Scientific Basis for Efficient Extraction of Uranium from Seawater. I

Jan 27, 2016 - User Resources. About Us · ACS Members · Librarians · ACS Publishing Center · Website Demos · Privacy Policy · Mobile Site ...
1 downloads 0 Views 724KB Size
Download PDF
Article pubs.acs.org/IECR

Scientific Basis for Efficient Extraction of Uranium from Seawater. I: Understanding the Chemical Speciation of Uranium under Seawater Conditions Francesco Endrizzi,# Christina J. Leggett,§ and Linfeng Rao* Chemical Sciences Division, Lawrence Berkeley National Laboratory, One Cyclotron Road, Berkeley, California 94720, United States S Supporting Information *

ABSTRACT: In recent years, the prospective recovery of uranium from seawater has become a topic of interest owing to the increasing demand for nuclear fuel worldwide and because of efforts to find sustainable alternatives to terrestrial mining for uranium. To date, the most advanced and promising method of extracting and concentrating uranium from seawater involves the use of polymeric sorbents containing the amidoxime binding moiety. Among a number of different moieties investigated, glutaroimide-dioxime is the most promising one, forming very stable complexes with U(VI) even in the presence of carbonate. To properly assess the affinity of uranium toward the amidoxime substrates, a comprehensive knowledge of the aqueous chemical equilibria of uranium is required. With this aim, in this paper we review the chemical equilibria of uranium (as UO22+) in seawater, focusing on the solution equilibria leading to the formation of the stable complexes, Mm(UO2)(CO3)3(2m−4)(aq) (M = Ca or Mg, m = 0−2). These binary and ternary species dominate the chemistry of uranium in seawater and have recently been the object of study in several papers in the literature. The solubility equilibria of UO22+ in seawater leading to the formation of the known minerals, including Liebigite, Ca2(UO2)(CO3)3·10H2O(cr), Swartzite, CaMg(UO2)(CO3)3·12H2O(cr), Bayleyite Mg2(UO2)(CO3)3·18H2O(cr), and Andersonite, Na2Ca(UO2)(CO3)3·6H2O(cr), are also critically reviewed. Newly calculated values of the solubility products (log K0s) for these solid compounds are presented based on the currently proposed speciation model that includes the most recent aforementioned data for the aqueous speciation of UO22+. Based on these data, simulated speciation diagrams are calculated, both at zero ionic strength and in seawater-like media. In combination with the speciation data for uranium with glutaroimide-dioxime, these models provide a better, more comprehensive picture of the chemical equilibria of U(VI) in seawater while also providing useful tools to help assess the feasibility of its recovery through amidoxime-based collection systems.

■

INTRODUCTION

the highest concentrations include sodium chloride, which accounts for 3.5% of the total ions by weight and essentially defines the seawater ionic strength, along with calcium- and magnesium chlorides with concentrations of 0.010 and 0.053 mol dm−3, respectively. At substantially lower concentrations, other heavier metal cations are present such as iron, copper, vanadium, and, importantly, uranium (14 × 10−9 mol dm−3).1 Though uranium in seawater is very dilute, it is of interest because the world’s oceans collectively contain nearly 4.5 billion tons of uranium,2 approximately 1000 times the total known terrestrial supply. In light of the rapid expansion of nuclear power in countries such as China and India, the development of an efficient and economical method of recovering this uranium could provide an essentially limitless source of fuel for nuclear reactors. To date, the most advanced and promising method of extracting and concentrating the dilute uranium from seawater involves the use of polymeric sorbents containing the amidoxime binding moiety. These poly(amidoxime) sorbents

Seawater is a matrix with a rich chemistry due to the presence of a variety of dissolved ions (see Table 1). Dissolved salts in Table 1. Concentrations of Selected Elements in Seawaterc concentration mg kg−1

element Na Cl Ca Mg K Li U TICa Fe Pb Ni Cu V S DIPb a

10.8 19.4 0.413 1.29 0.40 0.18 3.3 0.291 3.4 30 5 1 1.83 0.90 0.71

× × × × ×

103 103 103 103 103

× × × × × × × × ×

10−3 103 10−3 10−6 10−3 10−3 10−3 103 10−3

mol dm−3 0.468 0.546 10.3 53 10.2 26 14 24.2 0.5 0.01 8 3 36 28 2.3

× × × × × × × × × × × × ×

b

10−3 10−3 10−3 10−6 10−9 10−3 10−9 10−9 10−9 10−9 10−9 10−3 10−6

Special Issue: Uranium in Seawater Received: October 2, 2015 Revised: January 16, 2016 Accepted: January 27, 2016

c

Total Inorganic Carbon. Dissolved Inorganic Phosphorus. From ref 1, p 224. © XXXX American Chemical Society

A

DOI: 10.1021/acs.iecr.5b03679 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Article

Industrial & Engineering Chemistry Research Table 2. Thermodynamic Parameters for Relevant in-Solution Equilibria and the Solubility of UO22+ Speciesd aqueous species, reactions Uranyl hydrolysis UO22+ + H2O = (UO2)(OH)+ + H+ UO22+ + 2H2O = (UO2)(OH)2(aq) + 2H+ UO22+ + 3H2O = (UO2)(OH)3− + 3H+ UO22+ + 4H2O = (UO2)(OH)42− + 4H+ 2UO22+ + 2H2O = (UO2)2(OH)22+ + 2H+ 3UO22+ + 4H2O = (UO2)3(OH)42+ + 4H+ 3UO22+ + 5H2O = (UO2)3(OH)5+ + 5H+ Uranyl-carbonato and ternary complexes UO22+ + CO32− = UO2(CO3)(aq) UO22+ + 2CO32− = UO2(CO3)22− UO22+ + 3CO32− = UO2(CO3)34− Ca2+ + UO22+ + 3CO32− = CaUO2(CO3)32−

log β0 ± 3σ I≈0 −5.25 −12.15 −20.25 −32.40 −5.62 −11.9 −15.55

9.94 16.61 21.84 27.00 27.18 27.27 2Ca2+ + UO22+ + 3CO32− = Ca2UO2(CO3)3(aq) 30.84 29.8 30.8 30.70 29.8 Mg2+ + UO22+ + 3CO32− = MgUO2(CO3)32− 26.25 26.11 26.2 Solubility products of known (Na, Ca, Mg)-UO2-CO3·xH2O minerals Liebigite, Ca2(UO2)(CO3)3·10H2O(cr)

