Article Cite This: Environ. Sci. Technol. 2018, 52, 5902−5910
pubs.acs.org/est
Incorporation Modes of Iodate in Calcite Sebastien N. Kerisit,*,† Frances N. Smith,‡ Sarah A. Saslow,‡ Megan E. Hoover,§ Amanda R. Lawter,‡ and Nikolla P. Qafoku‡ †
Physical and Computational Sciences Directorate, Pacific Northwest National Laboratory, Richland, Washington 99352, United States ‡ Energy and Environment Directorate, Pacific Northwest National Laboratory, Richland, Washington 99352, United States § Environmental Engineering and Earth Sciences Department, Clemson University, Anderson, South Carolina 29625, United States S Supporting Information *
ABSTRACT: Iodate (IO3−) incorporation in calcite (CaCO3) is a potential sequestration pathway for environmental remediation of radioiodine-contaminated sites (e.g., Hanford Site, WA), but the incorporation mechanisms have not been fully elucidated. Ab initio molecular dynamics (AIMD) simulations and extended X-ray absorption fine structure spectroscopy (EXAFS) were combined to determine the local coordination environment of iodate in calcite, the associated charge compensation schemes (CCS), and any tendency for surface segregation. IO3− substituted for CO32− and charge compensation was achieved by substitution of Ca2+ by Na+ or H+. CCS that minimized the I−Na/H distance or placed IO3− at the surface were predicted by density functional theory to be energetically favored, with the exception of HIO3, which was found to be metastable relative to the formation of HCO3−. Iodine K-edge EXAFS spectra were calculated from AIMD trajectories and used to fit the experimental spectrum. The best-fit combination consisted of a significant proportion of surface-segregated IO3− and charge compensation was predominantly by H+. Important implications are therefore that pH should strongly affect the extent of IO3− incorporation and that IO3− accumulated at the surface of CaCO3 particles may undergo mobilization under conditions that promote calcite dissolution. These impacts need to be considered in calcite-based iodate remediation strategies.
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INTRODUCTION Iodine is a contaminant of significant environmental concern at facilities where irradiated nuclear fuel is stored, reprocessed, or disposed of because of its high toxicity, high bioaccumulation factor, extremely long half-life, and rapid mobility in subsurface environments.1 129I (half-life = 1.6 × 107 years) is present in groundwater at the U.S. Department of Energy’s Hanford Site in southeastern Washington State and Savannah River Site in South Carolina. Some of the challenges with predicting iodine behavior in the environment stem from its high solubility, multiple oxidation states, dynamic aqueous speciation across the entire pH range, and potential for interacting with organic matter.1 Potential remediation technologies for iodine contamination are currently being studied as part of environmental remediation activities at the Hanford Site. A speciation study2 of the groundwater at the Hanford Site determined that iodine was present mostly as iodate (IO 3−); however, small percentages of iodide (I−) and organically bound iodine were also detected. Moreover, iodine was found in association with calcite particles that precipitated following CO2 degassing during removal of groundwater from the deep surface,2 and a significant fraction of the total iodine content of sediments © 2018 American Chemical Society
from the Hanford Site was associated with the carbonate fraction.3 These findings are consistent with results from the literature on natural calcium carbonates4 and led to further research into the ability of calcite, and other calcium carbonate polymorphs, to sequester iodine as IO3−.5 Precedent for iodine association with calcium carbonate phases in nature exists in speleothems (i.e., cave features6,7). Other studies have used I/ Ca ratios in microscopic ocean-dwelling organisms, foraminiferans, or in marine carbonate deposits as a proxy for paleoredox conditions.4,8 Although not a calcium carbonate phase, the calcium iodate phase lauterite (Ca(IO3)2) is found in arid climates9 and serves as a useful benchmark for studying IO3− incorporation in calcium carbonate phases.4 Calcium carbonate is able to incorporate a wide range of metals and radionuclides,10,11 and a variety of studies have also examined oxyanion substitution into calcium carbonate phases.12−14,5 However, one of the challenges with the incorporation of IO3− into calcium carbonate is the resulting Received: Revised: Accepted: Published: 5902
January 18, 2018 April 20, 2018 April 26, 2018 April 26, 2018 DOI: 10.1021/acs.est.8b00339 Environ. Sci. Technol. 2018, 52, 5902−5910
Article
Environmental Science & Technology
Figure 1. Atomistic models illustrating IO3− incorporation at the CO32− position in calcite and charge compensation by Na+ in three nearestneighbor positions (Na B1, B2, and B3) and one distant position (Na B4) in the bulk and in one nearest-neighbor position (Na S1d) at the surface (the inset shows a top view of position Na S1d). The corresponding atomic positions in pure bulk calcite are also shown (0). Calcium is shown in blue, oxygen in red, carbon in brown, iodine in purple, and sodium in yellow.
