Uptake Mechanisms of Eu(III) on Hydroxyapatite: A Potential

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Uptake Mechanisms of Eu(III) on Hydroxyapatite: A Potential Permeable Reactive Barrier Backfill Material for Trapping Trivalent Minor Actinides Lin Xu,†,‡ Tao Zheng,†,‡ Shitong Yang,*,†,‡ Linjuan Zhang,§ Jianqiang Wang,§ Wei Liu,†,‡ Lanhua Chen,†,‡ Juan Diwu,†,‡ Zhifang Chai,†,‡ and Shuao Wang*,†,‡ †

School for Radiological and Interdisciplinary Sciences (RAD-X), Soochow University, 215123 Suzhou, P. R. China Collaborative Innovation Center of Radiation Medicine of Jiangsu Higher Education Institutions, 215123 Suzhou, P. R. China § Shanghai Institute of Applied Physics and Key Laboratory of Nuclear Radiation and Nuclear Energy Technology, Chinese Academy of Sciences, 201800 Shanghai, P. R. China ‡

S Supporting Information *

ABSTRACT: The permeable reactive barrier (PRB) technique has attracted an increasing level of attention for the in situ remediation of contaminated groundwater. In this study, the macroscopic uptake behaviors and microscopic speciation of Eu(III) on hydroxyapatite (HAP) were investigated by a combination of theoretical modeling, batch experiments, powder X-ray diffraction (PXRD) fitting, and X-ray absorption spectroscopy (XAS). The underlying removal mechanisms were identified to further assess the application potential of HAP as an effective PRB backfill material. The macroscopic analysis revealed that nearly all dissolved Eu(III) in solution was removed at pH 6.5 within an extremely short reaction time of 5 min. In addition, the thermodynamic calculations, desorption experiments, and PXRD and XAS analyses definitely confirmed the formation of the EuPO4·H2O(s) phase during the process of uptake of dissolved Eu(III) by HAP via the dissolution−precipitation mechanism. A detailed comparison of the present experimental findings and related HAP−metal systems suggests that the relative contribution of precipitation to the total Eu(III) removal increases as the P:Eu ratio decreases. The dosage of HAP-based PRB for the remediation of groundwater polluted by Eu(III) and analogous trivalent actinides [e.g., Am(III) and Cm(III)] should be strictly controlled depending on the dissolved Eu(III) concentration to obtain an optimal P:M (M represents Eu, Am, or Cm) ratio and treatment efficiency.

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INTRODUCTION The quality of surface water and groundwater resources is threatened by direct exposure to radionuclides (e.g., 60Co, 154 Eu, 232Th, 235U, 235Np, 239Pu, 241Am, 247Cm, etc.) leaching from the medical diagnosis, nuclear weapon production, and nuclear materials industries.1,2 These radionuclides can be ingested by aquatic organisms and humans via biological accumulation in the food chain. Their high and long-term radioactivity as well as strong complexation affinity for organic ligands can induce radiation and biochemical injuries to organs.3 Excess exposure to 241Am and 244Cm can trigger acute radiotoxicity (e.g., bone marrow damage) and chronic diseases (e.g., incisor malformation, thyroiditis, nephritis, hepatitis, and various cancers) in marine animals (e.g., crucian carp) and terrestrial mammals (e.g., mice, rats, dogs, and humans).4,5 Consequently, practical remediation strategies for radioactive contamination in aquatic environments are urgently needed. A variety of engineering technologies, such as irrigation, electrodynamics, permeable reactive barriers (PRBs), pumpand-treat, bioremediation, and monitored natural attenuation, © XXXX American Chemical Society

have been developed for the purification of contaminated groundwater. Among these, PRB technology is widely used for the remediation of contaminated groundwater because of its diverse advantages, including dependence on natural groundwater flow rather than active pumping, a weak requirement for monitoring and maintenance, utilization of cost-effective backfill materials, and favorable disposal capacity for multiple contamination plumes.6−8 A series of in situ studies at field scale have been conducted to assess the removal performance, treatment cost, and long-term effectiveness of this remediation approach.9−13 For instance, the U.S. Environmental Protection Agency constructed three PRB facilities at the Fry Canyon site to purify U(VI)-contaminated groundwater with a dissolved U(VI) concentration of ∼85 μM. Three types of reactive materials, i.e., pelletized hydroxyapatite [Ca5(PO4)3OH] (HAP) in the form of bone charcoal, pelletized metallic iron, Received: December 2, 2015 Revised: March 9, 2016 Accepted: March 10, 2016