Swartzite, CaMg(UO2)(CO3)3·12H2O(cr)

Bayleyite, Mg2(UO2)(CO3)3·18H2O(cr)

Andersonite, Na2Ca(UO2)(CO3)3·6H2O(cr)

± ± ± ± ± ± ±

log βM ± 3σa I = 0.50 M NaCl −5.6 −12.5 −20.3 −31.8 −6.1 −12.7 −16.4

± ± ± ± ± ± ±

0.3 0.1 0.4 0.8 (p.w.)c 0.1 0.4 0.3

± 0.03 8.61 ± 0.09 15.26 ± 0.04 21.85 ± 0.12 24.28 ± 0.18 ± 0.42 ± 0.12 26.81 ± 2.0 ± 0.8 ± 0.15 ± 0.6 ± 0.12 23.62 ± 0.12 ± 0.4 log K0s ± 3σ (reactions)

± ± ± ±

0.04 0.10 0.07 0.21

0.24 0.07 0.42 0.68 0.04 0.3 0.12

ΔH0 ± 3σ, kJ/mol

ref 4 4 4 4 4 4 4

5.0 18.5 −39.2 −47

± 0.21

± 0.21

− Δf G0 ± 3σ, kJ/mol

± ± ± ±

2.0 4.0 4.1 6

4 4 4 11 8 10 −47 ± 7 11 6 7 8 10 11 8 15 − Δf H0 ± 3σ, kJ/mol

Ca2(UO2) (CO3)3·10H2O(cr) = 2Ca2+ + UO22+ + 3CO32‑ + 10H2O −(29.6 ± 2.7)16b 6226 ± 3616 7037 ± 7216 −(36.9 ± 6.3)17 6226 ± 3616,17 7029 ± 2418 c c −(33.9 ± 3.0) (p.w.) (5820 ± 9) (p.w.) CaMg(UO2) (CO3)3·12H2O(cr) = Ca2+ + Mg2+ + UO22+ + 3CO32− + 12H2O −(30.1 ± 5.7)16b 6607 ± 2416 7535 ± 6016 −(37.9 ± 4.2)17 6607 ± 2416,17 −(33.4 ± 6.0) (p.w.)c (6199 ± 9) (p.w.)c Mg2(UO2) (CO3)3·18H2O(cr) = 2Mg2+ + UO22+ + 3CO32− + 18H2O −(29.1 ± 4.2)b 7924 ± 2416 9192 ± 6016 −(36.6 ± 4.2)17 7924 ± 2416,17 9110 ± 2718 c c −(31.7 ± 4.5) (p.w.) (6682 ± 4) (p.w.) Na2Ca(UO2) (CO3)3·6H2O(cr) = 2Na+ + Ca2+ + UO22+ + 3CO32− + 6H2O −(30.3 ± 8.7)16b 5209 ± 5116 5916 ± 10816 −(37.5 ± 6.3)17 5651 ± 7216

Calculated in this paper from log β0 with the Specif ic-ion Interaction Theory empirical parameters19 selected in ref 4. bEstimated in this paper from the analysis of the original solubility data by Alwan and Williams16 (Figure 2, see also details in the Supporting Information). c(p.w.): proposed in this paper (calculated with the corresponding values of the (suggested) solubility products (b), and up-to-date values of the standard formation energies of individual species in ref 4 see also Table S2 in the Supporting Information). dFor the solid minerals, the values of log K0s and − Δf G0 recommended by this paper are in bold face. The uncertainties on the values are given at 3σ, statistically defined as the 95% confidence interval. In the text, the uncertainties are given at 1σ confidence interval, as reported in the original sources. a

more selective for uranium and lead to higher sorption of uranium. Simple thermodynamic models can facilitate the development of these new sorbents by predicting parameters such as the binding strength of uranium with various ligand moieties that could be present on the sorbent. Since the model describing this reaction requires knowledge of the nature of the uranium species present in seawater, it is evident that an understanding of the solubility and composition of uranium species in seawater is a crucial factor for model development and ultimately for designing new sorbents that meet the abovementioned criteria. Although hexavalent uranium (UO 22+) has a strong propensity toward hydrolysis and precipitation4 in ligand-free

are typically in the form of long braids, one end of which is anchored to the bottom of the ocean while the other end is allowed to float freely in the ocean. The interaction of the freestanding braids with ocean currents facilitates sorption of uranium (and other cations) to the poly(amidoxime) braids. After a predetermined amount of time, the braids are removed from the ocean, and the sorbed uranium is eluted using a suitable eluant. Poly(amidoxime) sorbents have been shown to yield promising results in marine tests − up to 3.3 g U/kg sorbent were obtained after 8 weeks3 − but the cost of such a method is still prohibitive and high sorption of other seawater cations (e.g., vanadium) lowers uranium uptake. As a result, research is underway to develop polymeric sorbents that are B

DOI: 10.1021/acs.iecr.5b03679 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Article

