Thermodynamic Modeling of Actinide Complexation ... - Springer Link

Journal of Solution Chemistry, Vol. 28, No. 5, 1999

Thermodynamic Modeling of Actinide Complexation with Acetate and Lactate at High Ionic Strength Robert C. Moore,1 Marian Borkowski,2,3 Michael G. Bronikowski,2 Jianfeng Chen,2 Oleg S. Pokrovsky,2 Yuanxian Xia,2 and Gregory R. Choppin2* Received November 20, 1998 The stability constants of NpO2+, UO22+ Am3+, and Th4+ with acetate and lactate anions has been measured in 0.3-5.0m NaCl media at 25°C by the solvent extraction technique. For the 1:1 Complexation, the values of the stability constants increased in the order: NpO+2 < Am3+ < UO22+ < Th4+, in accordance with the actinide charge density and reflecting the strongly ionic bonding of the complexes. The Pitzer ionic interaction parameters were calculated and used to estimate the thermodynamic stability constants at / - 0. Because our data were collected mainly in the high ionic strength region values of p(1) were estimated from values reported in the literature. For all stability constants the Pitzer model gives an excellent representation of the data using three interaction parameters B(0), B (l) , and C* KEY WORDS: Actinides; acetate; lactate; Complexation; Pitzer parameters.

INTRODUCTION A geochemical model is under development to describe the solubility of actinides in underground water present in the salt layers of the nuclear 1

Sandia National Laboratories, MS 0733, Albuquerque, New Mexico 87185-0733. Department of Chemistry, Florida State University, Tallahassee, Florida 32306. 3 On leave from Institute of Nuclear Chemistry and Technology, 03-195 Warsaw, Poland. T We dedicate this paper to the memory of Kenneth S. Pitzer in recognition of his many invaluable contributions to solution chemistry. 2

521 0095-9782/99/0500-0521$16.00/0 C 1999 Plenum Publishing Corporation

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waste repository of the Waste Isolation Pilot Plant (WIPP) Project. The database for this model includes the standard chemical potentials of hundreds of species that may be present in the WIPP disposal room and the Pitzer parameters characterizing their interactions. Organic ligands used in separation and decontamination processes may be present in the nuclear wastes placed in the WIPP site and could have a significant impact on mobile actinide concentrations. In this work, the complexation constants for two organic ligands that may be present at WIPP, acetate and lactate, were determined in NaCl media and the data modeled using the Pitzer formalism. In the past, the complexation stability constants of actinides with acetate and lactate were measured in perchlorate(1-11) and nitrate(12) media of low ionic strength to eliminate possible competing complexation by the background anions. Since WIPP brine contains high concentration of sodium chloride as the main component, we have used NaCl as our background electolyte and Am 3+ , Th4+, NpO+2, and UO2+2 as representatives of tri-, tetra-, penta-, and hexavalent actinides. The solvent extraction technique using di-(2-ethylhexyI)-phosphoric acid (HDEHP) and dibenzoylmethane (HDBM) as extractant was applied to measurement of the stability constants of the complexes. The basis of the thermodynamic solubility model for the WIPP project in the Pitzer equation for activity coefficients. All Pitzer parameters required to model the brine system at WIPP have been reported by Harvie et al.(13) This database has been extensively augmented to include standard chemical potentials and Pitzer parameters to describe actinide interactions in the brine.(14) For complexation between NpO+2 and acetate, thermodynamic parameters have been reported by Novak et al.(l5) We report in this work Pitzer parameters to describe the first stability constants for UO2+2, Am 3+ , and Th4+ with acetate and NpO+2, UO2+2, and Th4+ with lactate and the second stability constant for Th4+ with acetate and lactate. 2. EXPERIMENTAL 2.1. Reagents All the regents used are of ACS certified grade. Dibenzoylmethane (1,3diphenyl-1, 3-propanedione, HDBM) from Aldrich was purified by sublimation at 45°C. Di-(2-ethylhexyl)-phosphoric acid (HDEHP, SIGMA, >95%) was used without purification. A stock solution of 0.050 M HDBM in nheptane was prepared and used as the organic phase for the extraction of Th4+. Stock solutions of 1.0 X 10 -2 M, 1.28 X l0-5 M, 5.0 X 10 -5 M HDEHP in n-heptane (Fisher, ACS) were prepared and used as the organic phase for extraction of NpO+2, UO2+2 and Am 3+ , respectively. The stock

