Actinide ions, reduction potentials

The reduction potentials for the actinide elements ate shown in Figure 5 (12—14,17,20). These ate formal potentials, defined as the measured potentials corrected to unit concentration of the substances entering into the reactions they ate based on the hydrogen-ion-hydrogen couple taken as zero volts no corrections ate made for activity coefficients. The measured potentials were estabhshed by cell, equihbrium, and heat of reaction determinations. The potentials for acid solution were generally measured in 1 Af perchloric acid and for alkaline solution in 1 Af sodium hydroxide. Estimated values ate given in parentheses. 

Fig. 5a. Standard (or formal) reduction potentials of actinium and the actinide ions in acidic (pH 0) and basic (pH 14) aqueous solutions (values are in volts...

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... 

Control of the particle valence/conduction band oxidation/reduction potential is not only achieved through a judicious choice of particle component material band edge redox thermodynamics of a single material are also affected by solution pH, semiconductor doping level and particle size. The relevant properties of the actinide metal are its range of available valence states and, for aqueous systems, the pH dependence of the thermodynamics of inter-valence conversion. Consequently, any study of semiconductor-particle-induced valence control has to be conducted in close consultation with the thermodynamic potential-pH speciation diagrams of both the targeted actinide metal ion system and the semiconductor material. 

The formal reduction potentials of actinide ions in aqueous solution are given in Table 28-6, from which it is clear that the electropositive character of the metals increases with increasing Z and that the stability of the higher oxidation states decreases. A comparison of various actinide ions is given in Table 28-7. It must be noted also that, for comparatively short-lived isotopes decaying by a-emission or spontaneous fission, heating and chemical effects due to the high level of radioactivity occur in both solids and aqueous solu-... 

Fig. 17.2 Standard (or formal) reduction potential diagrams for the actinide ions in (a) acidic (pH 0) and (b) basic (pH 14) solutions. (Values in volts versus standard hydrogen electro. ) Note that the solubility of PuOz(OH)i increases from 1 m KOH to 10 m KOH solution [57]. Thus there is evidence for amphoteric behavior of PuOi (OH) by forming Pu02(OH)1 in strong base.

Note that AG° q has a liquid-phase standard state of 1 moI/L, and we can use a gas-phase standard state of either 1 atm or 1 mol/L, as long as we use the same convention for the oxidized and reduced forms. Often the 1 mol/L standard state is used. There are several sources of uncertainties in the calculations of reduction potentials, and we will comment on them by scanning the published literature on actinide elements, for which the most studied redox systems are actinyl aqua ions, with the exception of one study on Pu(VII)/Pu(VIII) [172], The first comment is that redox potentials are defined with respect to the standard hydrogen electrode corresponding to the following half-equation... 

The oxidation-reduction behavior of plutonium is described by the redox potentials shown in Table I. (For the purposes of this paper, the unstable and environmentally unimportant heptavalent oxidation state will be ignored.) These values are of a high degree of accuracy, but are valid only for the media in which they are measured. In more strongly complexing media, the potentials will change. In weakly complexing media such as 1 M HClOq, all of the couples have potentials very nearly the same as a result, ionic plutonium in such solutions tends to disproportionate. Plutonium is unique in its ability to exist in all four oxidation states simultaneously in the same solution. Its behavior is in contrast to that of uranium, which is commonly present in aqueous media as the uranyl(VI) ion, and the transplutonium actinide elements, which normally occur in solution as trlvalent... 

The standard electrode potentials, E, for such reduction reactions are related to the free energy change for the process by equation 5.3. Since some elements may exist in a number of different oxidation states, it is possible to construct electrode potential diagrams, sometimes called Latimer diagrams, relating the various oxidation states by their redox potentials. Examples are shown in Figure 5.6 for aqueous solutions of some first-row d-block metals and for some actinides in 1 mol dm acid. In cases where the reduction involves oxide or hydroxide ions bound to... 

In view of the position of Cm in the actinide series, numerous experiments have been made to ascertain if Cm has only the +3 state in solution no evidence for a lower state has been found. Concerning the +4 state, the potential of the Cm4+/Cm3+ couple must be greater than that of Am44/ Am3+, which is 2.6 to 2.9 V, so that solutions of Cm4+ must be unstable. When CmF4, prepared by dry fluorination of CmF3, is treated with 15M CsF at 0°, a pale yellow solution is obtained which appears to contain Cm4+ as a fluoro complex. The solution exists for only an hour or so at 10° owing to reduction by the effects of a-radiation its spectrum resembles that of the iso-electronic Am34 ion. 

At a time when little was known about ionization potentials of lanthanide ions as well as about thermochemistry of non-tripositive lanthanide speeies, Johnson (1969a) showed that differences in the third ionization potentials /j of the lanthanides are primarily responsible for many of their apparent oxidation-reduction anomalies. In a subsequent paper (Johnson 1969b) he compared and contrasted the relative stabilities of the di-, tri- and tetrapositive oxidation states of the lanthanides and actinides, pointing out how much less is the change in ionization potential for actinides than lanthanides at the half-filled shell (see fig. 4). He elaborated (Johnson 1974) on the first paper by systematizing the properties of the dipositive lanthanide ions in conjunction with those of the alkaline-earth metal ions. 

All actinides from thorium to californium form tetravalent oxidation states. For the three elements of highest atomic number, however, viz. americium, curium, and berkelium, the hydrated ions are too strongly oxidizing to be stable in aqueous solution [7,10]. Their rates of reduction nevertheless vary widely, in the order Bk + < Am < Cm + < Cf, with Bk" being by far the most resistant species. This is also the order of thermodynamic stability, as indicated by the oxidation potentials of the couples [11]. 

These trends become even more marked in the thermodynamic equilibrium parameters of the two types of reaction (Tables 21.21 and 21.22). The values of AS° are much the same for all oxidations, just as they are for all oxidations. But while the values of AS are very negative in the latter reactions, on account of the formation of strongly solvated ions, they are very positive in the former ones, on account of the disappearance of ions. These very favorable changes of AS° are strongly counteracted, however, by unfavorable changes of AH°. The oxidations are all exothermic, the M oxidations all endothermic. On balance, these conditions create a mixture of large and small, negative and positive, values of AG in both types of reactions. This of course illustrates the intricate oxidation-reduction pattern of the actinide elements, also reflected in their oxidation potentials (Chapter 17). 

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