Industrial & Engineering Chemistry Research

complexes of the form Cam(UO2)(CO3)3(2m−4)(aq) (m = 1, 2), the individual spectral contribution of the two successive species is mainly determined by their different fluorescence lifetimes. The complexation constant of the aqueous species, Ca2(UO2)(CO3)3(aq), was determined to be log β0(213) = (29.2 ± 0.3) for the first time by Bernhard et al. using TRLFS.5 They later published a further study in 2001 in which they again used TRLFS to validate their previous results, conducted EXAFS experiments on solutions of uranyl carbonate, and conducted crystallographic analysis of Liebigite.7 TRLFS experiments were carried out using two independent methods: 1) direct titrations of uranyl triscarbonato solutions with Ca2+ solutions and 2) competitive titrations with EDTA. The two independent values of log β0(213) obtained using these methods were (30.45 ± 0.35) and (30.77 ± 0.25), respectively. The EXAFS experiments on solutions of uranyl carbonate in the presence and absence of Ca2+ were conducted with the aim of obtaining in-solution structural information about the ternary (Ca2+/Mg2+)-UO22+-CO32− species. In particular, the objective was to ascertain whether the interaction of calcium with the (UO2)(CO3)34− core led to an inner- or outer-sphere complex. Unfortunately, the analysis of the data did not provide definitive information since the spectral changes in the presence and absence of calcium were insufficient to provide a clear model. However, the EXAFS data did suggest that the structure of Ca2(UO2)(CO3)3(aq) contains the same (UO2)(CO3)34− unit as that in the solid Liebigite. Lastly, in the crystallographic analysis of Liebigite, identification of calcium proved to be difficult because the distance between uranium and calcium is probably identical to the distance between uranium and the terminal carbonate oxygen atom, rendering the analysis inconclusive. The calcium atoms in Ca2(UO2)(CO3)3(aq) may assume similar positions to those in the Liebigite mineral. Kalmikov and Choppin6 also used fluorescence spectroscopy to study the formation of Ca2(UO2)(CO3)3(aq) in aqueous solution at different ionic strengths (0.1 to 3.0 mol/kg NaClO4). They used the empirical relations of the Specif ic-ion Interaction Theory (S.I.T.)19 to correct the values of log β(213) as a function of the ionic strength and obtained the value log β0(213) = (29.8 ± 0.7) at infinite dilution. Speciation calculations with this value indicate that Ca2(UO2)(CO3)3(aq) is the predominant species at pH 8 in ground waters. Dong et al.8 used the anion exchange method to determine the values of log β of the ternary complexes of (UO2)(CO3)34− with Mg2+, Ca2+, Sr2+, and Ba2+. For the calcium system, values of log β0(113) = (27.18 ± 0.06) and log β0(213) = (30.70 ± 0.05) were obtained for the Ca(UO2)(CO3)32− and Ca2(UO2)(CO3)3(aq) complexes, respectively. For the magnesium system, only the value of log β0(113) = (26.11 ± 0.04) was obtained for Mg(UO2)(CO3)32−, while the Mg2(UO2)(CO3)3(aq) species was not identified. More recently, Geipel et al.15 used TRLFS to study the formation of the Mg complexes. In agreement with the previous observations by Dong,8 the formation constant of a single ternary complex, Mg(UO2)(CO3)32−, was calculated to be log β0(113) = (26.24 ± 0.13). Lee et al.10 also used TRLFS in combination with EDTA competitive titrations to determine the stability constants for Cam(UO2)(CO3)32(m−2) (m = 1, 2) with the values of log β0(113) = (27.27 ± 0.14) and log β0(213) = (29.81 ± 0.19). The values of the formation constants of the species M m (UO 2 )(CO 3 ) 3 2(m−2) (M = Ca, Mg), independently

aqueous solutions, it is fairly soluble (up to millimolar concentrations) in carbonate-buffered seawater (pH = 7.8− 8.4) due to complexation with carbonate to form very stable uranyl carbonato complexes bearing the general formula (UO2)(CO3)n2−2n (n = 1−3, Table 2).4 Historically, the (UO2)(CO3)34− anion in particular was considered to be the predominant uranium species in seawater. However, evidence from numerous studies5−11 indicates that (UO2)(CO3)34− can form soluble ternary complexes with divalent alkaline-earth cations, especially with Ca2+ and Mg2+, bearing the general formula Mm(UO2)(CO3)32m‑4 (M = Ca, Mg, m = 1, 2). Recently, Endrizzi and Rao11 proposed that these ternary species, although characterized by a relatively low stability (Table 2), could significantly affect the speciation of uranium in seawater due to the high concentrations of calcium and magnesium. Based on their thermodynamic investigations, they proposed that Ca2(UO2)(CO3)3(aq) and Mg(UO2)(CO3)32− are predominant in seawater, accounting for more than 90% of the overall uranium dissolved. Such a change in speciation would also impact the subsequent sorption behavior of uranium with the poly(amidoxime) sorbents intended for uranium recovery. Interestingly, these ternary species of calcium and magnesium also have a relatively high solubility similar to that of (UO2)(CO3)34−.11,12 Minerals with the general formula NajCakMgm(UO2)(CO3)3·xH2O(cr) (x = 6−18, j + 2k + 2m = 4) are known to be formed by concentration of aqueous solutions containing Ca2+, Mg2+, UO22+, and CO32− via slow evaporation.13,14 Although these minerals have been well characterized in the solid state, the currently available literature data on the solubility of these species show disagreement, and in some cases they need to be re-evaluated or updated. In order to fully understand the speciation of uranium in seawater, reliable solubility data should be coupled with the recent, more accurate data regarding its aqueous speciation. In the following sections, the current literature available on both the aqueous speciation and the solubility of ternary (Ca2+/Mg2+)-UO22+CO32− compounds and complexes are critically reviewed. In addition, to demonstrate the impact of these ternary complexes on uranium binding to amidoxime-type ligands under seawater conditions, calculated speciation diagrams are presented that compare the interactions of (UO2)(CO3)34− and Ca2(UO2)(CO3)3(aq) with glutaroimide-dioxime, a cyclic imidedioxime ligand that can be formed during synthesis of poly(amidoxime) sorbents.