A

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the standard hexagonal EuPO4·xH2O solid (JCPDS Card No. 20-1044). The number of H2O molecules in the crystal structure was calculated as ∼1.0 based on the weight loss as shown in the thermogravimetric curve (Figure S4). The specific analysis procedure is described in the Supporting Information. Equilibrium Modeling. To predict the feasibility of using a HAP sample for Eu(III) removal, the theoretical dissolution behaviors of HAP in the presence of Eu(III) in solution and the potential precipitation of europium-containing phases were simulated using Visual MINTEQ version 3.1.20 Specifically, the simulation was performed at molar P:Eu ratios [i.e., the ratio between the number of moles of total P in a specific amount of HAP and those of Eu(III) at a certain Eu(NO3)3·6H2O concentration] of 60:1 [i.e., 0.5 g/L HAP and 5 × 10−5 mol/L Eu(III)], 18:1 [i.e., 0.5 g/L HAP and 1.67 × 10−4 mol/L Eu(III)], and 3.6:1 [i.e., 0.1 g/L HAP and 1.67 × 10−4 mol/L Eu(III)]. The specific concentrations of Eu(III) and NaNO3 (0.01 mol/L) were input in the simulation window of Visual MINTEQ. NaNO3 was used as the background electrolyte due to the common distribution of the Na+ ion in the natural environment and the widespread presence of the NO3− ion in contaminated water.22 The concentration of NaNO3 was maintained at 0.01 mol/L to simulate the conventional ionic strength in aquatic systems.23,24 A pressure value of 0.00038 atm was also input to simulate the partial pressure of dissolved carbon dioxide in aquatic systems. The thermodynamic calculation was based on the theoretical dissolution−precipitation equilibrium of the HAP material and the potential europium-containing solids as well as their solubility product constants [i.e., 10−44.33 for HAP, 1015.49 for Eu(OH)3(s), 10−32.3 for Eu2(CO3)3(s), and 10−25.96 for europium phosphate] embedded in the software. The theoretical amounts of dissolved HAP and the resulting europium-containing phases were simulated in the pH range of 1.0−10.0. More detailed information about the equilibrium modeling is described in the Supporting Information. HAP and EuPO4·H2O Dissolution Experiments. The dissolution behaviors of the synthetic HAP as a function of the solution pH and aging time were explored to evaluate the capacity of the HAP as a source of P. For the time-dependent dissolution experiment, 0.25 g of the HAP sample was added to 500 mL of a NaNO3 solution (0.01 mol/L) to achieve a solid:liquid ratio of 0.5 g/L. The pH of the mixture was adjusted to 6.5 with aliquots of HNO3 and/or NaOH solutions. The pH adjustment resulted in a deviation of the solid:liquid ratio of HAP of 7.0. Similar trends are observed for P:Eu ratios of 18:1 and 3.6:1, with the complete dissolution of HAP in a wider pH range of 1.0−5.8. As illustrated in Figure 1, the solid europium phosphate is present in the HAP−Eu(III) systems along with the dissolution of HAP. Specifically, the theoretical molar concentration of this phase rapidly increases as the solution pH increases from 1.0 to 2.0 and then remains nearly identical to that of the total Eu(III) at pH >2.0. Correspondingly, the concentration of the dissolved Eu(III) species sharply decreases to near zero as the pH increases from 1.0 to 2.0 and then remains at this extremely low value at pH >2.0 (Figure S5). The modeling results also reveal that the presence of Eu(III) partly alters the dissolution behaviors of HAP in solution (Figure S6A,B). Specifically, at pH >6.0, the dissolution of HAP alone cannot provide a sufficient source of P for the precipitation of the total dissolved Eu(III). The addition of Eu(III) disrupts the solubility equilibrium of HAP and provides a thermodynamic gradient for the enhanced dissolution of HAP followed by the precipitation of solid europium phosphate. The equilibrium modeling suggests that the HAP material may be suitable for the purification of Eu(III)-bearing groundwater. Specifically, the aqueous P resulting from the HAP dissolution would interact with the dissolved Eu(III), leading to the precipitation of solid europium phosphate as described by the following two equations: Ca10(PO4 )6 (OH)2 + 2H+ ⇔ Ca 2 + + H nPO4 n − 3 + 2H 2O (1)