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solutions of acetate and lactate were prepared by dissolution of sodium acetate (Aldrich, ACS) or lactic acid (Fisher, ACS) in deionized water. Their pcH (negative logarithm of H+ concentration in molar units) were adjusted using either HC1 or NaOH solutions to between 6.0-7.2 for NpO+2, 3.0-4.1 for UO22+, 4.7-5.1 for Am3+ and 2.0-3.0 for Th4+. NaCl (Fisher, ACS) was used to established the desired ionic strength. All aqueous solutions were filtered with a 0.2-um pore size Nalgene disposable filter to minimize the possible adsorption of the radioactive tracers by suspended particles. The radioactive tracers were obtained from Oak Ridge National Laboratory. 237Np and 233U tracers were purified by an ion-exchange technique using Dowex 1 X 4 resin. The radionuclides were electrodeposited on stainless steel discs from the samples of the purified solutions and were checked for radioactive purity by alpha spectrometry using a Silicon Surface Barrier detector connected to a pulse height multichannel analyzer. 233Pa, the decay daughter of 237Np, was removed from the 237Np solution at pH 2.5 by extraction with HDEHP prior to the extraction experiments. The oxidation state of neptunium was verified by measurement of the absorption spectroscopy at 980 nm on a CARY-14 spectrophotometer (OLIS upgraded). The 241Am tracer was prepared in 0.001 M HC1 solution, while the 230Th tracer was prepared in 0.1 M HC1 solution to avoid hydrolysis. The purity of both 241Am and 230Th tracers were checked by y-spectroscopy using a Ge(Li) detector connected to a multichannel analyzer. 2.2. Solvent Extraction The solvent extraction experiments were conducted in 20 mL liquid scintillation vials at a fixed volume ratio of 5.0 mL organic : 5.0 mL aqueous solution. The organic solutions were pre-equilibrated with the aqueous stock solutions. A 10 u-L volume of the radioactive tracer solution were added into each vial and the vials were shaken for 3-6 h at constant temperature (25°C), which had been determined to be sufficient to obtain equilibrium for all the systems. After the vials were centrifuged for 3-5 minutes, 0.500-1.00 mL of duplicate samples were taken from both phases. The y-count rates of 241 Am samples were measured with an ISOFLEX automatic gamma counter over the energy window width of 10 - 70 keV. The a-activity of 237Np, 233 U, and 230Th samples were measured on a Tri-Carb (Packard) Liquid Scintillation Counter in plastic scintillation vials in which an aliquot of the sample was mixed with approximately 10 mL Ecolume (ICN Biomedicals, California) liquid scintillation cocktail. The pH meter reading (denoted as pHr) of all the aqueous samples were measured after the extraction using a Corning Semi-Micro Combination glass electrode equipped with an Accumet 950 pH/ion Meter. The pHr were converted to pcH values using(l5)

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where mNaCl denotes the molal concentration of NaCl. 2.3. Calculation of the Stability Constants Under the experimental conditions used, the complexation reactions in the aqueous phase between the actinide ions (Mn+ denotes the four actinide ions) and the acetate or lactate ligands (L - ) can be generalized as:

