4.2.1 Trivalent actinides

In aqueous media, americium exists as the trivalent ion except under strongly oxidising conditions, where the penta- and hexavalent trans-dioxo americyl cations, AmO₂⁺ and AmO₂²⁺, are formed [3]. Thus, a sufficient number of experimental studies could be reviewed and thermodynamic data could be selected for Am(III) oxalate, citrate and edta complexes (Table 3).

A consistent picture emerges concerning the simple complexes. Whereas in the case of oxalate data have been selected for Am(ox)⁺, Am(ox)₂⁻ and Am(ox)₃³⁻, citrate may form Am(cit)(aq) and Am(cit)₂²⁻ complexes, but edta forms only Am(edta)⁻, reflecting the increasing size and denticity of the ligands from bidentate (oxalate) to tridentate (citrate) to hexadentate (edta). Additional ligands may coordinate with Am(III), but the interactions are expected to be weak for steric reasons.

The data also reveal a consistent pattern for protonated complexes. The complexes Am(Hedta)(aq), Am(Hcit)⁺ and Am(Hcit)₂⁻ are important species in acidic solutions. Although complexes with Hox⁻ may be possible, their stabilities are expected to be much smaller than those with ox²⁻. The few experimental data concerning protonated Am(III) oxalate complexes have been considered inconclusive in the organics review.

The selected values for Am(III) oxalate, citrate and edta complexes describe well the respective speciation in acidic and neutral to slightly alkaline solutions. However, in alkaline solutions mixed hydroxo complexes may also form.

In the case of edta, a single study proposed the complex Am(edta)(OH)²⁻. Although the data interpretation and the derived constant are not unreasonable, the organics review did not select a value based on a single study of uncertain accuracy. For citrate, mixed hydroxo complexes have also been proposed, but the data were considered as unreliable by the organics review. No information about mixed hydroxo complexes with oxalate could be found in the literature. However, one may infer from the analogous Am(III)–carbonate system, where no evidence for the formation of ternary hydroxide–carbonate complexes was found [6], that they are always minor species if they form at all.

For isa the complex Am(isa_−3H)⁻ has been proposed to form in alkaline solutions. The subscript −3H indicates the assumption that three protons dissociated from the alcoholic OH groups of isa⁻ to form the Am(III)–isa complex in alkaline solutions. This stoichiometry is ambiguous, and the available experimental data may be more appropriately represented by the following reaction:

No value has been selected for these complexes by the organics review. Nevertheless, the reported data are of value by providing “first estimates” of Am(III) isa complexation in alkaline solutions in the absence of more reliable data.

In the case of Pu(III), only equilibrium constants for the edta complexes Pu(edta)⁻ and Pu(Hedta)(aq) could be selected in the organics review (Table 3). Their numerical values are consistent with the corresponding Am(III)–edta complexes. No data for the Pu(III)–edta system in alkaline solutions could be identified.

The data found for the Pu(III)–citrate system are not considered reliable enough by the organics review to derive selected values. However, the values reported for Pu(cit)(aq) and Pu(Hcit)⁺ may be used as guidelines, together with the selected Am(III)–citrate values, for scoping calculations and data estimation procedures.

Several Pu(III)–oxalate complexes have been proposed in the literature. Approximate values of equilibrium constants for Pu(ox)⁺, Pu(ox)₂⁻ and Pu(ox)₃³⁻ are provided for qualitative modelling in the organics review, but no value could be selected.

No data for the Pu(III)–isa system could be found. The discussion of the Am(III)–isa and Eu(III)–isa systems in the organics review may serve as guidelines for qualitative modelling.

Only one experimental study of the Np(III)–edta system could be identified. However, the reported details about the measurements are too poor in order to base a selection on this single study. For environmental modelling studies, the data selected for Pu(III)–edta and Am(III)–edta might be considered as guidelines for data estimation procedures concerning Np(III)–edta complexes.

There is no evidence in the literature on the formation of Np(III)–oxalate complexes in aqueous solutions. However, Np(III) is oxidised by oxalate [16] and hence, the lack of data on unstable complexes is of no consequence for modelling of environmental systems.