Eu

3+

+ HnPO4

n−3

+ x H 2O ⇔ EuPO4 ·x H 2O(s) + nH

+

(2)

However, the theoretical simulation cannot predict the reaction rate of the HAP−Eu(III) systems. Moreover, the thermodynamic calculations based solely on the solubility product constants of different solids neglect the potential retention of C

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× 10−5 mol/L. In addition, nearly all of the dissolved Eu(III) in the solution was removed by HAP at pH >3.0 at the three P:Eu ratios. These two experimental phenomena contradict the conventional uptake trends of metal ions at the solid−water interface via adsorption mechanisms; i.e., the sorption behaviors are greatly dependent on the solution pH, and the sorption percentage is expected to be lower at higher metal concentrations. Taken together with the theoretical simulation results as shown in Figure 1, one can speculate that the uptake of Eu(III) by HAP is due to precipitation rather than adsorption. As illustrated in Figure S7, even the lowest concentrations of dissolved P (∼3.0 × 10−4 mol/L) and Eu(III) (5.0 × 10−5 mol/L) are supersaturated to form the EuPO4·H2O(s) solid (pKsp = 26).34 Theoretically, the interaction process and trends for the HAP−Eu(III) systems are mainly controlled by the relative thermodynamic stabilities of the HAP and EuPO4·H2O(s) phases. As demonstrated by the dissolution experiments (Figure S7), the thermodynamic stability of EuPO4·H2O(s) is much higher than that of HAP. Consequently, the dissolved Eu(III) may be coordinated by the phosphate leaching from the dissolution of HAP, resulting in the formation of the stable EuPO4·H2O(s) phase. The equilibrium time is commonly regarded as an important factor for evaluating the application potential of a material during water purification. According to the batch kinetic experiment (Figure S9), ∼100% of the initially added Eu(III) was removed within a short reaction time of 5 min. In this system, the fast uptake kinetics can be attributed to the dissolution−precipitation mechanism as indicated above. The dissolution kinetic curve of HAP (Figure S8) indicates that the amount of P leached in the first few minutes (∼2.0 × 10−4 mol/ L) is sufficient to completely precipitate the dissolved Eu(III) (5.0 × 10−5 mol/L). The precipitation of solid europium phosphate is a rapid process, and thus, fast uptake kinetics is observed for the HAP−Eu(III) system. PXRD Analysis. To verify the proposed dissolution− precipitation mechanism, the PXRD patterns of Eu(III)-loaded HAP samples were collected and analyzed in detail. For the pHdependent HAP−Eu(III) samples at a P:Eu ratio of 60:1, there was no clear evidence to indicate the formation of a new phase (Figure 3 and Figure S10). As mentioned above, formation of

Eu(III) by HAP via surface complexation mechanisms, as observed for other metal ions [e.g., Co(II), Zn(II), and U(VI)].25−27 Therefore, additional experiments are needed to verify and supplement this theoretical deduction. HAP and EuPO4·H2O Dissolution Results. The dissolution of the synthetic HAP and EuPO4·H2O as a function of solution pH is presented in Figure S7. The amount of P leached from the HAP sharply decreased from ∼1.9 × 10−3 to ∼3.0 × 10−4 mol/L as the pH increased from 1.5 to 5.0 and then remained constant at pH >5.0. Similarly, the amount of dissolved P from EuPO4·H2O decreased from ∼5.0 × 10−4 to ∼5.0 × 10−5 mol/L as the solution pH increased from 1.5 to 3.0 and then remained nearly unchanged at pH >3.0. Thus, the amount of dissolved P from EuPO4·H2O was much lower than that from HAP throughout the whole pH range. This phenomenon suggests that the EuPO4·H2O phase possesses a thermodynamic stability higher than that of the HAP solid. As illustrated in the dissolution kinetic curve (Figure S8), the amount of P dissolved from HAP at pH 6.5 and 293 K gradually increased with a prolonged aging time and reached an equilibrium value of ∼3.0 × 10−4 mol/L after 1440 min (i.e., 24 h). Under the assumption of congruent dissolution, the experimental solubility of the synthetic HAP sample at 293 K was determined to be ∼0.005 g/100 g of water based on the amount of P (∼3.0 × 10−4 mol/L) leached at the equilibrium time. This value is higher than those of the HAP minerals used in the equilibrium modeling (i.e., ∼0.0025 g/100 g of water) and in a previous dissolution study (∼0.0034 g/100 g of water)32 but much lower than that of the apatite material (0.0167 g/100 g of water) used for the removal of U(VI).33 These differences may be attributed to the diverse properties of the applied HAP samples (e.g., solubility product constant, specific components, particle size, etc.). Batch Uptake Data. Figure 2 presents the pH dependence of Eu(III) uptake by HAP at three different P:Eu ratios. The

