Dissolution Study of Metatorbernite ... - ACS Publications

Aug 30, 2010 - AND. ANDREW R. FELMY †. Pacific Northwest National Laboratory, 902 Battelle. Boulevard, Richland, Washington 99352, and Department of...
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Environ. Sci. Technol. 2010, 44, 7521–7526

Dissolution Study of Metatorbernite: Thermodynamic Properties and the Effect of pH and Phosphate E U G E N E S . I L T O N , * ,† JOHN M. ZACHARA,† DEAN A. MOORE,† JAMES P. MCKINLEY,† ALISON D. ECKBERG,† CHRISTOPHER L. CAHILL,‡ AND ANDREW R. FELMY† Pacific Northwest National Laboratory, 902 Battelle Boulevard, Richland, Washington 99352, and Department of Chemistry, The George Washington University, Washington, DC 20052

Received May 12, 2010. Revised manuscript received August 3, 2010. Accepted August 10, 2010.

The uranyl copper-phosphate, metatorbernite, has been identified in the shallow vadose zone of the 300 A area at the Hanford site, WA, USA. Consequently, modeling the evolution of U concentrations in vadose zone porewaters driven by meteoric water recharge requires accurate knowledge of metatorbernite solubility. Previous determinations of the solubility constant for metatorbernite were under constrained. In the present contribution, the dissolution of natural metatorbernite crystals was studied at target pH 2.5 and 3.0, using both nitric and phosphoric acid. Steady state was approached from underand supersaturation. The experiments and calculations yielded a preferred log Ksp ) -28.0 ( 0.1 that is significantly different than previously determined values. Further, both stoichiometric and nonstoichiometric dissolution was observed as a function of pH and aqueous phosphate concentration.

Introduction Uranium is a common contaminant in the environment and the most abundant radionuclide in the subsurface at U.S. Department of Energy facilities (1). It is well-known that the solubility, and hence mobility, of U in the subsurface is strongly governed by its oxidation state. For example, U(VI) is generally more soluble than U(IV) which can form sparingly soluble minerals such as UO2. Consequently, one strategy for restricting U mobility is to reduce it to and maintain it as U(IV). However, maintaining reducing conditions poses an ongoing stewardship problem. This issue has been recognized and has stimulated research on the viability of precipitating uranyl as sparingly soluble phosphates, such as the autunite and meta-autunite mineral groups, at reactive barriers (2) and using more dispersive techniques such as injecting polyphosphate into wells (3). In addition to their promise as environmental engineering solutions to U mobility, the autunite and meta-autunite mineral groups are of geochemical and environmental importance because they are ubiquitous oxidation products of uranium ores (4-6) and potentially control the mobility * Corresponding author phone: 509-371-6387; fax: 509-371-6354; e-mail: [email protected]. † Pacific Northwest National Laboratory. ‡ The George Washington University. 10.1021/es101619f

 2010 American Chemical Society

Published on Web 08/30/2010

of U in uranium contaminated sediments (7, 8) including those at DOE sites such as Oak Ridge, TN (9) and at the Hanford site, WA (10-14). In particular, detailed investigations of U contaminated sediments beneath historic processing ponds in the 300 A area at the Hanford site indicate that metatorbernite, with the chemical formula Cu[(UO2)(PO4)]2 · 8H2O, is present in certain shallow vadose zone sediments. It was unambiguously identified in the vadose zone at ∼4′ below the excavated level of the North Processing Pond 2 (NPP2-4) with microdiffraction and transmission electron microscopy, (12, 14) where it was estimated that metatorbernite accounts for ∼55% of total uranium (11). Further, analytical electron microscopy of the metatorbernite grains only detected U, P, Cu, and O (14). Other unidentified U(VI) phases were also observed in this sediment as well as immediately above and below this depth interval. These cumulative discoveries have led to a conceptual model (10) for the release and transport of U from vadose zone source terms to groundwater that involves the dissolution of metatorbernite and other precipitated phases in the shallow vadose zone during recharge events (rain and snowmelt), followed by transport and adsorption of uranyl to mineral coatings in the deep vadose zone. Cumulative recharge, however, is generally low at this location because of the semiarid climate. Release of U to groundwater likely occurs during periods of spring high water table. Transforming this conceptual model into one that can make quantitative predictions of U flux through the vadose zone and into groundwater requires, in part, accurate information on the solubility of metatorbernite and other U(VI) phases that may be present. Further, mechanistically based models of metatorbernite dissolution kinetics require accurate estimates of its Gibbs free energy of formation in order to calculate the free energy of reaction. This information would also be useful for characterizing current or historical uranyl transport at other localities where metatorbernite has formed, such as down gradient of the Koongara ore deposit (4, 5). Previous determinations of metatorbernite solubility indicate that it is one of the least soluble members of the autunite family (6). However, past experimental work has been under constrained due to lack of reversals (i.e., approach to equilibrium from above and below saturation) and assumptions concerning reaction stoichiometry (e.g., 6). In this contribution, we investigated the dissolution of very pure and well crystallized natural metatorbernite over the pH range 2.5-3.0, where low pH was preferred due to the likely low solubility of the material, using either nitric acid or phosphoric acid to adjust the pH, and approached equilibrium from both undersaturation and supersaturation. All major components (i.e., U, P, and Cu) as well as potential minor cations and anions were measured in solution. The objectives were to obtain a well constrained solubility product for metatorbernite and to examine dissolution stoichiometry under varying solution composition.