No information about the Np(III)–citrate and the Np(III)–isa systems could be found in the literature. As it is not clear whether citrate and isa oxidise Np(III), as oxalate does, or may form stable complexes, as edta does, this lack of information may represent a real gap in the database.

No experimentally determined thermodynamic data could be identified in the organics review for U(III) oxalates, citrates, edta and isa. However, U(III) is unstable in aqueous solution and is thermodynamically capable of being oxidised to U(IV) by many anions [17]. For example, it was observed [17] that U(III) was almost instantaneously oxidised to U(IV) by ammonium oxalate, while oxalate was reduced to glyoxylic acid (OCHCOOH). Also, when edta is added to an U(III) acetate solution, the strongly complexing ligand edta does not stabilise the trivalent state of uranium, but accelerates its oxidation, which indicates that edta plays the role of an oxidising agent with respect to U(III) [17]. Hence, the lack of thermodynamic data for unstable U(III) complexes with organic ligands does not represent a gap in the database.

4.2.2 Tetravalent actinides

The tetravalent oxidation state of actinides is difficult to explore with respect to aqueous complexation. The very strong hydrolysis of U(IV), Np(IV) and Pu(IV) and the tendency of colloid formation in An(IV) aqueous solutions pose serious obstacles for experimental investigations.

Although colloid formation can be of importance in the environmental behaviour of An(IV) aqueous solutions, a general discussion of the colloidal properties of tetravalent actinides solutions was outside the scope of the NEA organics review. Colloid formation has been considered in the review of experimental data only with respect to its influence on the quality of thermodynamic data.

A few binary aqueous U(IV), Np(IV) and Pu(IV) oxalate complexes are reported in the literature (Table 2). Their stability constants are usually determined in conjunction with solubility measurements for solid An(IV) oxalates, which are associated with various shortcomings as discussed in detail in Ref. [10]. Hence, in spite of the qualitative self-consistency and the apparent reasonableness of the reported stability constants, no thermodynamic data for the complexes shown in Table 2 are recommended by the organics review. Nevertheless, the data may serve as qualitative guidelines in modelling exercises.

The mentioned solubility measurements of An(IV) oxalates have been carried out in acidic solutions in order to minimise problems with the very strong hydrolysis. A ternary complex, U(ox)₂(OH)₂²⁻, was assumed to form when U(ox)₂·6H₂O was dissolved in solutions of ammonium bicarbonate, but no stability constant was determined. Generally, the behaviour of U(IV), Np(IV) and Pu(IV) oxalates in neutral and alkaline solutions is unexplored and the missing stability constants for ternary An(IV)–ox–OH complexes represent serious gaps in the database.

The situation is even worse in the case of U(IV), Np(IV) and Pu(IV) citrates. Either no data at all could be identified, as in the case of Np(IV) citrates, or the few published data for the U(IV) and Pu(IV) citrate systems are ambiguous with respect to the stoichiometry and stability of the formed complexes. Although the limited information for An(IV) citrates suggests very strong complex formation, the ambiguous data cannot be used for scoping calculations.

In the system U(IV)–edta, thermodynamic data for the complex Uedta(aq) could be selected by the organics review (Table 3). The hydrolysis of Uedta(aq) occurs already in acidic medium at pH > 3 with formation of UedtaOH⁻ and of the dimeric species (UedtaOH)₂²⁻. In addition, the formation of Uedta(OH)₂²⁻ has been reported at pH > 7. Although no values could be selected by the organics review for U(IV) edta hydrolysis species, they may serve as qualitative guidelines in modelling exercises.

Furthermore, thermodynamic data for the complex Np(edta)(aq) could be selected (Table 3), which is consistent with the Uedta(aq) data. The Np(IV)–edta–OH system has not been explored experimentally, but data for the analogous U(IV)–edta–OH mixed complexes can be used as qualitative guidelines.