Figure 2. Uptake percentages of Eu(III) by HAP and the remaining concentrations of Eu(III) in solution at different P:Eu ratios. T = 293 K, and CNaNO3 = 0.01 mol/L.

removal percentage of Eu(III) exhibited a sharply increasing trend as the pH increased from 1.5 to 3.0. Correspondingly, the concentration of the remaining Eu(III) in solution rapidly decreased with increasing pH, in good agreement with the equilibrium modeling results (Figure S5). Unexpectedly, the removal curve for a P:Eu ratio of 18:1 at a higher Eu(III) concentration of 1.67 × 10−4 mol/L was much higher than that for a P:Eu ratio of 60:1 at a lower Eu(III) concentration of 5.0

Figure 3. PXRD patterns of HAP−Eu(III) samples prepared at different P:Eu ratios. T = 293 K; pH = 6.5; CNaNO3 = 0.01 mol/L. D

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Figure 4. XAS data and the fit results of Eu(III) reference and uptake samples at different P:Eu ratios. (A) Solid lines represent the experimental k3weighted spectra, and dashed lines represent the fit results. (B) Symbols represent the RSFs. and solid lines represent the fit results.

the EuPO4·H2O(s) phase is expected because of its low solubility product (pKsp = 26) and the relatively high concentrations of dissolved P and Eu(III) in solution. However, this precipitate may not be detected by the PXRD technique because of its relatively low weight percentage relative to the entire bulk phase. For the HAP−Eu(III) samples prepared at pH 6.5 with P:Eu ratios of 18:1 and 3.6:1, the EuPO4·H2O(s) phase can be roughly identified on the basis of the gradually appearing diffraction peak at 20° (marked by the asterisks) relative to that of the pure HAP (Figure 3). Other peaks related to this phase (e.g., 26° and 32° as marked by the solid rhombus) may be masked by those of the HAP sample. These analyses preliminarily suggest that the interaction of Eu(III) with HAP leads to the formation of the EuPO4·H2O(s) phase at pH 6.5. For the samples with a P:Eu ratio of 3.6:1, the EuPO4· H2O(s) phase was clearly detected over a wide pH range, from an extremely acidic pH of 3.0 to a highly alkaline pH of 9.8 (Figure S11). An additional diffraction peak at 38.6° (marked by the hollow triangle) was observed in the PXRD pattern of the sample prepared at pH 9.8, suggesting the possible formation of a secondary solid phase. The FTIR analysis revealed some carbonate groups in the HAP structure (Figure S3). In addition, dissolved atmospheric carbon dioxide in the solution would be present as carbonate and/or bicarbonate ions at alkaline pH. Hence, the Na3Eu(CO3)3(s) phase is likely to form via the precipitation of Na+ ions and Eu(III) species with the carbonate ions therein. However, this phase was not identified in the equilibrium modeling, even though it was specified as a possible solid during the simulation. This discrepancy may result from the idealization of the computer modeling (based on the solubility product constants of different solids) and the complexity of the actual sorption systems.

