Mirek Icker, C. Walther*, H. Geckeis
Institute for Nuclear Waste Disposal, Karlsruhe Institute of Technology, PO 3640, D 76021 Karlsruhe, (D)
* Corresponding author: [email protected]
Abstract
Redox potentials were measured in acidic aqueous solutions (-log10[H+]=0.7) containing
different fractions of tri- and tetravalent plutonium. Eh values measured directly by a Pt electrode vs Ag/AgCl reference electrode agree very well with the redox potential calculated from the oxidation state distribution Pu(III)/Pu(IV). By monitoring the solutions over 120 days the kinetics of redox state distribution and dissolution of initially present Pu(IV)-colloids were studied. In solutions of Eh>950mV colloids dissolve and form Pu(VI), whereas at lower Eh the dissolution of colloids leads to formation of Pu(III). These findings corroborate the assumption that colloids are an integral part of the aqueous Pu redox chemistry and that formation and dissolution can be fully understood by means of Eh / pH stability calculations.
Introduction
Hydrolysis-, polymerization-, and redox reactions of tetravalent Pu in acidic aqueous solution take place simultaneously and it is often difficult to investigate one reaction without interference of the others. In solutions containing only one oxidation state of Pu, over time equilibrium of two or more oxidation states may form. In this case, Pu is oxidizing and reducing agent at the same time - so called disproportionation
Connick(1949). The prediction of the equilibrium oxidation state distribution in plutonium solutions was facilitated by use of more than ten disproportionation equations and corresponding equilibrium constants Clark et al.(2006), Katz(1986),
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oxidation state distribution and the acidity of the sample. In particular, those reactions involving formation or breaking of Pu-O bonds proceed quite slowly and it may take weeks or even months to reach a steady state. A considerable number of studies on this topic were performed since the 1950s (reviews in Clark et al.(2006), Newton(2002)) and Pu reaction kinetics are still an active area of research (e.g. Haschke(2007)).
In the present work we show that redox reactions of Pu are fully determined by knowing pH and redox potential of the solution. Redox potentials are measured by Pt-electrodes combined with a Ag/AgCl reference electrode and agree very well with those obtained from the distribution of oxidation states. As one of the main results we find that Pu(IV) colloids can dissolve and form Pu(III) or Pu(VI) respectively, depending on whether the solution is in the stability field of Pu(III) or Pu(VI).
Hydrolysis and redox reactions
The terminology used in the present work is adapted from Walther et al.(2007): The oxidation state of aquatic species is designated Roman numerals, e.g. Pu(III). In the case of simultaneous presence of aquatic species and solid the subscript “aq” refers to the aquo ions, e.g. Pu(IV)aq, whereas tetravalent colloidal Pu particles are designated
Pu(IV)coll. The exact ion is written with Arabic numbers and charges (e.g. Pu4+ or
Pu(OH)y4-y). Due to the strong hydrolysis reactions of (mononuclear) Pu4+ ions starting
already at ~pH 0.3 the total Pu(IV) equilibrium concentration is expressed by
4 4
4 4 y 4 y
aq y sp 1y
y 1 y 1
[Pu(IV)] [Pu ]+ [Pu(OH) − ] (K [OH ] )(1− − β [OH ] )−
= =
′ ′
= +
∑
= +∑
(1)with the conditional solubility constant of amorphous Pu(OH)4, Ksp′ =K (γ ) (γ )osp Pu -1 OH -4. β′1y are the conditional formation constants of Pu(OH)y4-y complexes (see Table 2).