In the case of Pu(IV)–edta, a paper of Boukhalfa et al. [18] became available to the reviewers only in the final stage of preparation of their review. The reported species and their stability constants are consistent with the U(IV) and Np(IV)–edta data and the experimental procedures seem to be reliable, but the results need to be re-evaluated using NEA auxiliary constants and the SIT approach before an eventual data selection can be made.

For the U(IV), Np(IV) and Pu(IV)–isa systems, the few reported studies qualitatively agree that isa forms very strong complexes with these tetravalent actinides in alkaline solutions. However, the proposed number and stoichiometry of formed complexes widely disagree (Table 2). The most detailed study has been published for the Np(IV)–isa system, but due to the uncertainties in oxidation state analysis, the scarcity of data and the uncertainty in defining the reaction stiochiometries, no thermodynamic values have been selected by the organics review. However, the results give an indication of the order of magnitude of Np(IV)–isa complexation and the reported values may be used for scoping calculations. They may also serve as guidelines for estimating U(IV) and Pu(IV)–isa complexation effects.

Tetravalent americium is unstable in non-complexing solutions and is reduced spontaneously to its more stable III oxidation state [19]. Am(IV) complexes can be obtained in the presence of high concentrations of strongly complexing agents such as carbonates [3]. A thermodynamic constant was selected for the complex Am^IV(CO₃)₅⁶⁻ by the NEA americium review [3] based on experimental data of Bourges et al. [20]. In the latter study, Am(IV) was prepared electrochemically by anodic oxidation in sodium bicarbonate–carbonate medium 1.2 M < [HCO₃⁻ + CO₃²⁻] < 2.3 M in the pH range from 9.5 to 10. The Am(IV) solutions were “stable for ∼12 h in the best conditions” [20]. No study about the interaction of Am(IV) with oxalate, citrate, edta or isa could be found in the literature, except that Am(IV) is mentioned as reaction intermediate in the reduction of Am(VI) by oxalic acid [21] (see Section 4.2.4). However, this lack of experimental data probably is of no importance for environmental modelling studies, as we might infer from the Am(IV)–carbonate system that Am(IV) in the presence of oxalate, citrate, edta or isa will be slowly reduced to Am(III).

4.2.3 Pentavalent actinides

Oxalate complexes of U(V) were identified only as reaction intermediates in the reduction of U(VI). No experimental stability constants of aqueous U(V) oxalate complexes were identified in the literature. No information about U(V) citrate, edta or isa complexes could be found. There is no indication that U(V) could be stabilised by complexing ligands, and hence, the lack of thermodynamic data for unstable U(V) complexes with organic ligands does not represent a gap in the database.

Np(V) is the most common redox state of neptunium, reflected by a large body of experimental studies, which allowed us to select thermodynamic constants for Np(V) oxalate, citrate and edta complexes (Table 3).

In the Np(V)–oxalate system, values for the complexes NpO₂ox⁻ and NpO₂(ox)₂³⁻ have been selected. There is no evidence for the formation of higher complexes. A postulated protonated complex, NpO₂Hox(aq), has been rejected by the organics review after re-analysis of a comprehensive experimental data set. On the other hand, although there were clear indications for the formation of mixed Np(V)–OH⁻–ox²⁻ complexes at pH > 9, the sparse experimental data in the alkaline region do not allow to resolve ambiguities with respect to their stoichiometry.

A literature search by the organic review on the thermodynamics of neptunium–citrate systems revealed only information concerning the aqueous complexes of Np(V). Because of the small formal charge of Np(V) and the steric hindrance due to the linear dioxo structure O–Np–O, NpO₂⁺ forms only one-to-one complexes with citrate. A value for the complex NpO₂cit²⁻ has been selected. Although the formation of a weak NpO₂(Hcit)⁻ complex is probable, the thermodynamic data reported are not reliable enough to be accepted. The formation of NpO₂(cit)(OH)³⁻ at pH > 9 has been claimed, the data are considered inconclusive by the organics review.