To further verify the presence of EuPO4·H2O(s) and determine its relative proportion, the PXRD patterns of the uptake samples were simulated using TOPAS version 4.2.35 By adopting a least-squares minimization approach, we performed the fitting on the basis of the theoretical PXRD patterns of HAP and EuPO4·H2O(s) as refined from their standard crystalline structures.36 Optimized fitting results were achieved by scaling the collected PXPD patterns with those of the standards [herein, HAP and EuPO4·H2O(s)] on an equalintensity basis with the minimum deviation.37,38 The fitting results revealed that the sample with a P:Eu ratio of 60:1 was composed of a large mass fraction (∼99.9%) of HAP and a small mass fraction (∼0.1%) of EuPO4·H2O(s) (Figure S12). Specifically, for the samples with P:Eu ratios of 18:1 and 3.6:1, the mass fractions of EuPO4·H2O(s) were fitted to be ∼19 and ∼56%, respectively. According to these PXRD fitting results, the EuPO4·H2O(s) phase formed in the HAP−Eu(III) systems, and its mass fraction increased with a decreasing P:Eu ratio. As mentioned above, some diffraction peaks of the EuPO4·H2O(s) phase (e.g., 26° and 32° as marked by the solid rhombus in Figure 3) overlapped with those of the HAP sample. Considering the poor crystallinity of the HAP substrate and the high crystallinity of the refined EuPO4·H2O(s) model, the mass fractions of EuPO4·H2O(s) derived from the PXRD fitting are somewhat higher than the actual values, particularly for the P:Eu ratios of 18:1 and 3.6:1. Hence, additional analytical approaches are required to quantitatively determine the exact amounts of EuPO4·H2O(s) under these conditions. Desorption Experiments. According to the preliminary experiment, the EuPO4·H2O(s) phase was nearly insoluble in the EDTA-2Na solution at pH 6.5 due to the insignificant amounts of dissolved P and Eu(III). Thus, the EDTA-2Na solution can extract adsorbed Eu(III) only on HAP surfaces. E

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demonstrated that the central Eu atoms were coordinated with an average of ∼8.4 O at REu−O ∼ 2.43 Å, ∼1.3 P at REu−P1 ∼ 3.22 Å, ∼2.5 P at REu−P2 ∼ 3.67 Å, and ∼2.1 Eu at REu−Eu ∼ 4.08 Å. These structural parameters are consistent with those of the reference EuPO4·H2O(s) sample. In addition, the XAS fit results are also consistent with the crystallographic data for SmPO4·0.667H2O (Figure S14), whose crystal structure is extremely similar to that of EuPO4·H2O.39,40 These results clearly confirm the formation of the EuPO4·H2O(s) phase in our HAP−Eu(III) systems. This finding is consistent with the equilibrium modeling and PXRD fitting. However, the XAS analysis results provide no evidence of some of the adsorbed species indicated by the desorption experiments. XAS fitting revealed the average coordination environment surrounding the central Eu atoms. The masking effect of the strong oscillation signals of the precipitated species hindered the detection of the weak XAS oscillation signals of the adsorbed species. Hence, advanced fitting approaches (e.g., linear fitting and principal component analysis) should be adopted in future studies to distinguish the adsorbed and precipitated species. Comparison with Other HAP−Metal Systems. A series of studies has been performed to explore the underlying removal mechanisms of heavy metal ions by HAP with the aid of multiple experimental techniques.11,26,41−45 As listed in Table S4, the uptake of Zn(II) by HAP at P:Zn ratios of 24:1 and 6:1 was driven by surface complexation over a wide pH range (5.0−9.0).26,41 By contrast, the hopeite [Zn3(PO4)2· 4H2O(s)] phase was formed at a lower P:Zn ratio of 2:1 via the dissolution−precipitation mechanism.26 Similarly, the interaction of Pb(II) with HAP at low P:Pb ratios (e.g., 5.3:1 and 3.4:1) led to the dissolution of HAP and precipitation of the hydroxypyromorphite [Pb10(PO4)6(OH)2] phase.42 Previous synchrotron XRD, SEM-EDS, and XAS analyses have indicated that U(VI) is adsorbed on HAP via surface complexation at a relatively higher P:U ratio of 303:1.11 A similar adsorption mechanism was also proposed for the HAP−Np(V) and HAP− Pu(VI) systems over a wide pH range (e.g., 6.0−11.0) at relatively high P:Np (6000:1 and 96:1) and P:Pu (15000:1) ratios.43−45 Interestingly, for the HAP−U(VI) systems, a decrease in the P:U ratio from 303:1 to 132:1 resulted in a decrease in the proportion of surface complexation from 100 to 63% along with an increase in the proportion of chernikovite [H2(UO2)2(PO4)2·10H2O] precipitation from 0 to 37%. Finally, the removal of U(VI) was completely induced by the precipitation of the chernikovite phase at a lower P:U ratio of 18:1. In the study presented here, the decrease in the P:Eu ratio from 60:1 to 3.6:1 suppressed the occurrence of surface complexation and facilitated the formation of EuPO4·H2O(s). By comparison and integration of the research findings as mentioned above, the relationship between the specific experimental conditions and the underlying removal mechanisms can be tentatively summarized. In brief, the reaction process and trends for the HAP−metal systems are closely related to the molar P:M ratio (M represents the metal ions). Specifically, the metal ions are expected to be sequestered on the HAP surfaces via surface complexation at higher P:M ratios. Alternatively, the dissolved metal ions tend to be immobilized via the dissolution of HAP and the precipitation of phosphate solids at lower P:M ratios. The marginal P:M values for these two removal mechanisms will vary for different HAP−metal systems. When used as a PRB backfill material for the purification of contaminated groundwater, the dosage of HAP

