Uranium incorporation into amorphous silica.

Article pubs.acs.org/est

Uranium Incorporation into Amorphous Silica Michael S. Massey,*,†,‡ Juan S. Lezama-Pacheco,‡ Joey M. Nelson,§ Scott Fendorf,‡ and Kate Maher§ †

Department of Earth and Environmental Sciences, California State University, East Bay, Hayward, California 94542, United States Department of Environmental & Earth System Science, Stanford University, Stanford, California 94305, United States § Department of Geological and Environmental Sciences, Stanford University, Stanford, California 94305, United States ‡

S Supporting Information *

ABSTRACT: High concentrations of uranium are commonly observed in naturally occurring amorphous silica (including opal) deposits, suggesting that incorporation of U into amorphous silica may represent a natural attenuation mechanism and promising strategy for U remediation. However, the stability of uranium in opaline silicates, determined in part by the binding mechanism for U, is an important factor in its long-term fate. U may bind directly to the opaline silicate matrix, or to materials such as iron (hydr)oxides that are subsequently occluded within the opal. Here, we examine the coordination environment of U within opaline silica to elucidate incorporation mechanisms. Precipitates (with and without ferrihydrite inclusions) were synthesized from U-bearing sodium metasilicate solutions, buffered at pH ∼5.6. Natural and synthetic solids were analyzed with X-ray absorption spectroscopy and a suite of other techniques. In synthetic amorphous silica, U was coordinated by silicate in a double corner-sharing coordination geometry (Si at ∼3.8−3.9 Å) and a small amount of uranyl and silicate in a bidentate, mononuclear (edge-sharing) coordination (Si at ∼3.1−3.2 Å, U at ∼3.8−3.9 Å). In iron-bearing synthetic solids, U was adsorbed to iron (hydr)oxide, but the coordination environment also contained silicate in both edge-sharing and corner-sharing coordination. Uranium local coordination in synthetic solids is similar to that of natural Ubearing opals that retain U for millions of years. The stability and extent of U incorporation into opaline and amorphous silica represents a long-term repository for U that may provide an alternative strategy for remediation of U contamination.

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INTRODUCTION Uranium derived from anthropogenic and natural sources is a contaminant in groundwater globally, and as such, the development of effective geochemical models and remediation strategies for U remains a key challenge. Uranium is highly toxic and radioactive and can cause organ damage in humans and animals, which makes both natural and anthropogenic U contamination a human and ecosystem health risk. Uranium has an average concentration in continental crust of approximately 3 mg kg−1, and it can be naturally concentrated through a variety of low- and high-temperature processes. Common sources of uranium contamination include relatively undisturbed natural ore bodies such as those at Koongarra, Australia,1 along with anthropogenic activities such as nuclear weapons production, and U mining and milling.2 The United States Department of Energy (US DOE) alone manages 1.5 billion m3 of contaminated groundwater and 75 million m3 of contaminated soil or sediment,2 with U being one of the most common radionuclide contaminants.3 Globally, U water contamination from nuclear activities results from 900 million m3 of tailings distributed across nearly 6000 ha of land area on every continent except Antarctica.4 As uranium is transported by groundwater and retained by soil and sediment solids, an understanding of U retention processes is crucial for maintaining ecosystems and human health. Soil and sediment solids can be both sources and sinks © 2014 American Chemical Society

of U, either supplying U or removing it from water. Typical sinks of U from groundwater include precipitation of U minerals, U adsorption on solids, and U incorporation into non-U minerals. Uranium minerals include U(IV) phases such as uraninite (UO2, UO2+x) and U(IV) phosphates.5 Uranium(V/VI) minerals such as schoepite, 6 minerals of the uranophane groups (e.g., sodium boltwoodite, Na(UO2)(SiO3OH)·1.5H2O),7,8 and other, rarer minerals such as vorlanite.9 Uranium-bearing palagonite and metatorbernite10 are also found in some contaminated sediments. Additionally, U can adsorb to a variety of solids, ranging from iron (hydr)oxides11 to clay minerals.12,13 Furthermore, U(V/VI) can be incorporated into minerals such as iron (hydr)oxides14 or opal (SiO2·nH2O).15 Due to the apparent stability of U within opal, 230Th−U and U−Pb geochronology have been used to study the age of soil deposits16 as well as the geologic history of sites such as Yucca Mountain, Nevada,17−19 although the controls on U retention are not well-known. In near surface environments, opal, amorphous silica, and silica gels (hereafter referred to generally as “amorphous silica”) precipitate from Sibearing solutions that reach Si supersaturation due to Received: Revised: Accepted: Published: 8636