The complexation of Np(V) with edta has been studied extensively, and the organic review could select thermodynamic values for NpO₂edta³⁻, NpO₂(Hedta)²⁻ and NpO₂(H₂edta)⁻. A single value reported for the species NpO₂edtaOH⁴⁻ has not yet been confirmed by any other study. Although the reported stability constant does not look unreasonable, indicating the formation of this species at pH > 11, the value is not selected by the organics review, but it may serve as qualitative guidelines in modelling exercises.

No information has been found for the Np(V)–isa system. As we may infer from the above discussion that Np(V)–isa complexes might be stable with respect to redox effects, and as there is no chemical analogue for this system, the lack of data represents a gap in the database.

Stability constants have been reported for Pu(V) oxalate and edta complexes (see Table 2). However, Pu(V) disproportionates strongly in oxalate solution [22] and forms Pu(IV) and Pu(VI) oxalate complexes. The Pu(VI) oxalate complex was observed to be slowly reduced to a Pu(IV) oxalate complex without the buildup of a Pu(V) oxalate complex [23]. Furthermore, Pu(VI) is reduced by edta to Pu(V) and finally to Pu(IV) [23]. Hence, no reliable stability constants are available for Pu(V) oxalate and edta complexes, and no values are selected in the organics review for these systems. No data could be found for the Pu(V) citrate and isa systems. However, there is no need to estimate values for highly unstable Pu(V) organic complexes, except perhaps for short-term laboratory studies involving the kinetics of Pu(V) disproportionation and reduction in the presence of organic ligands.

Am(V) in oxalate and edta solutions slowly changes into the trivalent state. The rate of reduction of 1 mM solution of Am(V) at 25 °C is similar in both cases, i.e. 2–3% per hour in 0.05 M edta [24], and about 2% in 0.1 M oxalate [25,26]. Stability constants were proposed in the literature for the complexes AmO₂ox⁻, AmO₂(ox)₂³⁻, and AmO₂(Hedta)²⁻. Qualitatively, AmO₂⁺ forms oxalate and edta complexes with similar stabilities as NpO₂⁺ complexes, but no values were selected in the organics review because of various shortcomings in the experimental procedures. However, considering the instability of Am(V) in the presence of oxalate and edta, their values may only be of importance in short-term laboratory studies. They are of no importance in any long-term environmental modelling study.

No information about Am(V)–citrate and isa systems could be found, but we may infer that citrate and isa have similar reducing effects on Am(V) as oxalate and edta, as higher oxidation states of americium are generally reduced by organic complexing agents [19].

4.2.4 Hexavalent actinides

The U(VI)– oxalate, citrate and edta systems have been studied extensively, and in all cases the NEA organics review could select thermodynamic values for several complexes (Table 3). In the case of oxalate, only data for the mononuclear species UO₂ox(aq), UO₂(ox)₂²⁻ and UO₂(ox)₃⁴⁻ were selected, whereas in the citrate and edta systems, data for dinuclear species also, (UO₂)₂(cit)₂²⁻ and (UO₂)₂edta(aq), and complexes with protonated ligands, UO₂(Hcit)(aq) and UO₂(Hedta)⁻, were selected.

The existence of the protonated U(VI) oxalate complexes that were assumed to have formed in some experiments, UO₂(Hox)⁺, UO₂(Hox)₂(aq) and UO₂(H₂ox)²⁺, is still open for debate. These species, difficult to identify by physical methods, were usually postulated to improve the fitting of potentiometric or spectrophotometric data. If they form, it is most likely that they occur in strongly acidic solutions. Taking into consideration the contradictory information in the literature on the presence of these complexes, stability constants reported in the literature were not accepted by the organics review.

A few dinuclear U(VI) oxalate complexes were assumed to exist in solutions, i.e. (UO₂)₂(ox)₃²⁻ and (UO₂)₂(ox)₅⁶⁻. The inclusion of such species in models to interpret experimental data only slightly improved the overall fit, as they remained minor species (<5%) under the chosen experimental conditions. More studies are needed to confirm the presence of such complexes in solution and to obtain reliable stability constants.

