The oxidation state of actinides is a primary determinant of their behavior in solutions. As an example, Table 2.3 and 2.4 summarize their color and stability in aqueous solutions.

The trivalent oxidation state of f elements, which is so markedly stable in the lanthanides and heavier actinides, is not typical for Pu, Np and U. The ability of the light actinides to exhibit multiple oxidation states leads to their rich and complex redox chemistry. Comparison of the stability of their oxidation states (Table 2.4) which are supported by data on their standard reduction potentials at pH = 0 in different acids (Figs 2.1-2.2), can lead to interesting conclusions. For example, while uranium (III) can reduce water, U(IV) and Np(IV) can be stabilized in aqueous solution only under anaerobic conditions. U(III) exists at -0.631 V vs. saturated hydrogen elec­trode, indicating that U(III) liberates H2 from water. Np(III) (0.155 V) is

Table 2.4 Stability of actinide ions in aqueous solution (usually acidic) Z+ Ion Stability 2+ No Stable Md Stable to water, but readily oxidized 3+ Ac Stable U Oxidizes to u4+, reducing water (evolving h2) Np Stable to water, but easily oxidized by air Pu Stable to water and air, but easily oxidized to pu4+ Am Stable; can be oxidized with difficulty Cm Stable Bk Stable; can be oxidized to bk4+ Cf, Es, Fm, Lr Stable No, Md Stable, but rather easily reduced to 2+ 4+ Th Stable Pa Stable to water, but readily oxidized U Stable to water, but slowly oxidized by air to uo22+ Np Stable to water, but slowly oxidized by air to npo2+ Pu Stable in concentrated acid, e. G. 6 M hno3, but disproportionates to pu3+ and puo22+ at lower acidities Am In solution only in presence of strong complexants Cm In solution only in presence of strong complexants Bk Marginally stable; easily reduced to bk3+ 5+ PaO2+ Stable, hydrolyses readily UO2+ Disproportionates to u4+ and uo22+; most nearly stable at ph 2-4 NpO2+ Stable; disproportionates only at high acidities PuO2+ Tends to disproportionate to pu4+ and puo22+ (ultimate products); most nearly stable at ph=8 AmO2+ Disproportionates in strong acid to am3+ and amo22+; (as 241am rapidly reduced at lower acidities by alpha-autoradiolysis) 6+ UO22+ Stable; difficult to reduce NpO22+ Stable; easy to reduce PuO22+ Stable; easy to reduce; (as 239pu slowly reduced by alpha-autoradiolysis) AmO22+ Easy to reduce; (as 241am isotope reduces fairly rapidly by alpha-autoradiolysis) 7+ NpO4(OH)32- Observed only in alkaline solution PuO4(OH)32- Observed only in alkaline solution; oxidizes water

still a strong reducing agent, but less so than hydrogen. Plutonium Pu(III) is stable at 0.982 V.

The reduction potentials for the four common oxidation states of pluto­nium (III-VI) under acidic conditions are all near 1 V. As a result, all four oxidation states can coexist in aqueous solutions (Newton, 1975). The

	-0.631
U(III)		U(IV)	U(V)	U(VI)	(1)
			0.327
	0.982			0.925	(2)
Pu(III)		Pu(IV)	Pu(V)	Pu(VI)
			1.022
		0.938
	1.137		0.739	0.155
+ ЛІ CJ о CL 2		NpO2+	Np(IV)	Np(III)	(3)

0.447 0.677

HNO2 094 HNO3

2.1 The scheme of standard redox potentials (in volts) for U, Pu, Np and 1 M HNO2/HNO3 (Miles, 1990; Drake, 1990).

equilibrium concentrations of plutonium species existing simultaneously will be determined by Equations 2.1 and 2.2.:

Pu4+ + PuO2+ о Pu3+ + PuO22+ 2.1

Pu4++ 2H2O о Pu3+ + PuO2+ + 4H+ 2.2

The equilibrium constant for Equation 2.2 is dependent on [H+]4, hence the position of this disproportionation equilibrium changes significantly with acidity.