March 3, 2014 June 29, 2014 July 1, 2014 July 1, 2014 dx.doi.org/10.1021/es501064m | Environ. Sci. Technol. 2014, 48, 8636−8644

Environmental Science & Technology

Article

(hydr)oxides, and the potential effectiveness of U/Si precipitates as a strategy for remediating U contamination. Accordingly, here we elucidate the local molecular structure of U in amorphous silica and iron oxide-bearing amorphous silica, and we compare natural U-bearing opals with synthetic amorphous silica to determine U uptake mechanisms in natural samples.

weathering in aquifers and/or evaporation in soils and exposed sediments.20,21 Other ions, such as U, Fe, and Al, can be incorporated into or adsorbed onto the amorphous silica. Reductive precipitation of U(IV) minerals such as uraninite has been studied within the context of groundwater remediation.22−24 However, the resulting precipitates can be highly sensitive to reoxidation by oxygen, nitrate, Fe(III) (hydr)oxides, or Mn(IV/III) oxides.25−29 Adsorption of U to minerals such as iron (hydr)oxides,11 clay minerals,12,13,30 quartz,31 and amorphous silica32,33 has also been studied in order to understand the environmental fate and transport of U.34 Uranium adsorbed on redox-insensitive solids is not prone to oxidative dissolution/desorption, but changes in aqueous geochemical conditions may result in desorption of U. Thus, adsorption and reductive precipitation of U do not necessarily result in stable sinks for this contaminant, especially in temporally variable environmental conditions. Uranium inclusion within solids that are not composed of redox active elements represents a possible stable sink for U. Uranium will only be rereleased into groundwater (or surface water) from these minerals if the mineral itself dissolves. Further, uranium minerals, particularly U(VI) minerals such as uranophane and metatorbernite, are found in contaminated areas7,10,35 and are less vulnerable to changes in groundwater chemistry. Indeed, uranyl phosphate precipitation has been used in attempts to remediate U contamination.36 Uranium in silicates may serve as another long-term sink and offer another remediation strategy, as U in opals has been shown to persist in the solid for millions of years.17 Moreover, natural U-bearing opals have been found with U at concentrations of hundreds of milligrams per kilogram or more,15,17,21,37 and amorphous silica gels may contain U at tens of mass percent.33 Thus, opal and amorphous silica represent potentially stable sinks for U in the environment, able to retain U despite possible changes in redox status, pH, or other groundwater conditions. Several authors have investigated the molecular basis of U adsorption on natural and synthetic Si gel.12,32,33,38,39 For example, Allard et al.33 found that U was often retained as small uranyl silicate domains within the bulk gel. Reich et al.32 found little to no association between U and Si at low pH (3.5−4.5). Soderholm et al.40 examined Si/U coprecipitates resulting from solutions having 50 mM uranyl and 100 mM silica concentrations and found U−Si and U−U pair correlations indicating long-range ordering of the precipitate. Uranyl and silicate concentrations of this magnitude may not be easily achieved in surface or near-surface environments far from sources of contamination, and thus may not represent structures relevant to geochronology or groundwater remediation. In soils and sediments, various components of the mineral matrix can bind uranium and serve to compete with partitioning into opaline silica. In particular, iron (hydr)oxides such as ferrihydrite are excellent sorbents for U under a widerange of conditions and may preferentially bind U relative to precipitating silica gels. However, it is unclear whether iron oxides in silica-saturated environments preferentially retain U or, if they do, whether the silica may encase the U-bearing iron (hydr)oxide. The issue is compounded by analytical difficulty in distinguishing between U coordinated by second-neighbor Si or Fe. Distinguishing between U retained by silicate and U retained by iron (hydr)oxides provides insights into the structure of U retained in natural opals, the competitiveness of silicates in sequestering U in the presence of iron

