The oxidation of U(IV) ions in the diluted solvent phase, 30% TBP/n-dodecane, has been investigated in the presence of plutonium ions, which can act as catalysts for U(IV) oxidation. The reaction was shown to follow the cycle below, with the first and third stages being rate determining.
U4+ + 2Pu4+ + 2H2O → UO22+ + 2Pu3++4H+
2Pu3+ + HNO3 + 2H+ → 2Pu4+ + HNO2 + H2O
Pu3+ + HNO2 + H+ → Pu4+ + NO + H2O
2NO + HNO3 + H2O ⇔ 3HNO2
The overall reaction stoichiometry is the same as for the oxidation of U(IV) by HNO3 in TBP:
The rate equations of both these rate limiting steps have been determined, with that for the U(IV)-Pu(IV) reaction (5) being given by the equation below, where k1=74.4±6 M-1.2 min-1 at 25.2 °C and the activation energy is 72±11 kJ mol-1 (in 0.5 M HNO3).
The rate of the second slow stage, the Pu(III)-HNO2 reaction, is given by the equation below, where the rate constant is k2=627±28 M-1 min-1 at 25.2 °C and the activation energy is 87.2±1.4 kJmol-1 (in 0.5 M HNO3).
Mechanistically, it was shown that the U(IV)-Pu(IV) reaction may proceed via the interaction of the hydrolysed actinide ions U(OH)22+ and PuOH3+ and the Pu(III)-HNO2 reaction was found to most probably involve oxidation of Pu(III) ions by nitrinium nitrate (NONO3) ions in its rate determining step.
The rapid reduction of NpO22+ ions to NpO2+ by U(IV) ions in an aqueous nitric acid solution has been studied and, in many ways, is similar in character to the same reaction in HClO4. The major difference is that in HNO3 the reaction proceeds via two parallel routes. The first is via the hydrolysed UOH3+ ion, as in the reaction in HClO4 and the second is via the non-hydrolysed U4+ ion. These parallel routes lead to the observed order of reaction with respect to H+ ions being reduced from −1 in HClO4 to −0.7 in HNO3. After accounting for the reaction mechanism the rate equation is described by:
where k8 = 404 min−1, k19 = 275 M−1 min−1 and β = 0.009 M at 10 °C and μ = 2. The activation energy was 66.5 ± 4.9 kJ mol−1. Nitrate ions had no effect on the reaction rate but sulphamic acid increased the observed rate, probably through catalysis by sulphate ions arising from sulphamic acid reaction with HNO2 and hydrolysis.
where respectively: k26 = (9.2±0.2)×10-4 l2.2mol-2.2min-1 at 60° C; k19 = 254±10 min-1 at 26.0° C; k6 = 25.3±1.9 mol1.1 l-1.1min-1 at 19.5°C. The activation energies were found to be: E19 = 62.6±2.6 kJmol-1 and E26 = 87.7±9.8 kJmol-1. Np(V) was generally found to be stable for long periods in nearly all the kinetic experiments and the reduction of Np(V) to Np(IV) could only be studied at elevated temperatures and reactant concentrations.
Possible reaction mechanisms for the reduction of Np(VI) and Pu(IV) have been suggested; proceeding, in both cases, via the hydrolysis product and an intermediate CH3CHN¤O radical. Simple solvent extraction experiments have shown that Np(VI) and Pu(IV) can be reductively stripped from 30% TBP/n-dodecane in the presence of U(VI).
The oxidation of hydroxylamine by nitric acid in the presence of technetium ions at temperatures above ~60°C is an autocatalytic process comprising an induction period and then a catalysed reaction involving HNO2, which has accumulated in the solution. Tc ions have no appreciable effect on the reaction rate, which is governed only by the nitric and nitrous acid oxidation reactions of hydroxylamine, but the presence of Tc ions does extend the initial induction period. The rate of hydroxylamine oxidation by HNO3 in the presence of HNO2, that is, after the induction period, was found to be:
-d[NH3OH+]/dt = k[NH3OH+][HNO2][HNO3]3.5
where k = 120±10 l4.5 mol-4.5 min-1 at T = 80°C, μ = 2 and [H+] ≤ 2 M. Under these conditions, the reaction apparently has a high activation energy of 160-180 kJ mol-1. At low temperatures (20-40°C) hydroxylamine is effectively stable in solutions of HNO3 up to concentrations of ~2 M, whether or not Tc(VII) ions are present. Tc(V) was also observed to form at least one complex on reduction with excess hydroxylamine with an absorption maximum between 467 and 480 nm dependent on the solution acidity.