A first series of calculations was performed to determine incorporation energies (Einc) using a generalized “products minus reactants” approach and solid-state reference phases.26 One CO32− was substituted by one IO3− in the supercell and the net charge thus introduced was compensated by replacing one Ca2+ cation either by Na+ or H+. Multiple initial positions of Na+ and H+ relative to IO3− were considered. Surface calculations were also performed, whereby slabs representing the lowest-energy and morphologically dominant (104) surface were “cleaved” from the optimized calcite unit cell.27 Hereafter, bulk incorporation configurations are labeled as BX, where X is a digit used to differentiate between the different positions of the charge compensating species, and surface incorporation configurations are labeled as SXd/h, where X represents the surface atomic layer iodate is incorporated in and d/h indicates whether the surface is dry (d, no water adsorbed) or hydrated (h, 1 water monolayer adsorbed). In each case, the label is preceded by either Na or H to indicate the charge compensating species. A second series of calculations employed key configurations determined in the first series to perform AIMD simulations and calculate EXAFS spectra. NVT (constant number of particles, constant volume, and constant temperature) AIMD simulations were performed and a minimum of 100 configurations were collected from each simulation at 50 fs intervals to calculate the I K-edge EXAFS spectrum. For each configuration, a cluster centered on the I atom was generated to calculate all the scattering paths using FEFF928−30 and the spectra of all configurations were averaged for comparison with experiment. The Fourier transform (FT) was applied to the averaged EXAFS spectra using IFEFFIT.31 The same approach was applied to standard iodate compounds (e.g., NaIO3). Computational details and additional information can be found in the Supporting Information (SI) document. Coprecipitation and EXAFS Measurements. An iodatedoped calcite sample was synthesized at room temperature from CaCl2, (NH4)2CO3, and NaIO3 solutions (without calcite seed crystals) using the approach detailed in the SI and was prepared for EXAFS analysis immediately after synthesis completion. EXAFS spectra for the iodate-doped calcite sample and three iodate standards, NaIO3, Ca(IO3)2, and KIO3, were collected at the Stanford Synchrotron Radiation Lightsource (SSRL) beamline 11−2 at the iodine K edge (33,169 eV).
charge imbalance from the aliovalent substitution of the CO32− group. On the basis of X-ray absorption spectroscopy measurements and density functional theory (DFT) calculations, Podder et al.5 concluded that IO3− substituted for CO32− in calcite and formed ionic bonds with two or three additional oxygen atoms. Although their electron microprobe analysis showed a positive correlation between sodium and iodine concentrations, Podder et al.5 were not able to determine the nature and location of the charge compensating species from their extended X-ray absorption fine structure spectroscopy (EXFAS) measurements. Questions regarding the role of IO3− segregation to the calcite surface also remain unanswered. The objective of this work was therefore to determine the local environment around IO3− incorporated in calcite, the most likely charge compensation schemes (CCS), and any tendency for surface segregation. This objective is critical, not only for evaluating the energetics and mechanisms of iodine incorporation, but also as a basis for determining its effects on calcite stability and solubility. The approach employed in this work consisted in first undertaking a systematic evaluation of IO3− incorporation schemes in bulk and surface environments to identify key representative configurations. Calculated EXAFS spectra were then generated from ab initio molecular dynamics (AIMD) trajectories of these configurations and used as components in a direct fit to the EXAFS measurements performed in this work. This approach has been shown to be a powerful method for extracting more structural information from EXAFS spectra than via traditional shell-by-shell fitting.15,16