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METHODS Our experimental approach began with synthesis of uranyl/ silica precipitates (with or without iron inclusions). The U coordination environment of these well-constrained synthetic precipitates was then compared with that of natural U-bearing opals using extended X-ray absorption fine structure (EXAFS) spectroscopy. Synthetic Sample Preparation. Samples were prepared by titrating 200 mL of 37 μM uranyl acetate, 10 mM PIPES buffer, and 12 mM sodium metasilicate to an initial pH of 5.6 ± 0.1, and the silicate solution served as the primary buffer to maintain a (nearly) constant pH. The PIPES buffer was included as an extra precaution to ensure a pH < 7.0, as MuñozAguado and Gregorkiewitz41 found that pH ≤ 7.0 is critical for the precipitation of amorphous silica from solution. The initial Si concentration of 12 mM is higher than typically found in most waters, but it may reflect evapoconcentrated fluids where amorphous silicates are observed, or forced-precipitation remediation conditions. In one solution, ferrihydrite slurry was also present at a concentration of 0.75 mM (as Fe), approximately equivalent to a solid concentration of 80 mg L−1. Ferrihydrite was prepared as described previously.42 Briefly, ∼150 mM ferric chloride was vigorously mixed using a mechanical stirrer and rapidly (∼15 min) titrated to a pH of 7.0−7.5 using 0.4−1 M sodium hydroxide. After 1−2 h of equilibration and further addition of sodium hydroxide to establish a stable pH between 7.0 and 7.5, the slurry was allowed to settle, and the supernatant was decanted. The slurry was then centrifuged and washed with deionized water (18 MΩ) five times to remove residual salts. Two-line ferrihydrite was confirmed using X-ray powder diffraction, and the slurry density was measured by dissolving the solid in 6 M HCl along with complete reduction of Fe(III) to Fe(II) using 0.5 M hydroxylamine hydrochloride, followed by the ferrozine assay.43 All solutions were stored in plastic bottles, and the incubation itself was performed in 250 mL polypropylene (Nalgene) bottles. Solutions were incubated at 25 °C on a rotary shaker at 120 rpm for 28 days. The end-point of the experiment was chosen when silicate precipitate was visible. At the conclusion of the incubation, small globes of precipitate adhered to the container walls during shaking. The suspensions were vacuum-filtered through a large glass-fiber filter and rinsed with deionized water (18 MΩ). The wet solids were scraped from the filter, weighed for gravimetric water content, and left to dry in the air. After drying, the solids were ground for analysis. Solution Collection and Analysis. Triplicate aliquots of 1 mL of solution were withdrawn using a syringe, filtered through 0.22 μm syringe membranes, diluted with ∼2% trace metal grade HNO3 and analyzed via ICP-OES (Thermo Scientific ICAP 6300 Dual View, Thermo Fisher Scientific, Waltham, MA) for Si. Uranium was measured via ICP-MS (Thermo Scientific XSERIES 2, Thermo Fisher Scientific, Waltham, MA). 8637

dx.doi.org/10.1021/es501064m | Environ. Sci. Technol. 2014, 48, 8636−8644


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collected using an in-line ionization chamber, and fluorescence spectra were collected simultaneously using either a 13- or 30element Ge solid-state detector (Canberra, Connecticut, United States). An in-line Y foil was used to ensure energy calibration; the inflection in the Y K-edge spectrum (defined using the first-derivative peak position) was calibrated to 17 038.4 eV. Data calibration and averaging were performed using SixPack,45 and data were normalized and fit using the Athena and Artemis software packages46,47 and FEFF 6.48 Detailed normalization and fitting parameters are provided in the SI. High Resolution Synchrotron X-Ray Powder Diffraction. Powders of synthetic and natural solids were packed in 0.3 mm borosilicate glass capillaries (Hampton Research, Aliso Viejo, California), sealed with 5 min epoxy (ITW Devcon, Danvers, Massachusetts) and analyzed at SSRL beamline 7−2. The incident beam energy was maintained at 16 keV (λ = 0.77 Å), and precise energy calibration was achieved using a LaB6 calibration standard in a borosilicate glass capillary. Powder diffraction data were collected over a Q-space range of 0.8 to ∼8−10 Å−1 using a single-channel energy dispersive Vortex solid-state detector (Canberra, Connecticut, United States). The synthetic solids were identified as amorphous silica, having no crystalline structure visible over the capillary background; natural samples were all opal-CT. Opal-CT diffraction patterns were compared to known opal diffraction patterns.49,50 Microscale X-Ray Fluorescence Mapping. Fragments from each of the hyalite, fire opal, and mixed zones of the natural opal were encased in epoxy and polished as described above. X-ray fluorescence maps were collected at SSRL beamline 2−3 using a motorized sample stage and a singleelement Vortex Si-drift solid-state detector (Canberra, Connecticut, United States) to resolve X-ray fluorescence emission energies. Incident X-ray beam energy was maintained at 17 200 eV using a Si(220) double-crystal monochromator in the φ = 0° orientation. The X-ray beam was rastered across the sample with a pixel size of 10 μm × 10 μm and a count time of 35 ms. X-ray fluorescence maps were normalized and analyzed using the SMAK software package. Spatially Resolved Elemental/Isotopic Analysis of Natural Opal Using SHRIMP-RG. The epoxy-encased, polished opal fragments described above were coated in a thin layer of gold to prevent sample charging under an ion beam and analyzed using the SHRIMP-RG doubly focusing mass spectrometer. The fragments were analyzed in a manner similar to that described in Maher et al.16 Briefly, the primary 16 2− O ion beam was passed through a 100 μm Kohler aperture, with a brightness of 6, resulting in a primary beam intensity of ∼2 nA and a spot size of 30 μm. Each mass cycle monitored peaks for 30Si, 27Al16O, 54Fe, 74Ge, and 238U, as well as mass 224 (28Si416O7). Sample measurements were normalized using the 30 Si peak and compared to NIST SRM 611 to determine elemental ratios and concentrations.16