■

AQUEOUS SPECIATION OF URANIUM UNDER SEAWATER CONDITIONS Several studies on the solution chemistry and the formation of soluble aqueous complexes of calcium and magnesium with (UO2)(CO3)34− have been published. Table 2 summarizes the thermodynamic parameters of relevant species. Of the studies reported in the literature, a majority of them were carried out using Time-Resolved Laser Fluorescence Spectroscopy (TRLFS), which proved to be a rather convenient, effective technique for studying the ternary (Ca2+/Mg2+)-UO22+-CO32− aqueous systems. For instance, (UO2)(CO3)34− does not show a distinct fluorescence emission because of the signal quenching due to the presence of dissolved carbonate.5,7,10,15 However, upon complexation with calcium or magnesium, a noticeable increase in the emission intensity occurs, indicating the formation of ternary (Ca2+/Mg2+)-UO22+-CO32− complexes.10 In the case of calcium, which forms up to two successive C

DOI: 10.1021/acs.iecr.5b03679 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Industrial & Engineering Chemistry Research

■

Article

SOLUBILITY OF TERNARY (CA2+/MG2+)-UO22+-CO32− COMPOUNDS Solid ternary compounds of Na, Ca, and Mg with UO2(CO3)34−, bearing the general formula NajCakMgm(UO2)(CO3)3·xH2O(cr) where x = 6−18 and j + 2k + 2m = 4, are known to be naturally occurring minerals.14,16 [The number of water molecules is x = 6, 10, 12, 18, and the stoichiometry of the alkaline or earth alkaline cations (j, k, m) satisfies the electroneutrality rule in the formula.] They form by slow precipitation from saline solutions with [Ca2+] ≈ 0.05 M and are usually identified by their efflorescent deposits on the walls of old mining sites having high rates of evaporation.16 Liebigite, Ca2(UO2)(CO3)3·10H2O(cr) [According to most current databases on minerals (see http://webmineral.com/data/) Liebigite has a formula bearing 11 water molecules, Ca2(UO2)(CO3)3·11H2O(cr). The original sources we took the data from16,17 report instead the formula Ca2(UO2)(CO3)3· 10H2O(cr). In the present review we use for Liebigite the formula including 10 water molecules, in agreement with the original literature sources.], and Andersonite, Na2Ca(UO2)(CO3)3·6H2O(cr), are the most common and also the first characterized species.12−14,16 Other members of the Liebigite group include Bayleyite, Mg2(UO2)(CO3)3·18H2O(cr), and Swartzite, CaMg(UO2)(CO3)3·12H2O(cr). Bachelet et al.12 synthesized Liebigite, for which they qualitatively determined its solubility in water in the temperature range 23−55 °C. The solubility of Liebigite was reported to be 5.6 × 10−3 mol dm−3 (1.34 g-U dm−3) at 25 °C. Alwan et al. studied the solubilities of Liebigite, Andersonite, Bayleyite, and Swartzite in aqueous solutions.16 Experiments were conducted at five different temperatures ranging from 0.0−25.0 °C in a CO2-free atmosphere using pH-unbuffered solutions containing no background electrolytes. After 1 week, the uranium concentration in the supernatant in equilibrium with the undissolved mineral was measured for each sample. [The authors reported that the system was at equilibrium after one week equilibration, by verifying that the concentration of U in the supernatant in equilibrium with the solid phase was constant over time.] Assuming a stoichiometric dissolution of a given mineral, the authors used the measured uranium concentration to calculate the concentrations of the other ionic components. Using the calculated solubility product constants and the values of the free energies of the individual ionic species taken from the literature,20 the authors calculated the standard free energies of formation, Δf G0, for the four mineral species. Values of Δf H0 were also estimated from the temperature dependence of log Ks0. In a more recent calorimetric study, Chernorukov et al.18 experimentally determined values of Δf H0 for Liebigite and Bayleyite that were in good agreement with the previous estimation by Alwan (see Table 2). More recently, Gorman-Lewis et al.17 reviewed the solubility of various uranium minerals, including the carbonate species studied by Alwan.16 Since the solubility products for Liebigite, Swartzite, Bayleyite, and Andersonite were not tabulated in the original paper by Alwan,18 Gorman-Lewis and co-workers used the values of the standard formation energies for the minerals from ref 16 (Table 2) to derive the solubility products. However, the solubility product constants calculated in this manner differ by 6 to nearly 7 orders of magnitude from those originally derived from solubility experiments by Alwan (Table 2). We found that this discrepancy results from the use of