4 4-y 4 1 y Pu OH 1 4 Pu (OH) β (γ )(γ ) [Pu (OH) ] β [Pu ] [OH ] (γ ) + − − + − ′ = = o y y y y y y . (2)
[i] denotes the concentration of species i, γi its activity coefficient taking into account
(log K’w = -13.885 at I=0.25 M Rand et al. (2008)). β°1y are the formation constants at
zero ionic strength, γi are calculated using the SIT formalism Ciavatta(1980), Grenthe et
al.(1992) 2 10 log i i ij j j z D m γ = − +
∑
ε (3)zi is the charge of ion i, εij is the interaction parameter for ion i and an oppositely
charged ion j, and mj (mol/kg~H2O) is the molal concentration of ion j. D is the Debye-
Hückel term at 25°C: D=0.509 I / (1 1.5+ I) with I the molal ionic strength) The equilibria between Pu3+ and Pu4+ and between PuO
2+ and PuO22+ are fast because
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Pu4+ + e- ' Pu3+ , log10K° III/IV= -17.69 ± 0.04 Lemire et al.(2001) (4)
PuO22+ + e-' PuO2+ , log10K° V/VI = -15.82 ± 0.09 Lemire et al.(2001) (5)
In contrast, the conversion of the lower oxidation states (Pu(III) or Pu(IV)) to the higher ones (Pu(V) or Pu(VI)) and vice versa are slow since formation or rupture of Pu-O bonds is involved Katz(1986): Pu4+ + 2 H
2O ' PuO2+ + 4 H+ + e- (log10K°IV/V = -17.45
± 0.17 Lemire et al.(2001)). Note the fourth power dependence on hydrogen ion concentration which explains the increasing stability of the Pu(V) state with decreasing H+ concentration Fuger and Oetting(1976). Finally, if Pu(IV) oxyhydroxide colloids or precipitate are present, then solid Pu(IV) is in equilibrium with Pu(V)(aq). according to
Pu(OH)4 (am) ' PuO2+ + 2 H2O + e- (log10K° = -19.8 ± 0.9 Guillaumont et al.(2003)).
Besides measuring the redox potential by means of ion sensitive electrodes, it can be determined from the activities of oxidized (aOx) and reduced (aRed) species in solution. If
complexation of one of the redox partners takes place, e.g. by hydrolysis of Pu(IV)aq,
the amount of Pu4+ must be calculated from the spectroscopic information on the amount of Pu(IV)aq via equation (1). We used equilibrium constants from Yun et
al.(2007)) and performed ionic strength correction according to eq. (3) (I=0.25m). The redox potential E is obtained from the standard potential E° and the activities of Pu3+ and Pu4+, using Nernst’s equation.
ln = o+ Ox e Red a RT E E z F a (6)
R=8.31447 J/mol K, T is the temperature, ze is the number of electrons involved in the
redox reaction and F =96485 C/mol . Using the definition 10
pe= −log ae− =16.9 E (7)
pe is related to the Pu concentrations via 4 ' 10 3 10 / [Pu ] pe log - log [Pu ] KIII IV + + = (8) / III IV
K` is the conditional equilibrium constant, related to K°III/IV by
4 3
10 / 10 / 10 10
log K'III IV =log KIII IVo −log γPu+ +log γPu+ (9) Analogously, the redox couple Pu(V) / Pu(VI) is related to pe via
2 ' 2 10 10 / 2 [PuO ] pe log - log [PuO ] KV VI + + = (10) Sample preparation
A 242Pu stock solution in 0.5M HCl was electrolyticaly oxidized to Pu(VI) and subsequently reduced to Pu(III) (stock solution SIII: [Putot]= 8.24×10-3 M: [Pu(III)]
=8.05×10-3 M and [Pu(IV)aq] =0.191×10-3 M). A part of the stock solution was oxidized
to Pu(IV) (SIV: [Putot]= 6.22×10-3 M: [Pu(IV)aq] =5.64×10-3 M and [Pu(III)] =8.95×10- 5 M). From these stock solutions, five samples were prepared in 0.25M HCl (Table 1).