For the HAP−Eu(III) sample prepared at a P:Eu ratio of 60:1 and pH 6.5, the concentration of Eu(III) in solution after EDTA-2Na extraction was only ∼23% of the initial concentration, suggesting that ∼23% of the retained Eu(III) was due to surface complexation and/or incorporation mechanisms. Consequently, the remaining ∼77% of the sequestered Eu(III) was ascribed to the precipitation of EuPO4·H2O(s). This result is somewhat different from the results of the equilibrium calculation, which indicated that nearly all of the dissolved Eu(III) would form the europium phosphate phase at pH 6.5 (Figure 1). This slight difference may reflect the limitations of the theoretical modeling and the differences between the HAP materials used in these two systems. In the cases of P:Eu ratios of 18:1 and 3.6:1 at pH 6.5, an extremely small amount of Eu(III) was detected after the desorption reaction, suggesting that nearly 100% of the dissolved Eu(III) was transformed to EuPO4·H2O(s) during the uptake reactions and that this phase was stable in the EDTA-2Na solution. According to the desorption experiments, the weight percentages of the EuPO4·H2O(s) formed relative to the bulk phase increased from ∼9 to ∼33% as the P:Eu ratio decreased from 18:1 to 3.6:1. This variation trend is consistent with that obtained from the PXRD fitting. For the sample prepared at pH 9.8 with a P:Eu ratio of 3.6:1, the diffraction peak belonging to the Na3Eu(CO3)3(s) phase (38.6°) completely disappeared after desorption (Figure S13). According to the quantitative measurement results, the formation of this unstable precipitate phase consumed ∼5% of the total Eu(III) in solution. In other words, most of the dissolved Eu(III) (∼95%) tended to form a stable EuPO4· H2O(s) phase. In brief, the desorption experiments confirmed the formation of EuPO4·H2O(s) at lower P:Eu ratios (e.g., 18:1 and 3.6:1). However, the microscopic species of Eu(III) at the higher P:Eu ratio (i.e., 60:1) could not be definitively identified from the PXRD analysis. Therefore, an additional spectroscopic approach, such as XAS, is needed to provide more accurate and detailed information. XAS Analysis. The k3-weighted XAS spectra of the Eu(III) reference and uptake samples at different P:Eu ratios and pH values are presented in Figure 4A. The spectra of the three uptake samples are extremely similar to that of the reference EuPO4·H2O(s) sample but differ from that of Eu(OH)3(s). This phenomenon suggests the potential formation of a EuPO4· H2O(s) phase rather than a Eu(OH)3(s) precipitate. Fourier transformation was performed to obtain the corresponding radial structural functions (RSFs) for the typical oscillation features in the k3-weighted XAS spectra. As shown in Figure 4B, the RSFs for the uptake samples exhibited a single peak with a high intensity at ∼1.90 Å (uncorrected for the phase shift) arising from the signal of the O atoms in the first coordination shell. Two additional peaks appeared over the R ranges of 2.50−3.30 and 3.30−4.40 Å, suggesting higher coordination shells (e.g., Eu−P and/or Eu−Eu). To further identify the underlying interaction mechanisms between Eu(III) and HAP, the k3-weighted spectra and RSFs for the Eu(III)-containing samples were fit with a least-squares approach; the parameters obtained are listed in Table S3. The displayed interatomic distances (i.e., REu−O, REu−P, and REu−Eu) were corrected for the phase shift during the XAS data fitting. Therefore, the actual coordination environment of Eu(III) in the uptake samples can be confirmed directly from their values. For the three HAP−Eu(III) uptake samples prepared at different P:Eu ratios and pH values, the spectral analysis F