determined in the studies reviewed above, agree quite well. Nonetheless, the experimental techniques used were limited mostly to fluorescence spectroscopy5−7,10,15 or anion exchange.8 Although those techniques proved to be suitable for the study of the ternary Ca/Mg−U-carbonate system, an independent confirmation by means of different methods should be obtained whenever possible. Endrizzi and Rao11 have recently published a study in which the formation of the Mm(UO2)(CO3)32(m−2) (M = Ca, Mg, m = 1, 2) species was studied using techniques different from those in previous studies. In detail, the formation of the Mg2+-UO22+-CO32− complexes was followed by direct UV−vis spectrophotometric titrations, while the formation of the Ca2+-UO22+-CO32− complexes was studied by potentiometric titrations employing a calcium-sensitive electrode to monitor the change in the free calcium concentration in solution accompanying the formation of the ternary complexes. The spectrophotometric determination of the magnesium species in solution was hampered by difficulties identifying the subtle spectral changes due to formation of the ternary species and by the use of low concentrations of UO22+ and Mg2+ (10−40 μM) in order to avoid the precipitation of magnesium hydroxide/carbonate in the alkaline range. Consequently, only a single Mg(UO2)(CO3)32− complex was identified via spectrophotometric titrations, with log β0(113) = (26.25 ± 0.04). Though the experimental conditions precluded the observation of the Mg2(UO2)(CO3)3(aq) complex in their study, it is likely a very weak complex that cannot be completely excluded.11 For the Ca2+-UO22+-CO32− system, the authors identified two successive Ca2+ complexes, Cam(UO2)(CO3)32(m‑2), with stability constants of log β0(113) = (27.00 ± 0.04) and log β0(213) = (30.84 ± 0.04) by potentiometry with the Ca2+-selective electrode. In addition, the enthalpies of complexation of these complexes were determined for the first time by means of titration microcalorimetry.11 The results indicate that the complexation of Ca2+ with UO2(CO3)34‑ is almost athermic, with the first stepwise reaction enthalpy being only slightly exothermic, ΔH0(1(13)),step. = −(7 ± 6) kJ/mol, while the second stepwise reaction enthalpy is essentially athermic, ΔH0(2(13)),step. = (0 ± 7) kJ/mol. The values of the stability constants obtained in this study11 are suggested in the present paper as recommended values and are used to calculate the distribution of the different U(VI) species in seawater in a later section. The selection of these values was reasonable in the case of the log β0(113), which is in agreement within ±0.2 log units of the values given in the different papers discussed above. However, the selection of a reliable log β0(213) value for the Ca2(UO2)(CO3)3(aq) species required more careful consideration since the reported stability constants clustered around two values, log β0(213) ∼ 29.8 or 30.8. After comparing the methodologies and the body of experimental data collected, we decided to select the value near 30.8 because this value was determined (within error) independently by three different techniques. On the other hand, the two studies reporting a value of 29.8 only used TRLFS and, moreover, one of the two studies used an incomplete speciation model for the Ca2+-UO22+-CO 32− system. In order to provide a better understanding of the chemistry of U(VI) in seawater, data concerning the aqueous speciation of U(VI) under seawater conditions should be used in combination with accurate solubility models describing the formation of the aforementioned NajCakMgm(UO2)(CO3)3· xH2O(cr) solid species, as discussed in the next section. D

DOI: 10.1021/acs.iecr.5b03679 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Article

Industrial & Engineering Chemistry Research

Using the solubility products calculated in this paper, the solubility of Liebigite is found to agree with the one previously determined by Bachelet et al.12 In addition, none of the (Ca2+/ Mg2+)-UO22+-CO32− minerals shown in Table 2 controls the concentration of uranium in seawater. In other words, uranium is completely soluble at its natural concentration levels and the distribution of U(VI) aqueous species in seawater conditions is consistent with the current knowledge about its seawater speciation discussed in the previous section (Table 2). Given the above-mentioned considerations, in this paper we propose to use, as better estimates, the values of the solubility products for Liebigite, Swartzite, Bayleyite, and Andersonite that were recalculated along with the corresponding values of Δf G0 that we derived (Table 2) for each mineral. However, this does not preclude the necessity of more accurate solubility studies of these ternary systems. Effect of the Complexation of Calcium and Magnesium with Tris(carbonato) Uranyl on the Speciation of Uranium in Seawater. To properly assess the impact of the formation of the ternary (Ca2+/Mg2+)-UO22+-CO32− species on uranium speciation in seawater (∼0.5 M NaCl), we used the empirical equations of the Specif ic-ion Interaction Theory19 to correct the values of the (Ca2+/Mg2+)-UO22+-CO32− minerals, in agreement with the OECD/NEA recommendations (values in Table 2).4 Endrizzi and Rao11 already used the same approach19 to convert the values of the formation constants at I ≈ 0 to those in 0.5 M NaCl. Table S4 in the Supporting Information lists the values of the specific-interaction parameters used. Calculations indicate that, because of the effect of the ionic strength of seawater, the solubility products of the solid species increase up to 4 orders of magnitude with respect to conditions of infinite dilution (see Table S2 in the Supporting Information). The ternary Ca2+/Mg2+−UO22+− CO32− aqueous species are conversely predicted to be less stable in a seawater-like medium than at I ≈ 0, by 2 to nearly 4 orders of magnitude (see Table 2). Figure 2 shows the relative distribution of the different uranium species as a function of pH in a seawater-like medium. The speciation model used includes the formation of all the aqueous and solid (Ca2+/Mg2+)-UO22+CO32− ternary species, and the related stability constants and solubility products corrected to account for the effect of the ionic strength, as described above. Despite the weakness of the ternary Ca2+/Mg2+ complexes with (UO2)(CO3)34−, due to the overwhelmingly high seawater concentrations of Ca2+ and Mg2+ (10 × 10−3 and 53 × 10−3 mol dm3−, respectively), the aqueous speciation of U(VI) is dominated by the formation of the M m (UO 2 )(CO 3 ) 3 2(m‑2) complexes. The neutral ternary Ca2(UO2)(CO3)3(aq) complex is predicted to be the most abundant species in solution, accounting for ca. 50−60% of the total uranium in solution (Figure 2). Conversely, if the formation of (Ca2+/Mg2+)-UO22+-CO32− complexes were neglected, U(VI) would exist quantitatively as (UO2)(CO3)34−. The distribution diagram also shows that the formation of the ternary (Ca2+/Mg2+)-UO22+-CO32− minerals does not occur in such environment. The formation of the auxiliary solid species calcite, CaCO3(cr), and magnesite, MgCO3(cr), was also accounted for in the distribution diagram. Thermodynamic calculations indicate that noteworthy amounts of these two minerals should form in seawater at pH higher than 8.0. Nonetheless, it is known that upper oceanic waters are generally supersaturated with respect to calcium and magnesium carbonates (whereas deeper waters are undersaturated with respect to these minerals).1 The slight decrease