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pHC (i.e. –log10[H+]) of the samples were measured by a micro glass combination
electrode (ROSS type) from Thermo Scientific calibrated against ten different HCl solutions between [H+]=1M and [H+]=0.1M. The concentration [Putot] was measured by
ICP-MS. Over a time period of six months the redox potentials of the samples were measured (Pt wire against a Ag/AgCl, Kurt Schwabe, Meinsberg). The measured potential was corrected by ΔV=207mV to obtain Eh (SHE). The oxidation state distribution and the fraction of Pu(IV) colloids was observed by absorption spec- troscopy (Cary 5R) and was evaluated as described in Walther et al.(2007). [Pu(IV)coll]
was obtained from the optical absorption at 630nm. Errors are about 3% for Pu(III), Pu(IV) and Pu(VI), and some 10% for the Pu(IV)coll fraction, since the absorption cross
section depends on age and composition of the colloids Lloyd and Haire (1978).
Table 1. Concentrations, pHC and relative abundance of Pu(III) and Pu(IV) in the freshly prepared samples (1) through (5)
Sample 1 2 3 4 [Pu(III)] / % 8.00 9.10 50.00 80.00 [Pu(IV)] / % 92.00 90.9 50.00 20.00 pHC 0.63 0.62 0.64 0.64 [Putot] / M 3.25×10-4 3.28×10-4 3.43×10-4 3.46×10-4 0 500 1000 1500 2000 2500 3000 0 10 20 30 40 50 60 70 80 90 100 0 21 42 63 83 104 125 Time / h Time / d Sample 1 R el. Ab unda nce / % Putot Pu3+ Pu(IV) PuO22+ Colloids 0 500 1000 1500 2000 2500 3000 0 10 20 30 40 50 60 70 80 90 100 0 21 42 63 83 104 125 Time / h Time / d Sample 5 R e l. Ab un dan ce / % Putot Pu3+ Pu(IV) PuO22+ Colloids
Figure 1. Redox state distributions of (1) and (5) measured by absorption spectroscopy.
Results
The Pu oxidation state distribution in all five samples was measured over a time period of 120 days (Fig. 1). In all cases the fraction of Pu(IV)aq (PuIV(OH)nz+, n=0-3) decreases
while Pu(III) increases. A steady state is reached after some 30 days in samples (1) - (3), kinetics is more slowly in case of samples (4) and (5). Simple exponential decay functions (fD), or growth functions (fI), respectively,
º ( ) 0 ⎛ − /τ⎞ ⎜ ⎟ ⎝ ⎠ = + Δ t D f t c c e ; ( ) 0 ⎛1 − /τ⎞ ⎜ ⎟ ⎝ ⎠ = + Δ − t I f t c c e (11)
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fitted to the spectroscopically obtained oxidation state distributions of Fig. 1 yield time constants of τ=128 h, 128 h, 139 h, 351 h and 555 h, for the decrease of Pu(IV)aq in
samples (1) through (5), respectively. The increase of Pu(III) proceeds on similar time scales. The observed formation of Pu(VI), however, proceeds significantly faster in all cases. In samples (1) and (2) (hexavalent) PuO22+ increases up to a relative abundance
of some 10%, whereas in samples (3) to (5) PuO22+ contributes less than 5%. Pu(V)
cannot be measured at concentrations below 10-5M. Upper limit estimations show that Pu(V) contributions are negligible in all samples. At time zero all samples contain, presumably amorphous, PuIV-oxo-hydroxide colloids at some [Pucoll] = 4×10-5M
(Walther et al.(2009)). Over time the largest fraction of these colloids dissolve.
Discussion
The concentration of Pu4+ is obtained from the spectroscopically measured [Pu(IV)]aq
according to eq. (1) using the parameters of Table 2. Subsequently, the redox potentials calculated from the Pu(III)/Pu(IV) abundances according to eq. (8) (denominated “spectroscopic”, solid symbols) are compared to the redox potentials measured by the Ag/AgCl electrode (denominated Eh-Electrode, open symbols) in Figure 2. Samples (1) through (3) show very good agreement, indicating that at all times the redox couple Pu3+/Pu4+ reflects exactly the redox state of the sample. However, the sample does not reach a steady state until t~30 d. Hence we do not denominate the agreement of measured Eh and spectroscopically obtained Eh an equilibrium, at least not for fresh solutions. For samples (4) and (5) slight deviations within the limits of analytical uncertainties are observed. The systematically lower spectroscopically obtained values might originate from the small concentrations of [Pu(IV)aq]< 3×10-5 M close to the
detection limit of our absorption spectrometer.