Description and Sampling of Geologic Materials. Samples of U-bearing opal from the Virgin Valley, Nevada, United States15 were obtained from J. B. Paces and L. A. Neymark at the United States Geological Survey in Denver, Colorado, United States. These fragments are commonly known as “BZ-VV”.37 A second, nonprecious opal (presumed to be from Opal Butte Mine in Oregon, United States) was obtained from a commercial source (Opal Butte Mining Company, College Place, Washington, United States) and is hereafter abbreviated as “OB-OR”. The sample is similar to other Opal Butte Mine specimens, which are found in opalfilled geodes.44 The opal has three distinct zones attached to the host rock: (1) a clear zone of colorless, transparent silica material (hyalite), (2) an orange zone of silica mixed with colloidal yellow/orange iron precipitate (fire opal), and (3) a “mixed” zone of yellow/orange iron clasts encased in colorless, transparent hyalite. The suspended iron clasts in the “mixed” zone range in size from less than 1 mm to several millimeters in diameter. The zones are arranged in distinct bands, with the “mixed” zone nearest to the host rock, followed by the fire opal zone, then the hyalite zone. The host rock shows signs of substantial alteration, suggesting that the opaline silica precipitated from low-temperature hydrothermal fluids or acidic groundwater. The distinct zonation of the silica suggests a multistage formation history. Samples from each of the three zones of opaline silica were taken using a rock saw, and photographs of the sample and fragments are shown in Figure SI1. Solid Phase Analysis. Synthetic solids were weighed after harvest, dried, and reweighed for gravimetric water content. The dried synthetic solids were ground and analyzed using acid digestion and chemical analysis, XAS, and X-ray powder diffraction (XRD). Natural opal samples were also ground and analyzed using XAS and XRD. Fragments of natural opals were encased in epoxy (Struers EPOES resin and EPOAR hardener), polished to remove scratches greater than ∼1 μm in depth, and analyzed for U, Fe, and Si concentrations using the Sensitive High Resolution Ion Microprobe−Reverse Geometry (SHRIMP-RG) at Stanford University. The same sample mounts were also examined using microscale X-ray fluorescence at beamline 2−3 at the Stanford Synchrotron Radiation Lightsource (SSRL). Uranium and Fe concentrations were measured via point measurements using the SHRIMP-RG, while the larger-scale spatial distribution of elements was investigated using X-ray fluorescence mapping. Acid Digestion and Chemical Analysis. Subsamples of 20 mg of U/Fe/Si and U/Si solids were dissolved using sequential HF and aqua regia digestions to drive off silica. Digested samples were brought to analysis volume with 2% HNO3, ultrasonicated, and analyzed for Fe on a Nu AttoM high-resolution magnetic sector ICP-MS (Nu Instruments Limited, Wrexham, United Kingdom). X-Ray Absorption Spectroscopy. Uranium XAS spectra were collected for both of the synthetic U/Fe/Si and U/Si solids, as well as the BZ-VV opal and a sample from the “mixed” zone of the OB-OR opal. Powders of synthetic and natural solids were mixed with BN powder to obtain a working sample volume of ∼100 mg. The samples were then packed into thin Al sample holders, encased in Kapton tape, and analyzed at beamlines 4−1 and 11−2 at SSRL. Uranium L3-edge X-ray absorption spectra were collected; the X-ray beam incident energy was controlled via a Si(220) double-crystal monochromator in the φ = 0° orientation. Transmission spectra were