The couples in which only an electron is transferred, e. g. Pu3+/Pu4+ or NpO2+/NpO22+, are electrochemically reversible and the redox reactions are rapid. Redox reactions that involve forming or rupturing of the actinide — oxygen bond, e. g. Np4+/NpO2+ and Pu4+/PuO22+, are not electrochemically reversible and have a slower reaction rate because of the barrier introduced by the subsequent reorganization of the solvent shell and also because some of these are two-electron reductions (Edelstein, 2006).

Disproportionation of neptunium is appreciable only for the Np(V) oxi­dation state, and the reaction is favored by high concentration of acid:

2NpO2+ + 4H+ о Np4+ + NpO22+ + 2H2O 2.3

Table 2.5 summarizes the activation parameters of the An(V) disproportio­nation reaction in perchlorate medium [H+] = 1.0 M at 25oC.

2.04* ^ NpO3+ (6)

2.31*

(8)

(9)

Both the tetravalent and hexavalent cations, having higher effective cationic charge, are more strongly complexed by ligands, thus the dispro­portionation reaction will be accelerated toward completion by an addition of complexing agents. Obviously, the reproportionation reaction will be promoted by lower acidity of solution:

Np4+ + NpO22+ + 2H2O о 2NpO2+ + 4H+ 2.4

Examination of the redox potential diagrams of U, Pu and Np, compared with part of the nitrogen-potential diagram, leads to the conclusion that in

Table 2.5 Apparent second order rate constant and activation parameters for the disproportionation reaction (2.3) for actinyl ions An(V) in perchlorate medium, [H+] = 1M at 25°C (Ahrland, 1986) 2AnO2+ + 4H+ ^ An4+ + AnO22+ + 2 H2O k AG* AH* AS* Ion I(M) na M-1s-1 [kJmol-1] [kJmol-1] [JK-1mol-1) UO2+ 2 1 4 x 102 60 46 -46 NpO2+ 2 2 9 x 10-9 119 72 -159 PUO2+ 1 1 3.6 x 10-3 87 79 -24 na — acid dependence of the rate constant.

HNO3 solution containing NO2-, only U(VI), Pu(IV), Pu(VI) and Np(VI) should be present; trivalent species are not stable (Miles, 1990; Drake, 1990). The oxidation of Np(V) by nitrate ion is favored by high HNO3 and low HNO2 concentrations, and conversely, at high concentrations of HNO2, Np(VI) is rapidly reduced to Np(V) (Drake, 1990). Moulin found the rate of oxidation of Np(V) to be dependent on the [HNO2]/[Np(V)] ratio and the nitrate concentration; the mechanism of oxidation starts with the proton activation of Np(V), followed by oxidation of activated species by nitrate (Moulin, 1978). The rate-limiting step of the overall oxidation of Np(V) is the formation of the activated species, except at very low concentrations of HNO2, when oxidation of the activated species becomes comparatively slow (Siddall, Dukes, 1959).

Obviously, the role of the proton in redox processes is paramount, which is confirmed by the fact that the most dramatic changes in reduction poten­tials are observed with increase or decrease of acidity of aqueous solutions. For example, a change of medium from 1 molar acid to 10 molar base causes a change in redox potential of 2 V and makes possible the oxidation of Np(VI) to Np(VII), as seen in Equation 6 in Scheme 2 (Fig. 2.2). Np(VII) can in fact only be prepared in strongly basic solution; a similar effect is seen for Pu, though higher concentrations of base are required.

Generating unusual oxidation states opens new separation opportunities. For example, the typically trivalent Am, if oxidized to upper oxidation states, can be separated from lanthanides either as Am(VI) in a manner similar to other hexavalent actinides (e. g., PuO22+ or UO22+ by extraction with TBP) or, similarly to poorly extracted Np(V), pentavalent Am should be poorly extracted by most solvents, while Ln will be extracted by many extractants. Numerous candidates, including TRPO, CMPO’s and diamides (malonamides, diglycolamides and even picolinamides) are all potential reagents (Nash, 2009).

The equilibria and variety of actinide species discussed above confirm the necessity of careful control of the conditions and oxidation state of the actinides in separation processes.