Materials and Methods Minerals. Natural metatorbernite samples were purchased from the Mineralogical Research Company, San Jose, CA, and originated from the Musonoi Mine, Shaba Province, Zaire, Africa. Metatorbernite crystals grew as gem-like, greater than millimeter sized emerald-green tabular crystals, with the bottom layer of material embedded in a friable aluminosilicate matrix. It was relatively effortless to collect a few grams of metatorbernite from the top layers of two (samples MN and H) individual fist-size chunks of sample+matrix with VOL. 44, NO. 19, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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minimal disturbance of the matrix. The harvested metatorbenite crystals, free of contamination by visual inspection, were then sonicated in DI water; the DI water was decanted; and the remaining solid was crushed by mortar and pestle until the powder became a uniform greenish-white color. Solids were analyzed by bulk wet chemical analysis, bulk powder, and in situ XRD, EMPA, and SEM [details in the Supporting Information (SI)]. Powder XRD of the crushed material indicated no phases other than metatorbernite. Single crystal XRD and subsequent refinement indicated that the material had unit cell parameters (P4/n, a ) 6.9664(2), c ) 17.313(11) Å, R1 ) 1.48%, wR2 ) 4.29%), consistent with those of metatorbernite given by Locock and Burns (15). Bulk digestion and analysis by ICPOES for both major and potential minor cations indicated that the samples were >99% pure metatorbernite, with U:P: Cu ratios within error of expected stoichiometric values. These results were confirmed by EMPA analyses of polished mounts. Nonetheless, test dissolution experiments at pH 2.5 (adjusted with nitric acid) indicated a high nonstoichiometric release of Cu, which could be accounted for by complete dissolution of only ∼0.27 (sample MN) and ∼0.01 (sample H) wt % Cu equivalents of a secondary Cu-rich phase or possibly preferential release of Cu from metatorbernite. A salient question, then, is whether minor chernikovite (H3O+UO2PO4 · 3H2O) formed or whether a secondary Cu-rich phase was present. Although powder XRD and bulk wet chemical analyses of the original and pretreated material did not detect a secondary Cu phase or alteration products, neither method would detect the small amount of material required to produce the excess Cu(aq). However, reuse of material from the test dissolution experiments yielded release of U, P, and Cu in stoichiometric metatorbernite proportions at target pH 2.5. In fact, multiple reuse of material at pH 2.5 always yielded stoichiometric metatorbernite dissolution. This result strongly suggests that chernikovite is not a significant phase and certainly did not control dissolution. A search with SEM and EMPA did not find a high Cu-phase. Further constraints on this issue were provided by the dissolution experiments themselves including the approach to equilibrium from both below and above saturation, as described in the Results and Discussion. Although excess Cu(aq) is not necessarily an intrinsic problem for determining the solubility of metatorbernite, it was a practical issue, as the high copper severely depressed both U(aq) and P(aq) and interfered with ICP-OES determinations of P(aq). Consequently, the powdered material was subjected to pretreatment that consisted of acid washing (pH 2.5) and decanting of ultrafines prior to the dissolution experiments. SEM of powder mounts of the pretreated material indicated a marked decrease in the abundance of fines, with a particle size range from 1 to 50 µm across. Solubility Experiments. Batch experiments were conducted under N2 in a glovebox. The reaction vessels were acid-washed 80 mL Teflon bottles, which were continuously agitated on an orbital shaker at 80-100 rpm. All acids that were used to adjust the solution pH were of ultrahigh purity grade and arrived in Teflon bottles from the supplier (ultrapure nitric acid, GFS, double distilled, assay 66-70%; ultrapure phosphoric acid, GFS, assay 85% solution). Solutions were made from DI water (18 MΩ cm) that was sparged with N2. The pH measurements were always made on pipetted volumes (polypropylene pipet tips), never directly on active experimental solutions. In sum, during all preparation steps and the experiments themselves, great care was taken to make sure that solutions were never in contact with glass. This precaution was necessary as we found that even short exposure times of solutions with glass ware yielded elevated concentrations (∼30-100 µmol/L) of alkali and alkaline Earth cations, Si, and Al, whereas taking measures to avoid contact 7522