different standard formation energies of individual ionic species to calculate the standard formation energies of the carbonate minerals: Alwan and co-workers18 used the standard formation energies from a previous database,20 while Gorman-Lewis et al.17 used more recently updated values, most likely those selected by the Nuclear Energy Agency (NEA) Thermochemical Data Base of 2003.4 Although the Δf G0 values for most ions involved in the formation of carbonate minerals are in good agreement between the two reference databases, the value of Δf G0 (UO2+ 2 ) is quite different, namely, 994.1 kJ/mol in ref 20 and 952.55 kJ/mol in ref 4 (see reference data in Table S3 in the Supporting Information). Such discrepancy is responsible for the marked difference in the solubility product constants reported in the two publications. After a careful review of the discrepancies in the solubility products of the (Ca2+/Mg2+)-UO22+-CO32− minerals in the literature, it is evident that the standard free energies of the (Ca2+/Mg2+)-UO22+-CO32− minerals in ref 16 need to be recalculated from the experimental solubility data in conjunction with the updated values for the individual ions in the recently critically reviewed database.4 In the present paper, we have digitized and accurately replotted the Arrhenius plot of Ksp from ref 18 in order to obtain the numerical values of the solubility products at 25.0 °C (Figure 1). From the best fits of

Figure 1. Arrhenius plot of Ksp for Ca−Mg−U-carbonate minerals. ○ Bayleyite; ● Liebigite; □ Swartzite; ■ Andersonite (Digitized and reproduced with permission from ref 18 Copyright 2009 Springer.).

log Ks vs 1/T, the values of the solubility products at 25.0 °C for the four carbonate minerals were obtained and are listed in Table 2. Furthermore, we made an additional calculation of the solubility products of these different minerals using a speciation model that includes the formation of the aqueous complexes Mm(UO2)(CO3)3(2m−4)(aq) (M = Ca or Mg, m = 0 − 2), as discussed in the previous section. These species were originally not included by Alwan and Williams16 in their model since their existence was not known at the time of that publication. Nonetheless, they need to be accounted for in the calculation of the solubility products since they significantly affect the aqueous speciation of U(VI). Using the recalculated constants in conjunction with the most up-to-date values of Δf G0 for each of the corresponding ionic species4 (Table S3), we were also able to calculate more reliable values of the corresponding standard formation energies of the (Ca2+/Mg2+)-UO22+-CO32− minerals. The results are summarized in Table 2, and detailed information is available in the Supporting Information. E

DOI: 10.1021/acs.iecr.5b03679 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Article

Industrial & Engineering Chemistry Research

Impact of Ca2(UO2)(CO3)3 and Mg(UO2)(CO3)32− Species on the Recovery of Uranium from Seawater. Rao and co-workers have shown that thermodynamic modeling of uranium sorption to amidoxime-based sorbents can help guide the design of more selective and efficient uranium sorbents.21 For example, they recently carried out thermodynamic investigations of the complexation of uranium with a series of three simple amidoxime-type ligands that can be formed during the synthesis of poly(amidoxime) sorbents: a cyclic imidedioxime ligand, an acyclic diamidoxime, and a cyclic amidoxime.22,23 Of the three ligands investigated in separate studies, they found that uranium formed the strongest complexes with glutaroimide-dioxime, the cyclic imidedioxime ligand, followed by the acyclic diamidoxime and the monosubstituted amidoxime, which was by far the weakest ligand. For the two strongest ligands, they used their measured stability constants of uranylamidoxime complexes in conjunction with the uranyl carbonate stability constants to generate a simple speciation model demonstrating the competition of either ligand with carbonate for U(VI).22 However, the speciation model in ref 22 assumes that the UO2(CO3)34− anion is the dominant U(VI) species in seawater; ternary (Ca2+/Mg2+)-UO22+-CO32− complexes were not included. To obtain a model that realistically reflects the U(VI) speciation in the presence of the amidoxime ligand under seawater conditions, we recalculated the U(VI) speciation by including the ternary (Ca2+/Mg2+)-UO22+CO32− complexes.24 As compared in Figure 3, the speciation diagrams obtained by excluding (a) or including (b) the ternary (Ca2+/Mg2+)-UO22+-CO32− complexes are significantly different. Clearly, the ternary species shows a significant effect on the ability of glutaroimide-dioxime to bind to uranium.

Figure 2. U(VI) speciation in seawater conditions calculated using the values of log β and log Ks (0.5 M NaCl) in Table 2. The values for the formation of the aqueous species (Ca2+/Mg2+)-UO22+-CO32− are taken from Endrizzi and Rao.11 The speciation model also includes the solubility products suggested in this paper (see Table 2). [Ca2+] = 0.0102 M; [Mg2+] = 0.053 M; [UO22+] = 1.4 × 10−8 M; [CO32−] = 0.024 M. 1 − UO22+, 2 − (UO2)(CO3) (aq), 3 − UO2OH+, 4 − (UO2)(CO3)22−, 5 − Ca2(UO2)(CO3)3(aq), 6 − Mg(UO2)(CO3)32‑, 7 − Ca(UO2)(CO3)32−, 8 − (UO2)(CO3)34−. Dotted gray line at pH 8.2 indicates the distribution of U at seawater pH.

■

in the distribution of Ca2(UO2)(CO3)3(aq) in Figure 2 at pH > 8 is a consequence of the concentration decrease of Ca2+ in solution, following the theoretically predicted formation of calcite (values of log Ks for the auxiliary species summarized in Table S2 in the Supporting Information).