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METHODS DFT Calculations. All of the plane-wave DFT calculations were performed with VASP (Vienna Ab-initio Simulation Package)17−20 using the projector augmented-wave (PAW) approach21,22 and the generalized gradient approximation exchange-correlation functional of Perdew, Burke, and Ernzerhof23,24 (PBE) with Grimme dispersion corrections (G).25 For all solid phases, a constant-pressure energy minimization was first performed to determine the optimized crystal structure (Table S1) at the PBE+G level of approximation and the optimized unit cell was then scaled to achieve the desired supercell size. 5903
DOI: 10.1021/acs.est.8b00339 Environ. Sci. Technol. 2018, 52, 5902−5910
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Environmental Science & Technology Measurements were made at 8.0 ± 0.2 K using an Oxford Instruments cryostat cooled with liquid helium. X-ray absorption data were obtained from 240 eV below the edge to 1100 eV above the edge. The data from 30 eV below the edge to 20 eV above the edge were obtained with 0.5 eV spacing. The data beyond 20 eV above the edge were obtained with a k-spacing of 0.05 and a k2-weighted collection time. The monochromator was detuned 20% to reduce the harmonic content of the beam. Transmission data (iodate standards; NaIO3, 7 scans; Ca(IO3)2, 6 scans; KIO3, 5 scans) were obtained using Ar filled ion chambers. Fluorescence data (iodate-doped calcite sample; 9 scans) were obtained using a 100 element Ge detector and were corrected for detector dead time. A NaIO3 reference (diluted in boron nitride) was used to account for minor shifts in energy between samples. Spectra were generated from raw data using SIXPack32 and then normalized using Athena.33
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Figure 2. Incorporation energy (Einc) as a function of I−Na/H distance for the cosubstitution of H+/IO3− or Na+/IO3− for Ca2+/ CO32− in bulk calcite and at the (104) surface. For H+/IO3−, configurations in which H+ is on the IO3− or CO32− group are shown with different colors. Labels indicate the configurations used in the AIMD simulations (blue labels, Figure 1; red and dark yellow labels, Figure 3). See Methods section for nomenclature of incorporation configurations.
RESULTS AND DISCUSSION Energetics of Na+/IO3− Cosubstitution. In the pure calcite structure (model 0 in Figure 1), each carbonate oxygen is coordinated to two calcium ions. Because of the planar configuration of the carbonate ion, these two nearest-neighbor calcium positions are symmetrically equivalent (C−Ca distance of ∼3.2 Å). In contrast, IO3− assumes a trigonal pyramidal geometry with its oxygen atoms in the same plane as neighboring carbonate groups. As a result, the local symmetry is broken when CO32− is substituted by IO3− and the two positions are no longer equivalent, whereby Ca2+ substitution by Na+ (Figure 1) can occur at a first nearest-neighbor position (Na B1), with a long optimized I−Na distance of ∼3.7 Å, or at a second nearest-neighbor position (Na B2), with a short optimized I−Na distance of ∼3.3 Å. Similarly, a calcium atom is positioned directly above and below a carbon atom along the [001] direction in the calcite structure (C−Ca distance of ∼4.3 Å), but introducing IO3− breaks the local symmetry resulting in two inequivalent positions for Na+ substitution with the one leading to the shortest optimized I−Na distance (∼3.4 Å) being considered here (Na B3). When Na+ is in a distant position, calcium ions occupy the three nearest-neighbor positions (Na B4). 2 × 2 × 1 supercell calcite models were built with increasing I−Na separation distances, including the three nearest-neighbor positions shown in Figure 1 and a number of distant positions (Figure 2). Positions Na B2 and B3, which minimized the I−Na distance, were the lowest-energy positions differing only by 0.02 eV. In contrast, position Na B2 was more stable than position Na B1 by approximately 0.4 eV. Podder et al.5 only considered position Na B1 and a distant position and, therefore, did not consider the global energy minimum for the IO3−−Na+ pair. At I−Na distances greater than ∼5 Å, averaged Einc values were on par with position Na B1 (1.3 eV), suggesting that energy-lowering effects due to defect clustering are minimized at greater distances. Variations in Einc above and below 1.3 eV are attributed to the number of effective carbonate layers separating iodine from sodium in the [001] direction. Energies are higher when at least 2 carbonate layers separate the defects and lower when only one carbonate layer is between the iodine and sodium, again pointing to the energetic preference for defects to cluster. IO3− incorporation at the (104) calcite surface, which dominates the morphology of calcite crystals, was also