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with glass reduced concentrations of these elements by about 1 order of magnitude or more. Initial experiments that did not take such precautions invariably showed nonstoichiometric (i.e., elevated) release of Cu. In contrast, exclusion of all contact with glass, combined with the acid pretreatment described above, yielded near stoichiometric Cu release. We speculate that these “contaminant” cations were exchanging with interlayer Cu. Further, because the autunites are only sparingly soluble, even at relatively low pH, it was prudent to decrease the nonmetatorbernite components as much as possible to minimize potential effects on the solution speciation of U, Cu, and P. Solubility studies were conducted at 22 ( 1 °C and target pH 2.5 and 3.0 with two samples labeled MN and H. The samples have the same provenance, but were taken from different hand specimens. Weighed amounts of the pretreated material (0.75 g) were reacted with 60 mL of solution. The initial pH of the experiments was adjusted with nitric or phosphoric acid, and then measured at timed intervals, as the experiments were sampled, without further adjustment. At pH 2.5, steady state was approached from both undersaturation and supersaturation. For the approach from undersaturation, the solid phase was simply reacted with the acidified solutions. For the approach from supersaturation, the experiments were initiated at undersaturated conditions, reacted for 7 days, and then spiked with a small volume (0.05 mL) of 0.04 mol/L Cu(NO3)2 solution, preadjusted to pH 2.5, such that the final Cu(aq) concentration in the experimental solutions would have been 100 µmol/L in lieu of prior reaction (i.e., about a factor of 7 higher than present just prior to the spike). The Cu spike only increased the solution volume by 0.25%. The solutions were sampled at timed intervals for measurement of cations and anions and pH. The orbital shaker was stilled, the metatorbernite rapidly settled (about 5 min) due to its large grain size, and 3-5 mL of the clear supernatant pipetted off the top. A portion was set aside for pH analysis using an Orion, Ross Ultra glass combination semimicro pH electrode calibrated against pH 2 and 4 buffers which bracketed the experimental conditions. One ml was then sacrificed by passing it through a 4 nm pore-size centrifuge-filter tube that had been acid washed at the experimental pH. The remaining solution sample was then centrifuged-filtered and analyzed by ICP-OES (PerkinElmer 2100DV) and IC (Dionex ICS-2000) for cations and anions (Cl-, F-, SO42-, NO32-, and PO43-). P was analyzed primarily by ICP, with spot checks by IC. A wide range of ions were analyzed for, in part because the material was collected from natural samples. See SI for more detailed information on solution analyses.