The distribution coefficient of the actinide ions between the organic and aqueous phases in the presence of the ligand can be described by:

where D0 is the distribution coefficient measured in the absence of the ligand, [L - ] is the free ligand concentration calculated from the measured pcH after reaching extraction equilibrium, the total ligand concentrations, and the relevant protonation constants.(16) The first stability constant, 3,01, was calculated as a slope of the linear regression between (D0/D — 1) and the free ligand concentration [L-]. The second stability constants, B102, were calculated as the slope of the linear regression between [(D 0 /D) - 1]/[L-] and [L-]. Equal weight was given to each data point. Values with a deviation >3a from the regression curve were excluded in the final calculation of the stability constants. All experiments were performed at least twice and six to eight data points were obtained in each measurement. The reported stability constants were calculated as the average values of the data from these separate experiments. In the experiments, the pcH values measured after extraction sometimes varied slightly. The minor corrections to the D values for this pcH fluctuation were made by the relationship:

where n is the experimental slope of the log Dexp dependence on pcH in the absence of ligand for the certain NaCl ionic strength solution. 3. RESULTS AND DISCUSSION The distribution coefficients of NpO+2, UO22+, Am 3+ and Th4+ were measured between HDEHP in heptane and NaCl solutions of 0.3, 1, 2, 3, 4, and 5 m ionic strength in the absence and presence of acetate or lactate ligands. Typical plots of DJD for the thorium plus acetate and uranyl plus

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lactate as a function of the ligand anion concentration are shown in Fig. 1. For NpO+2, the extraction experiments were conducted in a pcH range 6.0-7.2, at which both acetic acid and lactic acid are completely deprotonated. A lower pcH value was used to minimize hydrolysis for the other actinides. The stability constants determined in this work in NaCl media are listed in Table I. These data were used to calculate parameters for the Pitzer model.(17) Standard state chemical potentials and Pitzer interaction parameters were calculated using a nonlinear least-squares analysis based on an algorithm for calculating the minimum Gibbs free energy change for the solution. The minimization algorithm is described in detail elsewhere.(13,18,19) All data points were given equal weight in the calculations. The standard chemical potentials used in this work are listed in Table II, while the Pitzer interaction parameters calculated are listed in Table III. The data were insufficient to obtain model parameters for Am3+ complexation with lactate and for the second stability constants for UO2+2 complexation with acetate and lactate. Model parameters describing the solubility and deprotonation of acetic acid were taken from the literature(15,21,25) From the lactic acid deprotonation reaction, a standard state chemical potential and Pitzer parameters for the lactate ion have been reported by Mizera et al.(22) These authors had insufficient data to determine parameters required to describe solubility of lactic acid in the NaCl solution and, therefore, set the standard chemical potential

Fig. 1. Plots of D0/D as a function of free ligand concentration for uranium with the lactate ion [I = 4.0 m (NaCl), pcH = 4.0] and for thorium with the acetate ion [I = 3.0 m (NaCl), pcH = 3.1].

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Table I. Stability Constants of NpO2+, UO2+2, Am 3+ , and Th4+ Complexes with Acetate and Lactate Measured in NaCl Media; at 25°C (mol-kg )

NpO2+ log B101

Am 3+ log P101

0.30 1.00 2.00 3.00 4.00 5.00

1.05 ±0.04 1. 13 ±0.05 1.25 ±0.05 1.55 ±0.05 1.70±0.20 1.80 ±0.02

1.73+0.02 1.51 ±0.05 1.44+0.02 1.65 ±0.03 1.83 ±0.08 2.20+0.10

0.30 1.00 2.00 3.00 4.00 5.00

1.78 ±0.03 1.43+0.04 1.48 ±0.05 1.76 ±0.02 1.93 ±0.06 1.95 ±0.04

— — — — — —

Im,

-1

UO22+ log

B101

Acetate 2.60+0.15 2.32±0.10 2.52 ±0.07 2.84 ±0.09 3.09+0.04 3.14±0.04 Lactate 2.60±0.01 2.36+0.02 2.45 ±0.02 2.73+0.04 2.50±0.03 2.64+0.01