Table 2. SIT parameters and equilibrium
constants used in the present work (I=0.25 m)
Assuming simultaneous (fast) equilibration of the Pu(III)/Pu(IV) and the PuO2+/PuO22+
couples, the concentration of PuO2+ (i.e. [Pu(V)]) can be calculated using eqs.(7), (10)
and K’V/VI = -15.32 (from Tab. 2). This approach was validated e.g. in Yun et al.(2007).
The maximum relative abundances of Pu(V) range from 0.25 to 2% of [Putot]
corresponding to [Pu(V)] = 5×10-7 M to 6×10-6 M, which is below the optical detection limit. It is justified to neglect Pu(V) in the total balance of Pu obtained from
SIT parameters i j εij (kg mol-1) H+ Cl- 0.12 OH- Na+ 0.04 PuO2+ Cl- 0.09 PuO22+ Cl- 0.21 Pu3+ Cl- 0.23 Pu4+ Cl- 0.4 Pu(OH)3+ Cl- 0.2 Pu(OH)22+ Cl- 0.1
Conditional constants Zero ionic strength
log10β11’ = 13.03 ( Pu(OH)13+ ) log10β11° = 14.0 (a)
log10β12’ = 25.10 ( Pu(OH)22+ ) log10β12° = 26.8 (a)
log10 K’III/IV = -16.72 log1 0K° III/IV = -17.69(b)
log10 K’V/VI = -15.32 log10 K° V/VI = -15.82 (b) (a) from Yun et al.(2007)
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spectroscopic data, but its presence is nevertheless important for explaining the redox chemistry and the kinetics of the present experiments.
It was shown before by Cho et al.(2005), Yun et al.(2007) that the formation of trans- dioxo species from Pu(IV)aq cannot be explained by disproportionation according to
3Pu4+ + 2 H2O ⇔ 2 Pu3+ + PuO22+ + 4 H+ since one would expect equal time constants
for consumption of Pu(IV)aq and formation of Pu(III) and Pu(VI). The central result of
the present work, however, is the role of the colloidal plutonium: The required ratios of consumption of three parts Pu(IV)aq for two parts of Pu(III) and one part Pu(VI) formed
are not fulfilled. As reported in Cho et al.(2005) decrease of Pu(IV)aq and increase of
Pu(III) are approximately equal under the conditions of samples (1) through (3). Samples (4) and (5) behave differently. The increase of Pu(III) exceeds the decrease of Pu(IV)aq, ruling out an exclusive formation of Pu(III) from Pu(IV)aq. However, in both
cases a large number of colloids were present in the freshly prepared solutions that dissolve over time. In recent publications Pu(IV)-oxy-hydroxide colloids were treated as part of the aqueous Pu redox system in so far as they are in equilibrium with (a) Pu(IV)aq as well as (b) with Pu(V) Neck et al.(2007). The reaction Pu(IV)coll ⇔ Pu(IV)aq
is directly plausible. The equilibrium with Pu(V) may be understood through the existence of hyperstoichiometric PuO2+x reported in Conradson et al.(2005) and further
supported by the observation of small mixed valence Pu(IV)/Pu(V) polymers described in Walther et al.(2009). Whether Pu(III) or Pu(VI) is formed at the expense of the colloidal fraction is determined by the redox potential in the solution: Fig. 3 displays the samples in the Pourbaix diagram for [Putot]=2×10-4M. Samples (1) and (2) are in the
stability field of Pu(VI) in agreement with the formation of Pu(VI) after dissolution of colloids. Samples (3)-(5), at lower Eh, are in the stabiliy field of Pu(III) and consequently dissolution of colloids leads to the formation of Pu(III).