Results and Discussion Solid Characterization. At the conclusion of the experiments, mineral residues, still in contact with the experimental solution, were characterized by in situ micro-XRD. In all cases, metatorbernite was the only phase detected. The XRD analyses were conducted under in situ conditions in order to check for potential additional hydration of the interlayer region and possible transformation of metatorbernite to torbernite (Cu[UO2PO4]2 · 12H2O). Although we cannot rule out the possibility that an outer alteration layer of torbernite formed that was below detection, the transformation involves more than simple hydration; the stacking geometry changes and the c cell dimension increases by ∼20% (15). Thus, regardless of the relative energetics of the two phases, there is likely a significant activation energy barrier to transformation of metatorbernite to torbernite. Solution compositions. Samples MN and H behaved in a very similar manner. Consequently, figures include data from only one sample but the full suite of solution data is

FIGURE 1. Nitric acid experiments at pH 2.5: dissolution of metatorbernite (sample H) as a function of time, where pH was adjusted by nitric acid. (A) Approach to equilibrium from undersaturation, where the insert only shows data from 0-8 days and (B) approach to equilibrium from oversaturation. Oversaturation was achieved by spiking the solution with Cu(NO3)2 after 7 days of equilibration. given in SI, Table S2. Experiments at the target pH 2.5 that used nitric acid and were conducted under initially undersaturated conditions reached steady state very rapidly, i.e., by ∼6 h (Figure 1a). Further, dissolution was stoichiometric for U, P, and Cu within experimental error, where the average U:P ratio ) 0.98 (2σ ) (0.03) for both samples MN and H and U:Cu ) 1.92 (2σ ) (0.18) and 2.03 (2σ ) (0.30) for samples MN and H, respectively. The pH values showed no trend from 0.25 to 29 days, where the average pH is 2.46 +0.04/-0.03 for sample MN. The next, and last, sampling occurred at 63 days, where the pH was higher than the highest pH reading over the 0.25-29 day period. The pH at 63 days is significantly different than the mean pH over 0.25-29 days for sample MN. However, for sample H, the pH values of the 63 day and the mean of the 0.25-29 day sampling periods overlap at 2σ uncertainty, indicating that the difference, although possibly real, is not significant. After solutions had reached steady state from undersaturation, equilibrium was approached from supersaturation (reversals) by spiking the experiments with Cu(NO3)2. Both U(aq) and P(aq) concentrations dropped significantly after Cu(NO3)2 was spiked into the nitric acid, target pH 2.5 solutions (e.g., Figure 1b). Ideal stoichiometric ∆P:∆U ratios, where ∆ is the difference between the pre- and postspike concentrations, should be 1 for metatorbernite precipitation, whereas measured values for sample H and MN were 1.21 and 1.11, respectively, 7 days after the Cu spike (or day 14 in Figure 1b). P(aq) concentrations dropped slightly more than U(aq) concentrations, indicating that the induced precipitation of P and U was nearly, but not quite, stoichiometric with respect to metatorbernite. Nonetheless, decreases

FIGURE 2. Phosphoric acid experiments at pH 2.5: dissolution of metatorbernite (sample H) as a function of time, where pH was adjusted by phosphoric acid. (A) Approach to equilibrium from undersaturation where the insert only shows data from 0-15 days and (B) approach to equilibrium from oversaturation where the insert only show U and Cu concentrations from 0-5 µmol/L. Oversaturation was achieved by spiking the solution with Cu(NO3)2 after 7 days of equilibration. in both P and U indicate that steady-state was approached from supersaturation and that precipitation of a metatorbernite-like phase is allowed by the data. Nitric acid experiments conducted at the target pH 3.0 indicated a close approach to steady state conditions from 7 to 35 days of reaction. Cu, U, and P concentrations were roughly a factor of 3 lower than at the target pH 2.5. However, dissolution was less stoichiometric at pH 3, where U:Cu and U:P ratios were 1.78 (2σ ) (0.25) and 0.90 (2σ ) (0.06), respectively. The results are consistent with preferential release of Cu and possibly P relative to U. The dissolution stoichiometry of metatorbernite as a function of pH is somewhat analogous to that for micas, where stoichiometric dissolution of the structural octahedral units and interlayer cations occurs at low pH, but preferential release of interlayer cations occurs at higher pH (16). Results for phosphoric acid experiments at target pH 2.5 under initially undersaturated conditions are illustrated in Figure 2. Concentrations of Cu and U were depressed ∼5-8 and ∼9-15 fold, respectively, relative to the equivalent nitric acid experiments. After the first 6 h of reaction aqueous U/Cu ratios ) ∼1.7-1.8. From 6 h to about 1 week, Cu(aq) rose sharply and then reached a plateau, but U(aq) decreased in mirror fashion such that U/Cu ratios approached ∼1. Sample H reached steady state by about 1 week. Sample MN reached steady state with respect to Cu by ∼4 days, whereas U appears to slowly evolve to lower concentrations up to 63 days; pH was invariant. It should be expected that higher concentrations of phosphate will depress metatorbernite solubility. In fact, Wellman et al. (17) showed that addition of phosphate depressed the dissolution rate of metatorbernite relative to the case of no added phosphate. However, a discussion concerning the relative behaviors of U(aq) and Cu(aq) is VOL. 44, NO. 19, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