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dissolution kinetic data of the synthetic HAP at pH 6.5, kinetic data of uptake of Eu(III) by HAP at pH 6.5, PXRD patterns of uptake samples at different pH values with P:Eu ratios of 60:1 and 3.6:1, PXRD fitting results of the HAP−Eu(III) samples with different P:Eu ratios, PXRD patterns of the uptake samples with a P:Eu ratio of 3.6:1 before and after soaking in the EDTA-2Na solution, crystallographic data of SmPO4·0.667H2O, and comparison of the uptake mechanisms for different HAP−metal systems (PDF)

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ENVIRONMENTAL IMPLICATIONS The complementary thermodynamic modeling, macroscopic data, and PXRD and XAS analyses in this study clearly demonstrate that dissolved Eu(III) can be effectively immobilized by HAP, resulting in the formation of a stable EuPO4·H2O(s) phase. Generally, the removal capacity of reactive materials for metal ions via the precipitation mechanism is higher than that via the adsorption mechanism. Although a small fraction of the dissolved Eu(III) was directly adsorbed on the HAP surfaces at a higher P:Eu ratio of 60:1, the HAP material exhibited favorable removal performance toward Eu(III) over a wide range of environmental conditions (i.e., molar P:Eu ratio ranging from 18:1 to 3.6:1; solution pH ranging from 3.0 to 9.8). These findings demonstrate the feasibility of using HAP-based PRBs for the in situ purification of groundwater containing trivalent lanthanide/actinides (e.g., Eu, Am, and Cm). Given the complexity of real groundwater systems, some significant issues remain to be clarified prior to the practical operation of HAP-based PRBs. Previous studies have proposed that the interaction mechanisms for HAP−metal systems are dependent on whether the HAP suspensions are allowed to pre-equilibrate (i.e., achieve saturation) before the addition of metal ions.26,46 This finding emphasizes the need to explore the reaction mechanisms of Eu(III) with pre-equilibrated HAP suspensions. A systematic comparison of nonequilibrated and pre-equilibrated systems may provide useful information for designing appropriate strategies for PRB applications. In addition, the role of carbonate/bicarbonate ions, which are present at much higher concentrations in the aquatic environment than under the experimental conditions used herein, requires more research. As summarized from previous studies,11,26,41−46 the reaction process and trends for HAP− metal systems are closely related to the molar P:M ratio. For HAP materials obtained using different approaches, the solubility values and capacity for providing a source of P obviously differ because of their diverse properties (e.g., solubility product constant, specific components, particle size, etc.).35,36,47 More studies are needed to seek the optimal P:M (M represents Eu, Am, or Cm) ratio to facilitate the dissolution−precipitation reaction. Moreover, powdered HAP samples with a fine particle size and low porosity would hinder groundwater flow passing through the PRB column. Consequently, granular or pelletized HAP materials with excellent porosity and permeability are currently being synthesized in our lab. Additional laboratory experiments and column tests will be performed to further evaluate the effectiveness of HAP materials for trapping trivalent Eu(III), Am(III), and Cm(III) from contaminated groundwater.

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AUTHOR INFORMATION

Corresponding Authors

*E-mail: [email protected]. Telephone: +86-51265883945. Fax: +86-512-65883945. *E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Financial support from the National Natural Science Foundation of China (41203086, 21422704, 41303006, and U1532259), the Science Foundation of Jiangsu Province (BK20140007), the “Young Thousand Talented Program” in China, the Jiangsu Provincial Key Laboratory of Radiation Medicine and Protection, and the Priority Academic Program Development (PAPD) of Jiangsu Higher Education Institutions is acknowledged.

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REFERENCES

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ASSOCIATED CONTENT

S Supporting Information *

The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/acs.est.5b05932. Synthetic and characterization of HAP and EuPO4·H2O samples, equilibrium modeling procedure, XAS data collection and analysis, speciation of Eu(III) in solution as a function of pH, dissolution behaviors of the synthetic HAP and EuPO4·H2O as a function of solution pH, G

DOI: 10.1021/acs.est.5b05932 Environ. Sci. Technol. XXXX, XXX, XXX−XXX

Article

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DOI: 10.1021/acs.est.5b05932 Environ. Sci. Technol. XXXX, XXX, XXX−XXX