0 500 1000 1500 920 940 960 980 1000 1020 Eh / mV Time / h Sample 1 spectrosc Sample 1 Eh-Electrode Sample 2 spectrosc Sample 2 Eh-Electrode Sample 3 spectrosc Sample 3 Eh-Electrode Sample 4 spectrosc Sample 4 Eh-Electrode Sample 5 spectrosc Sample 5 Eh-Electrode
Figure 2. Redox potentials of samples (1) – (5) measured with a Ag/AgCl electrode
(open symbols) and by evaluation of the redox state distribution (solid symbols
Summary and Conclusions
The present results fit into the picture drawn by various authors over the last 50 years:
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plutonium, disproportionation was violated. Costanzo et al.(1973) observed excess reduction of Pu(V) in rather concentrated solutions which were oversaturated with respect to formation of Pu(IV) hydrous oxide and most likely contained Pu(IV) polymers or colloids. Strong deviations of steady state distributions from the expected equilibria were reported by Romanovski et al.(2000) and Nitsche et al.(1994) over a wide range of acidity and Pu concentration. On the other hand, acceleration of redox reactions by Pu(IV) polymers was suggested by Madic et al.(1984) and Newton and
Rundberg(1983). Newton et al.(1986) assumed that Pu(V) disproportionation cannot proceed at all in the absence of Pu(IV) polymers. The effects of polymers become even more pronounced as the colloidal fraction increases for strongly oversaturated solutions.
Yun et al.(2007) observed decrease of Pu(IV)aq and dissolution of colloids at [Putot]=
1×10-5M, pHC 2, pe 14.0. However, Pu(V) was building up faster than Pu(IV)aq
decreased and the formation of Pu(V) correlated well with the decrease of the colloidal fraction -Δ[Pu(IV)coll]=Δ[Pu(V)]. Introducing this additional data point in Figure 3
(open green square) it is easily seen that the sample discussed by Yun is situated in the (meta)stability field of Pu(V) (dashed green lines).
10 11 12 13 14 15 16 17 18 1 2 3 4 0.6 0.7 0.8 0.9 1.0 pe pH 1 bar O2(g) Eh (V) Pu(VI)aq Pu(III)aq - - - Pu(V)aq (metastable) PuO2(am,hyd) [Pu]tot = 2.10-4M (1) (2) (3) (4) (5) Yun
Figure 3. Pourbaix diagram of [Putot]=2×10-4M, I=0.1 M (stability constants of Guillaumont et al.(2003) and β1n for n=1−3 from Yun et al.(2007)). Open circles represent samples (1)-(5), the open square data from Yun et al.(2007). The dashed lines (metastable field of Pu(V)) refer to the reducton of PuO2+ to Pu3+ or Pu(IV). In this field PuO2(am,hyd) is thermodynamically stable but the reduction of PuO2+ to PuO2(am,hyd) is kinetically hindered, because PuO2(am,hyd) cannot be formed directly; the first step requires the reduction of PuO2+ to aqueous Pu(IV) species
We conclude that colloids are in equilibrium with both, Pu(IV) and Pu(V) and act as an intermediate for formation and breaking of the Pu-O bond. The present observations provide further evidence for the hypothesis made in Cho et al.(2005) that the redox
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reactions of Pu(IV) proceed by a two step meachanism, namely the formation of PuO2+ followed by simultaneous equilibration of the redox couples Pu(V)/Pu(VI) and Pu(III)/Pu(IV). Whether the colloids dissolve and form Pu(III)/Pu(IV) or Pu(V)/Pu(VI) depends on the position within the pH-Eh stability fields. Non-equilibrium samples experience a driving force towards their respective stability fields which determines not only the end point but also the kinetics of the reaction. Hence, reactions might be kinetically hindered due to a low driving force and Pu species can be metastable on rather long time scales.
Acknowledgements
We thank Marcus Altmaier for help with the preparation and electrolysis of the Pu stock solutions and Sebastian Büchner and Oliver Schwindt for technical assistance
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