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RESULTS Uranium Incorporation into Synthetic Fe/Si Solids. Both the U/Si solid and the U/Fe/Si solid removed approximately 97% of the initial U from the solution. Additionally, ∼80−85% of initial 12 mM Si in solution was precipitated, with Si concentrations approaching equilibrium for amorphous silica after 28 days. On the basis of U and Si removed from solution, and Fe by solid dissolution, the U/Si 8638



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solid had a U content of ∼10 g kg−1 (dry weight) and a U/Si molar ratio of ∼1:250. The U/Fe/Si solid also had a U content of ∼10 g kg−1 (dry weight), 5.6 wt % Fe, and a U/Fe/Si molar ratio of ∼1:21:250. The solution pH remained at pH 5.6 ± 0.1 throughout the experiment, buffered by silicic acid. Gravimetric analysis of the synthetic U-bearing silicate samples (both with and without ferrihydrite) contained about 90%(±5%) water by weight. On the basis of the Si stoichiometry and mass balance, the dominant solid precipitate can be approximated by the formula SiO2·2H2O(s). Air-dried amorphous silica was used for the remainder of the analyses. Extended X-ray absorption fine structure (EXAFS) analysis (Table 1) indicated that uranyl in the synthetic U/Si solid was Table 1. Detailed U L3-Edge EXAFS Shell-by-Shell Fitting Results for Synthetic U-Bearing Silicate Solidsa coordination number U/Si solid U−Oaxial

2.00b

U−Oaxial MS U− Oequatorial, 1 U− Oequatorial, 2 U−Sinear

2.00c 2.84 (±0.19)d

U−Sifar

2.00 (±0.49)

U−U

0.55 (±0.25)

U/Fe/Si solid U−Oaxial

2.16 (±0.19)d 0.33 (±0.24)

2.00

b

U−Oaxial MS U− Oequatorial, 1 U− Oequatorial, 2 U−Fe

2.00c 1.74 (±0.55)d

U−Sinear

0.49 (±0.40)

U−Sifar

3.04 (±0.91)

4.26 (±0.55)d 1.0b

distance (Å) 1.81 (±0.02) 3.64c 2.39 (±0.02) 2.20 (±0.02) 3.15 (±0.06) 3.89 (±0.04) 3.85 (±0.05)

1.82 (±0.02) 3.66c 2.51 (±0.03) 2.29 (±0.01) 3.49 (±0.04) 3.13 (±0.05) 3.90 (±0.04)

σ2

ΔE0 (eV)

0.007 (±0.001) 0.014c 0.007b

7.19 (±2.29) 7.19c 7.19c

0.007b

7.19c

0.008b

7.19c

0.010b

7.19c

0.007b

7.19c

0.003 (±0.001) 0.007c 0.006c

7.73 (±4.38) 7.73c

Figure 1. Uranium L3-edge EXAFS shell-by-shell fits for a synthetic U/ Si precipitate obtained from reacting 37 μM uranyl with 12 mM silicate for 28 days, at pH 5.6. The data (solid yellow line), fit (dotted black line), and fit paths (colored lines) are shown in (a) k-space as a k3-weighted EXAFS χ function and (b) R-space as the Fourier Transform magnitude of a k3-weighted EXAFS function (paths are inverted in R-space). The fits illustrate U coordinated by axial and equatorial O atoms, as well as second-neighbor Si atoms.