considered. A first series of calculations (Figure S1) determined the energetically favored position of Na+ for IO3− incorporated in the topmost atomic layer (position Na S1d in Figure 1). When at the surface, the iodine atom moved toward the free surface with its oxygen atoms remaining in a relatively planar configuration and, as in the bulk case, the lowest-energy position for Na+ was the one that minimized the I−Na distance. A second series determined the depth-dependent energetics of Na+/IO3− cosubstitution for the lowest-energy Na+ position (Figure 2). Incorporation was much more favorable at the topmost atomic layer and rapidly converged with depth to the bulk values, indicating a strong preference for segregation to the calcite surface, in agreement with previous electronic structure calculations of selenite incorporation in calcite.34,35 The impact of surface hydration was also considered by adsorbing a water monolayer at the (104) surface, which had the effect of lowering the incorporation energy further (Figure 2). Energetics of H+/IO3− Cosubstitution. As with the first cosubstitution scheme, a series of calculations was performed to evaluate the effect of proton placement on Einc for H+/IO3− cosubstitution in calcite. Two distinct cases were tested (graphical description in Figure S2): (1) H+ associated with oxygen atoms of nearest-neighbor carbonate groups−either in the same carbonate layer as IO3− or in a nearest-neighbor carbonate layer and (2) H+ associated with each of the three iodate oxygen atoms. The incorporation energies indicated a preference for H+ to associate with oxygen atoms of nearest-neighbor CO32− groups over oxygen atoms of IO3−, even though the I−H distance was shortest for the latter (Figure 2). Key configurations labeled in Figure 2 and used in the AIMD simulations are shown in Figure 3. As for Na+/IO3− cosubstitution, the iodate oxygen atoms remained in plane with the carbonate groups in all cases. In the lowest-energy configuration (position H B2 in Figure 3), H+ was associated with the nearest-neighbor CO32− group thus forming HCO3−, the calcium vacancy was closest to the iodine atom, and IO3− accepted a hydrogen bond from HCO3−. 5904
DOI: 10.1021/acs.est.8b00339 Environ. Sci. Technol. 2018, 52, 5902−5910
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Environmental Science & Technology
In position H B1 (Figure 3), IO3− also accepted a hydrogen bond from HCO3− but the calcium vacancy and iodine atom were on opposite sides of the plane formed by the iodate oxygen atoms and were thus further apart, resulting in a less favorable incorporation energy. In cases where the H-bearing IO3− oxygen was too far from the Ca2+ vacancy to allow for proton transfer, the hydrogen atom pointed toward a neighboring CO3− instead. Configurations in which H+ was associated with iodate oxygen atoms yielded the least favorable incorporation energies (Figure 3), by at least 1 eV with respect to position H B2 (Figure 2). In these cases, the I−O(H) bond distance elongated to approximately 2.05 Å from 1.84 Å in IO3−. To compensate for the weakening of the I−O bond, one of the carbonate groups in the plane above that containing HIO3 distorted to reduce the I−O second-nearest-neighbor bond distance to approximately 2.4 Å compared to 2.7 Å in the case of IO3−. In two of the six cases in which it was initially positioned within 1 Å of an iodate oxygen (Figure S2), the distance between the H-bearing IO3− oxygen and the vacancy was short (Table S3), and the hydrogen atom was able to move to form a bond with a carbonate oxygen during the course of the energy minimization, indicating that there was no energy minimum associated with the formation of HIO3 for these configurations
Figure 3. Atomistic models illustrating IO3− incorporation at the CO32− position in calcite and charge compensation by H+ in three nearest-neighbor positions (H B1, B2, and B3) in the bulk and in one nearest-neighbor position (H S1d) at the surface (the inset shows a top view of position H S1d). Calcium is shown in blue, oxygen in red, carbon in brown, iodine in purple, and hydrogen in white. Calcium vacancies are shown by green dashed circles.