CONCLUSION Literature studies of the solution equilibria of ternary (Ca2+/ Mg2+)-UO22+-CO32− complexes have been summarized and

Figure 3. Speciation diagrams simulating the relative distribution of the different uranium species as a function of pH, in seawater solutions in the presence of glutaroimide-dioxime (H2L, 0.150 M). The formation of the aqueous ternary species (Ca2+/Mg2+)-UO22+-CO32− is neglected in diagram (a) and accounted for in diagram (b). [UO22+]tot = 3.3 ppb, [CO32−]tot = 0.024 M, [Ca2+]tot = 0.010 M, [Mg2+]tot = 0.053 M, I = 0.5 M (NaCl); T = 25.0 °C. The values of log βn in Table 2 were used. The values of the formation constants of the complexes of glutaroimide-dioxime with U(VI) are taken from Rao et al.22 Labels in the diagrams: 1 − (UO2)(HL)20 (aq), 2 − (UO2)L0 (aq), 3 − (UO2)(HL)L−, 4 − (UO2)(L)22− 5 − Ca2(UO2)(CO3)3(aq), 6 − Mg(UO2)(CO3)32−, 7 − Ca(UO2)(CO3)32−, 8 − (UO2)(CO3)34−; the dotted gray line at pH = 8.2 indicates the U(VI) species distribution expected at seawater pH. F

DOI: 10.1021/acs.iecr.5b03679 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Article

Industrial & Engineering Chemistry Research critically reviewed herein. Based on the results of these studies, it is evident that the aqueous speciation of uranium in marine environments is dominated by the formation of the successive, stable complexes Mm(UO2)(CO3)3(2m−4) (aq) (M = Ca or Mg, m = 0 − 2). For the ternary species, the following values of log β0 at 298 K are selected by the present authors: log β0(113) = (27.00 ± 0.04), log β0(213) = (30.84 ± 0.04), and log β0(113) = (26.25 ± 0.04), for the two successive complexes with Ca2+, and for the 1:1 complex with Mg2+, respectively. These values are taken from the most recent results by Endrizzi and Rao,11 which are in good agreement with those previously obtained by Bernhard et al. by TRLFS7 and by Dong et al. using anionexchange.8 Using these data, the relative distribution of U(VI) species under seawater conditions (pH ≈ 8.2, 0.5 M NaCl; naturally occurring concentrations of Ca2+, Mg2+, UO22+, CO32−) was calculated by the authors. Speciation diagrams indicate that, on the contrary to the previously accepted speciation that shows the negatively charged (UO2)(CO3)34− is the dominant species of U(VI) in seawater, the neutral ternary complex, Ca2(UO2)(CO3)3(aq) accounts for more than 50% of the total U(VI) (Figure 2). This implies that new types of sorbents that could interact more strongly with the neutral complex may extract U(VI) more efficiently under seawater conditions. The solubility equilibria leading to the formation of the known ternary NajCakMgm(UO2)(CO3)3·xH2O(cr) solid species Liebigite, Swartzite, Bayleyite, and Andersonite were also critically reviewed. After careful analysis of the results in the previous papers followed by additional calculations, we obtained new, updated values for the solubility products (log K0s, 298 K) of these solids: log K0s = (33.9 ± 1.0), (33.4 ± 2.0), (31.7 ± 1.5) for Liebigite, Swartzite, and Bayleyite, respectively. The authors of this paper propose these values as better estimates with respect to the ones reported to date. Nonetheless, new solubility studies on these systems may be required in order to obtain even more reliable solubility data. Solubility data calculated in this paper, together with U(VI) speciation data, provide a comprehensive picture of the most relevant reactions involving U(VI) in a seawater environment. With respect to the prospective recovery of uranium from seawater using poly(amidoxime) collection systems, these data provide useful insights by helping to assess the affinity of U(VI) toward the amidoxime binding sites in conditions closely simulating the real systems. Among different amidoxime moieties, the cyclic glutaroimide-dioxime is currently considered as the best candidate for this task.22,23 The data in this paper, in combination with the speciation model for the formation of UO22+ complexes with glutaroimide-dioxime,22−24 indicate that the concentration of this ligand in seawater should be ca. 0.100 M, in order to effectively compete with the formation of Ca2(UO2)(CO3)3(aq) and Mg(UO2)(CO3)32−.

■

■

calculate the distribution of the U(VI) species (Figure 1) as a function of pH; Table S3, Δf G0 and related standard formation enthalpies (PDF)

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Present Addresses

# Kalrlsruhe Institute of Technology, Institute for Nuclear Waste Disposal (KIT − INE), Hermann-von-Helmholtz Platz, 1−76344 − Eggenstein Leopoldshafen, Germany. § 11545 Rockville Pike, TWFN-4B34, U.S. Nuclear Regulatory Commission, Rockville, MD 20852, USA.

Author Contributions

F.E. prepared the first version of the manuscript. All authors participated in revising the manuscript. All authors have given approval to the final version of the manuscript. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This work was supported by the Fuel Resources Program, Fuel Cycle Research and Development Program, Office of Nuclear Energy, the U.S. Department of Energy, under Contract No. DE-AC02-05CH11231 at Lawrence Berkeley National Laboratory (LBNL).