7.73c 7.73c 7.73c

0.006 (±0.002) 0.008 (±0.003) 0.004b

7.73c

0.007b

7.73c

7.73c

Soderholm et al.40 The contribution (∼15% of the secondneighbor Si, on the basis of coordination number) from the edge-sharing U−Si distance of ∼3.1−3.2 Å (Figure 1, Table 1) indicates that uranyl is coordinated by silicate primarily in a corner-sharing geometry. These U−Si distances correspond to those observed in U−Si minerals like soddyite and are similar to the findings of other investigators.8,40 In the present study, only a minor contribution from a U−U pair was observed, in contrast to previous investigations.12,33,40 The dearth of longrange U−U pair correlations indicates a lack of long-range order in the small soddyite-like domains that contain U. Uranyl in the U/Fe/Si solid is coordinated by both secondneighbor Fe and second-neighbor Si (Table 1; Figure 2). The U EXAFS spectrum of the U/Fe/Si solid could not be described solely by U(VI) adsorbed to ferrihydrite (Figure SI4a) and rather illustrates attributes from second-neighbor Si in addition to Fe. The U EXAFS spectrum also could not be fit with a combination of U-ferrihydrite and U-silica (the Uferrihydrite is the only nonzero component of the fit in Figure SI4a), suggesting that the solid is not simply a mixture of two U-bearing phases (U adsorbed to ferrihydrite, and U in silica). The EXAFS spectrum indicates a single Fe at 3.49 Å (Table 1, Figure 2), indicative of an edge-sharing (bidentate, mono-

a Solids were obtained by reacting 37 μM uranyl with 12 mM silicate for 28 days, at pH 5.6 (“U/Si solid”) and 37 μM uranyl with ferrihydrite slurry and 12 mM silicate for 28 days, at pH 5.6 (“U/Fe/Si solid”). bParameter was set as a constant value in the fit. cParameter was defined as a function of other parameters in the fit. dParameter was constrained to sum to U−Oequatorial, total = 5 (U/Si) or 6 (U/Fe/Si) in the fits.

coordinated by second-neighbor Si atoms and a small amount of second-neighbor U atoms (Figure 1). The U−O coordination numbers and bond distances indicate that the central U atom is coordinated by two axial and about five equatorial O atoms, with substantial disorder present in both (axial and equatorial) planes. A U−Si distance of ∼3.1−3.2 Å (Table 1) is indicative of U coordinated by silica in an edgesharing (bidentate, mononuclear) coordination geometry, and a U−Si distance of ∼3.8−3.9 Å (Table 1) is indicative of U coordinated by silica in a (double) corner-sharing (bidentate, binuclear) geometry, as noted by Sylwester et al.12 and 8639



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Figure 3. Illustration of the molecular coordination environment of the (a) synthetic U/Si and (b) synthetic U/Fe/Si precipitates, as determined by U EXAFS spectroscopy. The U/Si precipitate has similarities to the coordination environment of U in soddyite (a uranyl silicate mineral) but with limited long-range structure and only a small contribution from a U−U pair. The U/Fe/Si precipitate is similar, though the uranyl-silicate complex appears to be associated with the iron oxide surface, and no U−U pair is evident. Precipitates were obtained from reacting 37 μM uranyl and 12 mM silicate for 28 days, in the presence or absence of ferrihydrite slurry, at pH 5.6.

Figure 2. Uranium L3-edge EXAFS shell-by-shell fits for a synthetic U/ Fe/Si precipitate obtained from reacting 37 μM uranyl with ferrihydrite slurry and 12 mM silicate for 28 days, at pH 5.6. The data (solid purple line), fit (dotted black line), and fit paths (colored lines) are shown in (a) k-space as a k3-weighted EXAFS χ function and (b) R-space as the Fourier Transform magnitude of a k3-weighted EXAFS function (paths are inverted in R-space). The fits illustrate U coordinated by axial and equatorial O atoms, as well as secondneighbor Fe and Si atoms.

diffractometer, no additional crystalline iron-bearing phases could be distinguished from the diffraction pattern. This suggests that the iron particles are very low in concentration, very small in particle size, or have limited crystallinity. Uranium concentration in the OB-OR sample was much lower than in the BZ-VV opal. In both the clear and “mixed” zones of the hyalitic opal (Figures SI1−SI2), U concentration varied from 0.03 mg kg−1 to 1.1 mg kg−1, with an average concentration of 0.6 mg kg−1, based on an average of n = 9 independent analysis points using SIMS; variability was quite high due to the low U concentration, and the uncertainty with these measurements is at least ±50%. In areas of the fragments near the parent material, the U concentration in silica was 2.3− 3.2 mg kg−1 (n = 7). In the fire opal silicate zone (Figures SI1− SI2), the U concentration was 0.6−1.8 mg kg−1 (n = 3). The parent material U concentration ranged from 2.1 to 5.0 mg kg−1 (n = 3). The uranium concentration was low enough in these samples that U could not be distinguished from Rb via synchrotron X-ray fluorescence mapping due to the overlap of X-ray emission lines. On the basis of the energy-dispersive fluorescence spectra, Rb makes up at least 95% of the counts in the U+Rb energy range. Iron, Si, and U+Rb fluorescence maps are shown in Figure SI2. The U L3-edge EXAFS spectra (Figure 4) of synthetic U/Fe/ Si and BZ-VV opal are similar, and a comparison of the spectra