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TABLE 1. Calculated log IAP Valuesa nitric acid experiments sample MN days 0.25 0.5 1 2 3 4 5 7 14 21 22 29 35 63

pH 2.5 v -28.12 -28.13 -28.17 -28.29 -28.25 -28.26 -28.44 -28.36 -28.24

pH 2.5

pH 3.0

pH 2.5

V

v

-28.39 -28.84

-27.92

v -28.4 -28.26 -28.31 -28.37 -28.34 -28.45 -28.55 -28.57 -28.16

b

phosphoric acid experiments sample H pH 2.5

pH 3.0

pH 2.5

V

v

-28.43 -28.63

-27.83

v -28.07 -27.94 -27.96 -28.06 -28.08 -28.03 -28.15 -28.02 -27.95

b

-27.98

pH 3.0

pH 2.5

V

v

-28.00 -28.18

-27.48

v -28.14 -28.06 -28.07 -28.14 -28.15 -28.11 -28.2 -28.06 -28.03

b

-28.18

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V

v

-27.70 -28.12

-27.61

-28.03 -28.2

-28.01

-27.92

a Steady state approached from undersaturation v and from oversaturation V. undersaturation, then spiked with Cu(NO3)2 to affect reversal.

warranted. As mentioned previously, Cu and U were released initially (0.25 day sampling) in near, yet significantly nonstoichiometric, proportions. If dissolution were to continue in like fashion, then U should have increased in proportion with Cu. Instead, U decreased as Cu increased. We speculate that this reflects dissolution processes at edge terminations where Cu and U are initially removed in near stoichiometric proportions but that the presence of high phosphate concentrations stabilized the uranyl-phosphate structural units such that the rate of U dissolution rapidly decreased relative to Cu, possibly stopping by 6 h. In contrast, Cu release continued up to about 4 days of reaction. In detail (see inset in Figure 2a), Cu(aq) was at a plateau from 0.25 to 2 days and then began to increase, reaching another plateau at about 4 days. U(aq) also appears to have recorded a plateau from 0.25-3 days, prior to decreasing. The same behavior was displayed by samples MN and H. The drop in U(aq) concentrations must indicate precipitation or adsorption of U(aq). Of course, one cannot exceed the saturation of metatorbernite from undersaturated conditions, so either another phase precipitated or uranyl adsorbed. Nonetheless, U(aq) and Cu(aq) concentrations reached near-steady state and the possible presence of an additional phase would not necessarily exclude using the data for determining Ksp for metatorbernite, as long as the system reached equilibrium. Thermodynamic modeling (see following section) suggests that this was indeed the case. Reversals, identical to those for the nitric acid experiments, were also attempted by spiking Cu(NO3)2 into the phosphoric acid, target pH 2.5 solutions that had already reached steady state values from undersaturation (e.g., Figure 2b). Measured U(aq) concentrations dropped by a factor of ∼7-9 after addition of Cu. Steady state was rapidly established, although any change in P(aq) was not detectable due to the high concentration of phosphate. Cu, P, and U concentrations in the phosphoric acid, target pH 3.0 experiments conducted at initially undersaturated conditions were at steady state by ∼7 days of reaction. The Cu(aq) and U(aq) concentrations were about a factor of 3 lower than the experiments at target pH 2.5, which parallels the behavior of the nitric acid experiments. Reversals were not attempted. As at pH 2.5, Cu was preferentially leached where U(aq):Cu(aq) ratios ) 0.9. Experimental solutions were analyzed for a wide range of cations and anions. In general, the cations Al, Ca, Mg, Na, K, Fe, and Si were detected at 0.2-10 µmol/L, with the majority of measurements at less than 5 µmol/L. Ca, Na, K,