A variety of ternary U(VI) oxalate complexes were assumed to exist in solution. However, thermodynamic data for such complexes are rare. Among the reported ternary complexes, the U(VI)–hydroxide–oxalate complexes are of some importance in predicting the chemical behaviour of U(VI) in pH neutral environments. Using the reported data on the U(VI)–hydroxide–oxalate complexes for scoping calculations, it is revealed that these ternary species could amount to 20–30% of the total U(VI) in the pH regions between 6 and 8. However, the stoichiometry of the formed complexes is ambiguous and further studies are needed to obtain reliable data on this system.

A similar situation is found in the U(VI)–citrate system: although the reported observations suggest the formation of U(VI) complexes with H₋₁cit⁴⁻ (deprotonation of the alcoholic hydroxyl group of citrate) or of mixed hydroxide–citrate complexes in the pH regions between 6 and 8, the available data are insufficient and no conclusions can be drawn.

The linear O-U-O configuration of the uranyl ion allows edta to coordinate only in the equatorial plane. Hence, due to sterical hindrance the complex formally written as UO₂edta²⁻ (Fig. 2) might actually be the ternary complex UO₂(Hedta)OH²⁻, forming above pH 5 from UO₂(Hedta)⁻ by the dissociation of a proton from a coordinated water molecule.

A variety of polynuclear U(VI)–edta species, including ternary U(VI)–hydroxide–edta complexes, have been suggested in the literature. No values for polymeric U(VI) edta other than (UO₂)₂edta(aq) are selected in the organics review, but as long as systems are modelled where edta is in excess with respect to uranium, these complexes are not of importance.

A complexation study of isa with U(VI) in acidic solutions could be interpreted in terms of the complexes UO₂isa⁺, UO₂(isa)₂(aq) and UO₂(isa)₃⁻. Although their reported stability constants appear reasonable, more studies are needed to confirm them. The values may be used for scoping calculations, but they only describe the U(VI)–isa system in acidic solutions. In alkaline solutions, other complexes of the type UO₂(OH)_x(isa_−yH)^1−x−y may dominate the U(VI)–isa system.

The reduction of Np(VI) and Pu(VI) by oxalate, citrate and edta has been investigated in low-ionic-strength media and brines [23]. At low ionic strength, Np(VI) was rapidly reduced to form Np(V) organic complexes, whereas Pu(VI) was predominantly reduced to Pu(IV). The presence of organic complexants also led to the rapid reduction of Np(VI) and Pu(VI) in brines, except in cases where carbonate and hydrolytic complexes predominated.

Hence, it is no surprise that no thermodynamic data have been found in the organic review for Np(VI) citrate, edta and isa complexes, and solely approximate values for Np(VI) oxalate complex formation are reported. However, these values are only of qualitative significance for laboratory studies. For any application modelling the long-term behaviour of neptunium, e.g., for nuclear waste disposal, the unstable Np(VI) oxalate complexes and the lack of experimental data for the Np(VI) citrate, edta and isa systems are of no importance, and there is no need to estimate values for highly unstable Np(VI) organic complexes.

Pu(VI) complexes with oxalate, citrate and edta have been proposed in the literature (see Table 2). However, none of these proposals has been considered reliable and no thermodynamic values are selected in the organics review. Considering the instability of Pu(VI) organic solutions, this non-selection has no consequences for environmental modelling studies.

Powerful oxidants oxidise Am(III) and Am(V) to Am(VI), which forms the linear trans-dioxo americyl cation AmO₂²⁺ [19]. In solutions with excess oxalic acid, Am(VI) was found to be rapidly reduced by oxalate to Am(IV), which disproportionates into Am(III) and Am(V) [21]. Also, the simultaneous formation of Am(V) and Am(III) was observed in solutions of edta, citrate and tartrate, although not all the reactions proceeded rapidly, and the yield of Am(III) depended on the pH [21]. No information is available for Am(VI)–isa interaction, but we may infer that isa has the same reducing effect on Am(VI) as citrate and tartrate. No thermodynamic data are available for these highly unstable Am(VI) complexes.