and electronic spectra of CdS crystallites In aqueous solution J. Chem. Phys. 79 1086... 

STANDARD REDOX POTENTIALS OF SOME COMMON METALS... 

A number of redox potentials for ion-ion systems are given in Fable 4.3 here again, state symbols are often omitted. 

Changes in ion concentration and temperature influence redox potentials by affecting the equilibrium... 

The change in the redox potential is given quantitatively by the Nernst equation ... 

The data in Tables 4.2 and 4.3 refer to ions in aqueous acid solution for cations, this means effectively [MlHjO), ]" species. However, we have already seen that the hydrated cations of elements such as aluminium or iron undergo hydrolysis when the pH is increased (p. 46). We may then assume (correctly), that the redox potential of the system... 

When the water ligands around a cation are replaced by other ligands which are more strongly attached, the redox potential can change dramatically, for example for the cobalt(II)-cobalt(III) system we have... 

When either hydrogen ions or hydroxide ions participate in a redox half-reaction, then clearly the redox potential is alTected by change of pH. Manganate(Vir) ions are usually used in well-acidified solution, where (as we shall see in detail later) they oxidise chlorine ions. If the pH is increased to make the solution only mildly acidic (pH = 3-6), the redox potential changes from 1.52 V to about 1.1 V, and chloride is not oxidised. This fact is of practical use in a mixture of iodide and chloride ions in mildly acid solution. manganate(VII) oxidises only iodide addition of acid causes oxidation of chloride to proceed. 

Other important effects of ligand and pH changes on redox potentials will be given under the appropriate element. 

USES OF REDOX POTENTIALS Reaction feasibility predictions... 

Thus under standard conditions chloride ions are not oxidised to chlorine by dichromate(Vr) ions. However, it is necessary to emphasise that changes in the concentration of the dichromate(VI) and chloride ions alters their redox potentials as indicated by the Nernst equation. Hence, when concentrated hydrochloric acid is added to solid potassium dichromate and the mixture warmed, chlorine is liberated. 

We have seen that the energetic feasibility of a reaction can be deduced from redox potential data. It is also possible to deduce the theoretical equilibrium position for a reaction. In Chapter 3 we saw that when AG = 0 the system is at equilibrium. Since AG = — nFE. this means that the potential of the cell must be zero. Consider once again the reaction... 

The following redox potentials are given for the oxidation of manganese(II) to manganese(III) in acid and alkaline solution. 

Discuss the factors which influence the redox potential of a half-reaction, illustrating your answer by as many examples as possible. 

Stannate(II) ions are powerful reducing agents. Since, for tin, the stability of oxidation state -b4 is greater than that of oxidation state -b2, tin(II) always has reducing properties, but these are greater in alkaline conditions than in acid (an example of the effect of pH on the redox potential, p. 101). 

Ozone is very much more reactive than oxygen and is a powerful oxidising agent especially in acid solution (the redox potential varies with conditions but can be as high as + 2.0 V). Some examples are 1. the conversion of black lead(ll) sulphide to white lead(II) sulphate (an example of oxidation by addition of oxygen) ... 

As the above redox potentials indicate, only in the presence of very powerful oxidising agents does hydrogen peroxide behave as a reducing agent. For example ... 

Two important redox potentials for reduction by sulphur dioxide in aqueous solution are ... 

Many of the reactions of halogens can be considered as either oxidation or displacement reactions the redox potentials (Table 11.2) give a clear indication of their relative oxidising power in aqueous solution. Fluorine, chlorine and bromine have the ability to displace hydrogen from hydrocarbons, but in addition each halogen is able to displace other elements which are less electronegative than itself. Thus fluorine can displace all the other halogens from both ionic and covalent compounds, for example... 

The acids are only known in aqueous solution all are oxidising agents the standard redox potentials for the reaction... 

In an aquo-complex, loss of protons from the coordinated water molecules can occur, as with hydrated non-transition metal ions (p. 45). To prevent proton loss by aquo complexes, therefore, acid must usually be added. It is for these conditions that redox potentials in Chapter 4 are usually quoted. Thus, in acid solutions, we have... 

In the presence of excess iodide ions, copper(II) salts produce the white insoluble copper(I) iodide and free iodine, because copper(II) oxidises iodide under these conditions. The redox potential for the half-reaction ... 

Most enones are reduced to anion radicals by organo cuprates. It is likely, that this reaction is connected with the alkylation. Both the formation of anion radicals and of conjugate adducts are not observed, when the redox potential of the enone becomes too negative (H.O. House, 1976). 

The standard redox potentials of inorganic oxidants used in organic synthesis are generally around or above + 1.0 V. Organic substrates do not have such high potentials. The values for the CH4/CH3OH and CjHj/CjHjOH couples are at +0,59 V and 0.52 V, respectively. The oxidation of alcohols and aldehydes corresponds to values around 0.0 V (W.M. 

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