b

pH 3.0

-28.19

-28.08 -27.36

-28.07

pH 2.5

-27.67

-28.21

-28.18 -27.51 -28.65

sample H

-27.54

-28.74

-28.67

pH 2.5

-27.68

-29.01 -28.84

sample MN

b

-28.2 -28.07

-28.07

Initially at steady state established from

and Fe were invariant with respect to reaction time. Al, Mg, and Si appeared to increase from about 14 to 63 days. The highest concentrations of Al, Mg, and Si were 11-22, 15-37, and 12-24 µmol/L, respectively, in the nitric acid experiments. Results were comparable for the phosphoric acid experiments. Because no cations were detected in blank controls, the presence of these cations either indicates that a secondary phase(s) was still entrained in the samples, despite the rigorous acid-washing procedures, or that metatorbernite contains these elements in solid solution, albeit at low concentrations. Cl and F were detectable in only a small subset of samples, whereas other anions such as SO42and CO32- were never detected. If metatorbernite was more soluble, such trace levels of “contamination” would not be of any concern. However, as described in the following section, the low aqueous concentrations of U, Cu, and P (in the nitric acid experiments) required an assessment of the potential sensitivity of calculated thermodynamic parameters to complete solution compositions. Thermodynamic Modeling. Thermodynamic modeling calculations were performed using the chemical equilibrium model GMIN (18) which utilizes the ion-interaction model of Pitzer and co-workers (19, 20). In this chemical system phosphate can form complexes with both U6+ and Cu2+ which could significantly impact the solution speciation and the calculation of the metatorbernite solubility product. In addition, in the lower concentration phosphate solutions, the presence of small concentrations of both Mg2+ and Al3+ could possibly impact the speciation of phosphate and the thermodynamic analysis. As a result, stability constants for U6+, Cu2+, Mg2+, and Al3+ hydrolysis species and phosphate complexes were included in the chemical modeling calculations as well as ion-interaction parameters for the most important electrolyte components. In these dilute solutions the principal impacts on the solution speciation are related to the formation of relatively strong U6+ phosphate complexes. In fact, in the highest phosphate solutions, >80% of the total U6+ is in the form of U6+ phosphate complexes the most important being UO2H2PO42+, UO2(H2PO4)2+, and UO2HPO4(aq). Hydrolysis species do not contribute significantly to the final speciation calculations owing to the low pH. The dissolution of Al and Mg from the samples at the longest time frames has only a minor impact on the final speciation of phosphate in these samples. Values of log IAP (Table 1) were calculated for each data point for the dissolution reaction

Cu(UO2)2(PO4)2 · 8H2O + 2H+ ) Cu2+ + 2UO22+ + 2HPO42- + 8H2O

(1)

Provisional log Ksp values, from averaged ion activity product (IAP) values for individual experiments, are given in Table 2; where the average of all the data yields a provisional log Ksp ) -28.04 + 0.2(1σ)/-0.8(1σ). The magnitude of error associated with this provisional solubility constant was investigated in detail, as discussed below. For the initially undersaturated phosphoric acid experiments at the target pH 2.5, log IAP values indicate no discernible trend from 0.25 to 63 days (Figure 3), despite sharply increasing and decreasing Cu(aq) and U(aq) concentrations from 0.25 to 5 days (Figure 2a). Further, there is no significant difference between the average log IAPs for the different samples or between the average log IAPs for the initially undersaturated and supersaturated experiments (Table 2). Averaging all the IAP values for both samples yields log Ksp(0.25-63) ) -28.07 ( 0.1. In general, the phosphoric acid experiments are consistent with a close approach to equilibrium, and show much less variation than the full data set. However, at issue is whether one should accept solubility data that is changing. Consequently, we calculated log Ksp values for both samples that only included the 7-63 and 63 day data for the initially undersaturated and oversaturated experiments, respectively. Both samples yield log Ksp values that are not significantly different from each other or the log Ksp values derived from the fuller data set (Table 2). Averaging the IAP values for both samples (using the 7-63 day data) yields log Ksp(7-63) ) -28.02 ( 0.1, which is marginally lower than, but within error of log Ksp(0.25-63) that was derived from all the phosphate data. The calculations indicate that the system was at near equilibrium. For the initially undersaturated nitric acid experiments at target pH 2.5, average log IAP values (0.25-63 days) for both samples were nearly identical and within error of the phosphoric acid experiments (Table 2). However, both samples yield average log IAP values that are ∼0.2-0.3 log units lower than for the phosphoric acid experiments with appreciably higher associated uncertainties. In essence, the agreement between the nitric and phosphoric acid experiments is due, in part, to the higher variability of the former. Similarly, the average log IAP values for the initially oversaturated experiments are the lowest recorded but not significantly different than the initially undersaturated experiments or the phosphoric acid experiments. However, they also show greater variability than their phosphoric acid counterparts, and the lack of significant differences is due to their larger associated uncertainty (Table 2). Averaging all