■

REFERENCES

(1) Chester, R.; Jickells, T. D. Marine Geochemistry, 3rd ed.; John Wiley & Sons: Hoboken, NJ, 2012. (2) Uranium 2014: Resources, Production and Demand; 2014. (3) Kim, J.; Tsouris, C.; Oyola, Y.; Janke, C. J.; Mayes, R. T.; Dai, S.; Gill, G.; Kuo, L.-J.; Wood, J.; Choe, K.-Y.; et al. Uptake of Uranium from Seawater by Amidoxime-Based Polymeric Adsorbent: Field Experiments, Modeling, and Updated Economic Assessment. Ind. Eng. Chem. Res. 2014, 53 (14), 6076. (4) Guillamont, R.; Fangänel, T.; Neck, V.; Fuger, J.; Palmer, D. A.; Grenthe, I.; Rand, M. Update on the Chemical Thermodynamics of Uranium, Neptunium, Plutonium, Americium and Technetium; OECD Nuclear Energy Agency, Data Bank: Issy-les-Moulineaux, France, 2003. (5) Bernhard, G.; Geipel, G.; Brendler, V.; Nitsche, H. Speciation of Uranium in Seepage Waters of a Mine Tailing Pile Studied by TimeResolved Laser-Induced Fluorescence Spectroscopy (TRLFS). Radiochim. Acta 1996, 74, 87. (6) Kalmykov, S. N.; Choppin, G. R. Mixed Ca2+/UO22+/CO32Complex Formation at Different Ionic Strengths. Radiochim. Acta 2000, 88 (9−11), 603. (7) Bernhard, G.; Geipel, G.; Reich, T.; Brendler, V.; Amayri, S.; Nitsche, H. Uranyl(VI) Carbonate Complex Formation: Validation of the Ca2UO2(CO3)3(aq.) Species. Radiochim. Acta 2001, 89 (8), 511. (8) Dong, W.; Brooks, S. C. Determination of the Formation Constants of Ternary Complexes of Uranyl and Carbonate with Alkaline Earth Metals (Mg2+, Ca2+, Sr2+, and Ba2+) Using Anion Exchange Method. Environ. Sci. Technol. 2006, 40 (15), 4689. (9) Doudou, S.; Arumugam, K.; Vaughan, D. J.; Livens, F. R.; Burton, N. A. Investigation of Ligand Exchange Reactions in Aqueous Uranyl Carbonate Complexes Using Computational Approaches. Phys. Chem. Chem. Phys. 2011, 13 (23), 11402. (10) Lee, J.-Y.; Yun, J.-I. Formation of Ternary CaUO2(CO3)32- and Ca2UO2(CO3)3(aq) Complexes under Neutral to Weakly Alkaline Conditions. Dalton Trans. 2013, 42 (27), 9862. (11) Endrizzi, F.; Rao, L. Chemical Speciation of Uranium(VI) in Marine Environments: Complexation of Calcium and Magnesium Ions

ASSOCIATED CONTENT

S Supporting Information *

. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.iecr.5b03679. Detailed information about the calculations of the solubility products of Liebigite, Swartzite, and Bayleyite; Table S1, calculated parameters used to fit the data points in Figure 1; Table S2, complete list of the reaction equilibria and the related stability constants or solubility products considered in the speciation model used to G

DOI: 10.1021/acs.iecr.5b03679 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX

Article

Industrial & Engineering Chemistry Research with [(UO2) (CO3)3]4- and the Effect on the Extraction of Uranium from Seawater. Chem. - Eur. J. 2014, 20 (44), 14499. (12) Bachelet, M.; Cheylan, E.; Douis, M.; Goulette, J. C. Preparation and Properties of the Uranyl Carbonates. II. The Alkaline Earth Series. Bull. Soc. Chim. Fr. 1952, 565. (13) Vochten, R.; Van Haverbeke, L.; Van Springel, K. Synthesis of Liebigite and Andersonite, and Study of Their Thermal Behavior and Luminescence. Can. Mineral. 1993, 31 (1), 167. (14) Vochten, R.; Van Haverbeke, L.; Van Springel, K.; Blaton, N.; Peeters, O. M. The Structure and Physicochemical Characteristics of a Synthetic Phase Compositionally Intermediate between Liebigite and Andersonite. Can. Mineral. 1994, 32 (3), 553. (15) Geipel, G.; Amayri, S.; Bernhard, G. Mixed Complexes of Alkaline Earth Uranyl Carbonates: A Laser-Induced Time-Resolved Fluorescence Spectroscopic Study. Spectrochim. Acta, Part A 2008, 71A (1), 53. (16) Alwan, A. K.; Williams, P. A. The Aqueous Chemistry of Uranium Minerals. 2. Minerals of the Liebigite Group. Mineral. Mag. 1980, 43 (329), 665. (17) Gorman-Lewis, D.; Burns, P. C.; Fein, J. B. Review of Uranyl Mineral Solubility Measurements. J. Chem. Thermodyn. 2008, 40 (3), 335. (18) Chernorukov, N. G.; Knyazev, A. V.; Vlasova, E. V.; Kuznetsova, N. Y. A Physicochemical Study of Alkaline-Earth Metal Carbonatouranylates. Radiochemistry (Moscow, Russ. Fed.) 2009, 51 (3), 244. (19) Ciavatta, L. The Specific Interaction Theory in Evaluating Ionic Equilibriums. Ann. Chim. 1980, 70, 551. (20) Barner, H. E.; Scheuerman, R. V. Handbook of Thermochemical Data For Compounds and Aqueous Species; Wiley: 1978. (21) Rao, L. Application of Radiation Grafting: Progress and Status of the Extraction of Uranium from Seawater in Japan. J. Isotopes 2012, 25 (3), 129. (22) Tian, G.; Teat, S. J.; Zhang, Z.; Rao, L. Sequestering Uranium from Seawater: Binding Strength and Modes of Uranyl Complexes with Glutarimidedioxime. Dalton Trans. 2012, 41 (38), 11579. (23) Tian, G.; Teat, S. J.; Rao, L. Thermodynamic studies of U(VI) complexation with glutardiamidoxime for sequestration of uranium from seawater. Dalton Trans. 2013, 42 (16), 5690−5696. (24) Endrizzi, F.; Melchior, A.; Tolazzi, M.; Rao, L. Complexation of Uranium(VI) with Glutarimidoxioxime: Thermodynamic and Computational Studies. Dalton Trans. 2015, 44, 13835−13844.

H

DOI: 10.1021/acs.iecr.5b03679 Ind. Eng. Chem. Res. XXXX, XXX, XXX−XXX