Th4+

log B102

log

B101

log B102

—

3.73+0.02 3.85 ±0.02 3.92+0.03 4.26±0.03 4.29 ±0.03 4.51+0.03

7.47 ±0.03 6.56+0.03 6.82±0.03 7.19+0.02 7.30±0.03 7.66+0.03

— — — — — —

3.85+0.03 3.83+0.03 3.81+0.02 3.91 ±0.02 4. 17 ±0.03 4.28+0.02

7.08±0.05 6.97 ±0.07 6.43+0.08 6.62 ±0.05 6.98±0.02 7.23+0.04

_ — 5.12±0.02 5.25+0.10 5.72±0.06

of the fully protonated acid to zero. The standard chemical potentials of the organic ligands and all complexes were calculated with reference to this value. Since our data is mainly in the high ionic strength region, we chose to estimate the B (l) Pitzer parameter and to obtain the standard chemical potential and the B0 and C* parameters from regression analysis of the data. The P(l) parameter has its greatest effect on data modeling in the dilute region, whereas C* is required to model data taken in very high ionic strength media. Recent work has demonstrated that setting B (l) = 0, is likely to can introduce systematic deviations in the concentration dependence of the activity coefficients.(20) Plyasunov et al.(26) have described an alternative procedure for obtaining Pitzer parameters from experimental stability constant data, especially when there is a shortage of data in the dilute region. These workers recommend using the simpler SIT model to obtain a standard state chemical potential for the complex under consideration and setting the B (l) parameter at a nonzero value, based on the charge type of the electrolyte. This further reduces the number of Pitzer parameters that must be obtained by the regression analysis of the data. Since data collected in this work include a single data point (at I = 0.30 m) for the dilute region below 1.0 m for each data set, extrapolation to zero ionic strength is questionable using either the SIT or the Pitzer model. For such an extrapolation to be accurate, much more data is required in the dilute region. In this study, the main goal was to accurately model the data over the concentration range where it was collected, i.e., 0.3-5.0 m NaCl. Therefore, average B (l) values were calculated for the different electrolyte

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Actinide Complexation Table II.

Standard Chemical Potentials for Actinide-Organic Complexation

Species

H2O Na+ ClH+ OHAm3+ NpO2+ UO2+2 Th4+ C2O2H4 (acetic acid) C2O2H3- (acetate ion) C3O3H6 (lactic acid) C3O3H5- (lactate ion) Am — acetate2+ Th — acetate3+ Th— lactate3+ NpO2-lactate0 UO2-acetate+ UO2-lactate+ Th-Oactate)22+ Th-(acetate)22+

u 0 /RT a

Ref.

-95.6635 -I05.651 -52.955

13 13 13 13 13 18 20 21 19 15 15 22 22 pwc pw pw pw pw pw pw pw

0 -63.435 -241.694 -369.1 -384.259 -284.23 -158.3 -147.347

0b

8.798 -395.2 -447.1 -291.2 -364.8 -538.6 -382.6 -292.4 -604.8

"Standard chemical potential (dimensionless). b Set to zero. c pw, present work.

charge using data compiled by Pitzer. For 1:1 electrolytes, a value of 0.29 was used for B (l) , which was based on an average of the values reported for salts of carboxylic acids. Similarly for 2:1 electrolytes, a value of 1.74 was used and one of 5.22 was used based on an average for 3:1 type electrolytes. All other parameters were calculated based on these approximations. This strategy of using "average" values of B(1) has been used in this laboratory to model protonation data of lactic acid, oxalic acid, citric acid, and EDTA measured in 0.3-5.0 m NaCl solutions with excellent results.(22) Using the data from the present study, equilibrium constants were calculated using both the SIT and Pitzer approaches. Table IV gives the results along with values reported in the literature. For the NpO2Lac° complex, no data were found at an ionic strength of zero (a value of 1.75 for NpO2Lac° was available at an ionic strength of 0.1). Since extrapolation to zero ionic strength would yield a higher value, the value of 1.97, determined using the Pitzer equation, is probably acceptable. The extrapolated log B0101 values are listed in the literature without error limits. However, the error is probably larger than the error of the corresponding experimental log B101 values mea-

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Table III(A). Pitzer parameters of the present work and reported in the literature, a. Binary Pitzer Parametersa Species i 3+