FIGURE 3. Values of log IAP for sample MN as function of time at pH 2.5, where pH was adjusted by phosphoric acid. Open and closed symbols indicate approach to equilibrium from under- and oversaturation, respectively. the IAP values from both samples yields log K(0.25-63) ) -28.35 +0.28/-1.12. At target pH 3, averaging all the data (12 points from undersaturation) yields log Ksp ) -27.68 +0.2(1σ)/ -0.4(1σ), which is within error of, but ∼0.3 log units less than, the provisional log Ksp that incorporates all the data. It is also within error of log Ksp values determined from the pH 2.5 data, although this is due to the higher uncertainty in the pH 3 data. The relatively high uncertainty in the target pH 3 data (compare to pH 2.5 phosphoric acid experiments) occurs under both nitric and phosphoric acid conditions, although their calculated log Ksp values are nearly identical (Table 2). Consequently, much of the uncertainty in the apparent log Ksp for the full data set is from variations associated with the pH 2.5 nitric acid and the pH 3 experiments. A source of this greater variability is, in part, the greater fluctuations in the measured pH compared to the pH 2.5 phosphoric acid experiments (Table 2). Given that the dominant phosphate species was H3PO4, such that 6 mol of H+ are consumed for every 2 mol of U and P released, the calculated IAP values are very sensitive to pH. Consequently, small changes in pH can induce noticeable variation in the calculated average log Ksp if the solubility does not respond, which appears to be the case. However, we do not necessarily expect the solubility to correlate to such small changes in pH because the pH fluctuations are not systematic with time and probably relate, in part, to measurement error (although we cannot discount the possibility of a small pH increase during the pH 3 experiments).

TABLE 2. Log Ksp (I = 0, 22 ± 1°C) Values for Cu(PO4UO2)2 · 8H2O + 2H+ = Cu2+ + 2HPO42- + 2UO22+ + 8H2Oa nitric acid sample

days

MN v MN V Hv HV (MN+H)vV (MN+H)vV MN vV H vV (MN+H) vV (Mn+H) v

0.25-63 14-63b 0.25-63 14-63b 0.25-63 0.25-63c 7-63d 7-63d 7-63d 7-35

Log Ksp (2σ) -28.30 -28.81 -28.31 -28.45 -28.35 -28.31

(+0.25/-0.69) (+0.26/-0.75) (+0.24/-0.57) (+0.39/e) (+0.28/-1.12) (+0.25/-0.62)

-27.66 (+0.34/e)

phosphoric acid pH (2σ)

Log Ksp (2σ)

pH (2σ)

2.48 (+0.1/-0.08) 2.49 (+0.1/-0.08) 2.44 (+0.1/-0.08) 2.48 (+0.1/-0.08) target pH ) 2.5

-28.03 -28.08 -28.09 -28.12 -28.07

(+0.11/-0.15) (+0.25/-0.68) (+0.10/-0.12) (+0.11/-0.14) (+0.10/-0.12)

2.51 (+0.07/-0.06) 2.51 (+0.03/-0.03) 2.49 (+0.06/-0.05) 2.51 (+0.02/-0.02) target pH ) 2.5

3.01 (+0.05/-0.05)

-27.99 -28.05 -28.02 -27.70

(+0.11/-0.15) (+0.04/-0.04) (+0.10/-0.13) (+0.33/e)