Figure 4. Experimental and calculated I K-edge EXAFS spectra (left; ΔE0 = 7 eV) and corresponding Fourier transform magnitudes (right; Hanning window, dk = 1 Å−1) for I in NaIO3 (top; 3.5 Å−1 ≤ k ≤ 14.7 Å−1), Ca(IO3)2·nH2O (n = 0 or 1; middle; 3.5 Å−1 ≤ k ≤ 14.7 Å−1), and KIO3 (bottom; 3.0 Å−1 ≤ k ≤ 14.0 Å−1) for temperatures ranging from 8 to 200 K. EXAFS data collected in this work at 8 K were complemented by data from Laurencin et al.36 and Yagi et al.37 5905
DOI: 10.1021/acs.est.8b00339 Environ. Sci. Technol. 2018, 52, 5902−5910
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Environmental Science & Technology Table 1. Comparison of Calculated I−O Bond Lengths (Å) in NaIO3 and Ca(IO3)2 with XRD Data NaIO3 I−O(1) XRD DFT Δ (%) a
1.802 1.827 1.4
a
Ca(IO3)2 I−O(2) 1.811 1.838 1.5
a
I(1)−O(1) 1.825 1.856 1.7
b
I(1)−O(2) 1.796 1.829 1.8
b
I(1)−O(3) 1.801 1.832 1.7
b
I(2)−O(1) 1.814 1.840 1.4
b
I(2)−O(2) 1.795 1.823 1.6
b
I(2)−O(3) 1.804b 1.828 1.3
Svensson and Ståhl 1988.38 bGhose et al. 1978.9
Figure 5. Effect of Na+ (1) and H+ (2) positions on the calculated I K-edge EXAFS spectra at 8 K (left; ΔE0 = 10 eV) and corresponding FT magnitudes (right; 3.5 Å−1 ≤ k ≤ 16.3 Å−1, Hanning window, dk = 1 Å−1). The effects of temperature (3) and surface segregation (4) for position Na B2 are also shown as examples. The experimental spectrum collected in this work at 8 K and its FT are also displayed for comparison in each panel.
adsorbed water monolayer reduced the incorporation energy further (Figure 2). Incorporation energies obtained for the lowest-energy configurations of the Na+/IO3− cosubstitution scheme were more favorable than those obtained for H+ associated with carbonate groups. However, the values of Einc were calculated using solid-state reference phases. To better reflect the aqueous conditions found in nature and in laboratory synthesis experiments, the energetics of the reaction
(Figure 2). These calculations therefore indicate that, depending on the relative positions of IO3−, the calcium vacancy, and H+, HIO3 may be metastable or may spontaneously dissociate to lead to the formation of HCO3−. This could explain the results of Podder et al.,5 who apparently only considered the case where HIO3 was metastable when H+ was in a nearestneighbor position, and thus only reported the formation of HIO3 and not the global minimum involving HCO3− (position H B2). As for cosubstitution with Na+, IO3− incorporation at the (104) calcite surface was also considered. Here again, a first series of calculations (Figure S3) was performed to determine the lowest-energy position for H+ charge-compensating IO3− incorporated in the topmost atomic layer. A second series evaluated the depth-dependence of the incorporation energy for this charge-compensating position (Figure 2). Consistent with the calculations already discussed, substitution of H+ on a nearest-neighbor CO32− was favored, the magnitude of the incorporation energy was greatly reduced when IO3− was present in the topmost atomic layer, and the presence of an
+ CaCO3 ·NaIO3 + H3O(aq) + ↔ CaCO3 ·HIO3 + Na(aq) + H 2O(l)
(1)
were evaluated by performing AIMD simulations of aqueous Na+ and H3O+ and of liquid water. The results of these simulations point to a negative energy of reaction and thus a preference for charge compensation of the incorporated IO3− by H+ over Na+ with respect to the aqueous ions (see Table S4 and associated text in the SI). 5906
DOI: 10.1021/acs.est.8b00339 Environ. Sci. Technol. 2018, 52, 5902−5910
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Environmental Science & Technology
Figure 6. Linear combination fits (k3 weight) to the experimental EXAFS spectrum collected in this work and to that of Podder et al.5 in reciprocal (left) and real (right) space using the EXAFS spectra calculated from AIMD trajectories of IO3− incorporated in calcite at 8 and 77 K (including S1d configurations). Contributions from each configuration are listed in Table S6.
EXAFS Standard Compounds. Simulated EXAFS spectra of four standard iodate compounds at various temperatures were compared to experiment to evaluate the simulation approach, develop an understanding of the effect of temperature on the EXAFS spectra, and determine how the local environment around IO3− might influence the EXAFS spectra. In addition to the EXAFS spectra collected at 8 K for NaIO3, Ca(IO3)2, and KIO3, data for NaIO3 and Ca(IO3)2·H2O obtained at 77 K by Laurencin et al.36 and data for KIO3 collected at 10, 100, and 200 K by Yagi et al.37 were also included in the comparison (Figure 4). Two main discrepancies were observed. First, a small offset between the calculated and experimental k3χ(k) spectra was apparent for high values of k, which was attributed to the slightly longer I−O bonds predicted by DFT compared to experimental data. Comparison of the I−O bond lengths obtained from an energy minimization of NaIO3 and Ca(IO3)2 with X-ray diffraction (XRD) data from the literature confirmed this result (Table 1). Because this difference was small (