Am NpO2+

2+

UO

Th4+ Na+ Na+ Na+ Na+ AmAc2+ ThAc3+ ThLac2+2 ThLac3+ ThLac22+ UO2Ac+ UO2Lac+

B0i,j

Species j

2

cl-clcl-

0.6120 0.1415 0.4274 1.092 0 0.1426 0 -0.0563 1.014 0.554 0.4671 0.668 0.5058 0.0124 -0.042

cl

HAc AcHLac LacClCl-clcl -

clclcl

cp,p

B2i,j 5.403 0.281 1.644 13.7 0 0.22 0 0.29 1.74 5.22 1.740 5.22 1.740 0.29 0.29

0 0 0 -160 0 0 0 0 0 0 0 0 0 0 0

-0.028 0 -0.0184 -0.122 0 -0.00629 0 0.047 -0.265 0.371 0.143 0.341 0.225 0.007 0.091

Ref. 18 23 21 24 15 21,25 pwb pw pw pw pw pw pw pw pw

a

Parameters for the brine system: Na+-K+-Mg2+-Ca2+-H+-Cr-SO2-4--OH--HCO3-CO2-3-CO2-H2O are from Harvie et al.(13) b Values set to zero. Table III(B). Neutral-Ion, Like-Ion and Ternary Pitzer Parameters Species i

Species j

AcNpO2Ac0 NpO2Lac°

cl-clcl

Ti,j;Li,j -0.09 0 0.015

Pi,j,Na+

Ref.

0.01029 0 —

15 15 pw

Table IV. Equilibrium Constants Extrapolated to I = 0 Using the Pitzer and SIT Equations Species

Lit. value

NpO2Ac° UO2Ac+ AmAc 2+ ThAc3+ ThAc22+ NpO2Lac0 UO2Lac+ ThLac3+ ThLac22+

1.30(1) 3.03(5) 2.97(5)

a

— 1.75 ±—0.02(4)a — 5.49(11) —

Reported at an ionic strength of 0. 1 .

SIT (log P0101)

Pitzer (log P0101)

3.01 2.49 5.24 9.06 1.70 3.16 5.12 9.12

1.46(15) 3.03 2.69 6.59 11.2 1.97 3.09 6.83 11.2


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sured in low ionic strength, which are in the range of ±0.05 to ±0.20. When the data from this study were modeling using average values of B ( l ) and, for the other parameters, the values obtained by regression analysis, including the standard chemical potential, gave a better representation of the data than the approach of Pokrovsky et al.(27) The experimentally determined first stability constants for the actinides with acetate and lactate along with results from Pitzer modeling (represented as solid lines) are plotted in Figs. 2 and 3, respectively. Figure 4 lists the experimental values and the model results for the second apparent stability constant of Th4+ with lactate and acetate. For all stability constants, the Pitzer model gives an excellent representation of the data at I > 0.30m (NaCl) using three interaction parameters B(0), B (l) , and Cb. The parameters B(2) as well as mixing parameters of cationic, anionic ( - j j ) and ternary (Lijk) interactions were set equal to zero since no experimental information was available to separate these effects from the binary interactions. For 1:1, 1:2, and 1:3 charge-type electrolytes, the typical values for B(0) parameter range from -0.0095 to 0.188, 0.019 to 0.42, and -1.12 to 0.82, respectively.(21) Therefore, using an averaged B (l) in this work yields reasonable values for B(0). However, values for C*, determined in this work, are higher than those typically reported by Pitzer for systems of the same charge type.

Fig. 2. Experimental points and fit from modeling for the 1:1 stability constant for Am 3+ , Th4+, and UO2+2 Complexation with the acetate ion as a function of ionic strength.

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Fig. 3. Experimental points and fit from modeling for the 1:1 stability constant for Th4+, NpO+2, and UO22+ complexation with the lactate ion as a function of ionic strength.

Fig. 4. Experimental points and fit from modeling for the 1:2 stability constant for Th4+ complexation with lactate (•) and acetate (D) ions as a function of ionic strength.