2.54 2.50 2.52 3.04

(+0.06/-0.06) (+0.04/-0.03) (+0.03/-0.05) (+0.15/-0.11)

a

Values of log IAP averaged over the time period (days) specified. Bolded log Ksp is a preferred value, see text. Approach steady state from under v and over V saturation. b Time period refers to first sampling at 14 days after initiation of experiments from undersaturation and 7 days after Cu(NO3)2 spike. c Excluding 14 and 22 day Cu(NO3)2 spike experiments. d Average log IAP only includes 7-63 and 63 day data for undersaturation and oversaturation experiments, respectively. e 2σ larger than IAP.

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The thermodynamic calculations also showed that all of the experimental points were significantly undersaturated with respect to Na(UO2PO4) · nH2O, HUO2PO4 · 4H2O, and (UO2)3(PO4)2 · 4H2O indicating that metatorbernite is the most thermodynamically stable phase in the system. Thus, it is not clear what uranyl bearing phase might have precipitated in the pH 2.5 phosphoric acid experiments. The general correspondence of calculated log Ksp values for experiments conducted under very different conditions (i.e., initially under or oversaturated, pH adjusted by nitric or phosphoric acid, and pH 2.5 or 3.0) indicates that equilibrium was approached. This correspondence reflects the fact that the solubility of metatorbernite responded, as predicted, to the different initial conditions. In particular, if minor chernikovite formed at edges it cannot be controlling equilibrium as one would not expect to obtain such close reversal brackets. We recommend a preferred solubility constant log Ksp ) -28.0 ( 0.1 (bolded in Table 2) calculated from the restricted pH 2.5 phosphoric acid experiments. We justify emphasis of the pH 2.5 phosphoric acid experiments because of the low uncertainty and very tight reversal brackets which suggest a close approach to equilibrium. It is likely that the high concentration of phosphate buffered the pH, which led to the relatively high precision of the calculated solubility constant. Nonetheless, this preferred solubility constant is nearly identical to the provisional solubility constant that was derived from all the data, where the primary difference is in their associated uncertainties. The preferred solubility constant is significantly higher than log Ksp ) -30 ( 0.1 given by Magalhaes and De Jesus (6). In contrast, our preferred value is closer to, but lower than log Ksp ) -27.2 derived from data in Vochten et al., (21) using the reaction stoichiometry given in the present contribution. The Magalhaes and De Jesus (6) study was performed at pH 2.83 with a phase they called torbernite that contained ∼9 H2O per unit formula (characterized thermogravimetrically and with powder XRD). The phase was formed by reacting Cu solutions with synthetic HUO2PO4 · 4.2H2O. Experiments were analyzed for U(aq) and Cu(aq) but not P(aq). Based on this they assumed that dissolution was stoichiometric. Vochten et al. (21) also assumed stoichiometric release, but did investigate dissolution over a range in pH. As we have seen, the assumption of stoichiometric release appears to be valid at pH 2.5 but not pH 3. Further, in both studies, steady state was only approached from undersaturation. We suggest that the log Ksp provided in the present study is better constrained than in previous work. Whether the effect of aqueous phosphate on dissolution stoichiometry holds for other autunites is an open question.

Acknowledgments Portions of this research were performed at the Environmental Molecular Sciences Laboratory at PNNL, a national user facility operated by Battelle on behalf of the U.S. DoE, OBER. Research performed by the PNNL Scientific Focus Area with funding support from DOE Biological and Environmental Sciences Division (BER) through the Subsurface Biogeochemistry Program (SBR). The authors are grateful to Mark Frisch and Paula Cantos (GWU) for assistance with single crystal X-ray data collection. X-ray Instrumentation at GWU was purchased with NSF Funding (DMR-0419754) and CLC was partially supported by DOE under Grant No. DEFG02-05ER15736.

Supporting Information Available Descriptions of facilities used for solid state and solution analyses; results of solid state analyses including bulk powderXRD, in situ micro-XRD, SEM, and EMPA; all cation and anion solution analyses; bulk analyses of starting materials; 7526

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thermodynamic data for both solution species and solid state phases. This material is available free of charge via the Internet at http://pubs.acs.org.

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