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ACKNOWLEDGMENTS This work was performed as part of the Waste Isolation Pilot Plant (WIPP) Actinide Source Term Program, supported at Sandia National Laboratories by the United States Department of Energy under Contract DE-AC0494AL85000 and at Florida State University under contract AH-5590. REFERENCES 1. V. A. Vodovatov, V. B. Kolokoltsov, T. V. Kovaleva, L. C. Macirov, D. N. Suglebov, and V. T. Sles, Radiokimiya 17, 889 (1975). 2. F. Rosch, S. Dittrich, G. V. Buklanov, M. Milanov, V. A. Khalkin, and R. Dreyer, Radiochim. Acta 49, 29(1990). 3. A. I. Moskvin, and M. P. Mefod'eva, Radiochim. Acta 7, 410 (1965). 4. S. H. Eberle, and J. B. Schaefer, J. Inorg. Nucl. Chem. 31, 1523 (1969). 5. A. I. Moskvin, Soviet Radiochem. 11, 447 (1969). 6. J. Stary, and V. Balek, Collection Czechoslov. Chem. Commun. 27, 809 (1962). 7. V. G. Voden, M. E. Obukhova, and M. F. Pushlenkov, Soviet Radiochem. 11, 633 (1969). 8. V. K. Rao, G. R Mahajan, and P. R. Natarajan, Inorg. Chim. Acta 128, 131 (1987). 9. I. Grenthe, Acta Chem. Scand. 16, 1695 (1962). 10. N. A. Skorik, and A. S. Artish, Russ. J. Inorg. Chem. 30, 1130 (1985). 11. L. P. Lisovaya, and N. A. Skorik, Russ. J. Inorg. Chem. 18, 599 (1973). 12. 1. Feldman, and L. Koval, Inorg. Chem. 2, 145 (1963). 13. C. E. Harvie, N. Moller, and J. H. Weare, Geochim. Cosmochim. Acta 48, 723 (1984). 14. C. F. Novak, R. C. Moore, and R. V. Bynum, Proc. 1996 Intern. Confe. Deep Geological Disposal of Radioactive Wastes, Winnipeg, Manitoba, Canada, 16-19 Sept., 1996, SAND 96-2695C. 15. C. F. Novak, M. Borkowski, and G.R. Choppin, Radiochim. Acta 74, 111(1996). 16. G. R. Choppin, A. H. Bond, M. Borkowski, M. G. Bronikowski, J. F. Chen, N. A. LabonneWall, S. Lis, J. Mizeria, O. Pokrovsky, Y. Xia, and R. C. Moore, Report in preparation for Sandia National Laboratory. Currently unpublished data furnished by personal communication. (1998) 17. K. S. Pitzer, J. Phys. Chem. 77, 268 (1973). 18. A. R. Felmy, D. Rai, J. A. Schramke and J. L. Ryan, Radiochim. Acta 48, 29 (1989). 19. A. R. Felmy, D. Rai, and M. J. Mason, Radiochim. Acta 55, 177 (1991). 20. Th. Fanghanel, V. Neck, and J. L. Kim, Radiochim. Acta 69, 169 (1995). 21. K. S. Pitzer, Activity Coefficients in Electrolyte Solutions CRC Press, Boca Raton, FL, (1991). 22. J. Mizera, A. H. Bond, G. R. Choppin, and R. C. Moore, Proc. Am. Chem. Soc. Orlando, FL, August, 1996, to be published (1999). 23. Neck V, J. I. Kim, and B. Kanellakopulos, Thermodynamic Behavior of Neptunium (V) in Concentrated NaCl and NaClO4 Solutions Raport KfK 5301 (1994). 24. R. N. Roy, K. M. Vogel, C. E. Good, W. B. Davis, L. N. Roy, D. A. Johnson, A. R. Felmy, and K. S. Pitzer, 7. Phys. Chem. 96, 11065 (1992). 25. R. E. Mesmer, C. S. Patterson, C. H. Busey, and H. F. Holmes, J. Phys. Chem. 93, 7483 (1989). 26. A. T. Plyasunov, T. Fanghanel, and 1. Grente, Acta Chem. Scand. 52, 250 (1998). 27. O. S. Pokrovsky, M. G. Bronikowski, R. C. Moore, and G. R. Choppin, Radiochim. Acta 80,23(1998).