nuclear) complex, consistent with the general findings of uranyl bound to iron (hydr)oxides having U−Fe bond distance of 3.40−3.49 Å. The U−Si bond distance of ∼3.9 Å (Table 1) indicates that U is coordinated by silicate in corner-sharing (binuclear, bidentate) geometry, and the coordination number of ∼2−3 suggests a double-corner (binuclear) moiety. As in the U/Si solid, a small contribution (10−20% second-neighbor Si) from a U−Si distance of ∼3.1−3.2 Å was present in the EXAFS spectrum, indicative of an edge-sharing Si contribution. Schematic representations of the U coordination environments in the U/Si and U/Fe/Si solids are given in Figure 3. Natural Materials. Uranium concentration in the BZ-VV opal (opal-CT) from Virgin Valley, Nevada ranges from ∼10 mg kg−1 to ∼600−900+ mg kg−1.15,37 On the basis of U−Pb geochronology, the opal formed at 2.35 ± 0.09 Ma.37 Iron concentration in the BZ-VV opal was reported to be 2 g kg−1, with iron inclusions and staining reportedly common.15 Similar fragments to the BZ-VV opal analyzed in this study contained 600 ± 25 mg kg−1 U and ∼500 mg kg−1 Fe. X-ray powder diffraction revealed that the OB-OR opal was opal-CT, as in, e.g., Guthrie et al.12,13,50 (Figure SI3). Despite the improved sensitivity to minor crystalline components via the use of synchrotron radiation and a high-resolution 8640



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U/Fe/Si solid is in fact such a mixture; this is a limitation of the analysis. Nonetheless, U adsorption to iron (hydr)oxides is relatively rapid, occurring within minutes to hours,6 while the observed precipitation of silica occurred over several weeks. This sequence of reactions (U adsorption to ferrihydrite, followed by precipitation of a silica coating on the ferrihydrite) could explain the observed U coordination environment without invoking a mixture of phases. For the synthetic and natural silica, uranyl was coordinated primarily by Si in an edge-sharing and double-corner sharing geometry. Though the U−O and U−Si distances and coordination numbers are consistent with corresponding EXAFS-derived bond distances in uranyl silicate minerals such as soddyite,8 a U−U interatomic distance was only observed in the U/Si synthetic solid in the present study. This is likely due to lower uranyl concentrations (∼1 mass % U). Allard et al.33 also found U−U distances using EXAFS in natural U-bearing silica gels, with ∼400 g kg−1 U. Soderholm et al.40 reported U−U interatomic distances of 5−7 Å in synthetic precipitates using total X-ray scattering pair distribution function analysis. However, the precipitates in Soderholm et al.40 formed at 1000+ times higher U concentrations (50 mM) compared to this study (37 μM). Higher U and Si concentrations also led to faster precipitation rates: Soderholm et al.40 reported nearly instantaneous precipitate formation, whereas in the present study precipitates took weeks to form. High U concentration and rapid precipitation rates may act in concert, increasing the prevalence of U−U interatomic distances and indicating longer-range structural order. Reich et al.32 reported an anomalously short U−Si interatomic distance of 2.72 Å for U(VI) adsorbed on silica gel at pH 3.5−4.5, and Wheaton et al.53 modeled possible structures that might be responsible for this short U−Si distance. However, no such interatomic distance was necessary to model the EXAFS spectrum of the U/Si precipitates in the present study. Even at U loadings of up to 59 g kg−1, Reich et al.32 observed no U−Si interatomic distances in the ∼3−4 Å range, in contrast to findings of other investigators.33,40 The differences in observations may be explained by the lower pH at which Reich et al.32 conducted their experiments (pH 3.5−4.5): the U may not have been predominantly adsorbed. Milonjic et al.51 found that U does not adsorb to Si gel at pH < 2.8, and even in the pH range of 3.5−4.0, adsorption of [Uinitial] = 0.2 mM was only ∼30−80%. Near-complete adsorption is only achieved at a pH greater than ∼4.5−5.0. Compared to previous studies of U/Si solids, we examined lower U concentrations40 and lower U loadings on some silica gels (such as certain samples in Allard et al.33 with U concentrations as high as 400 g kg−1). The system also had a higher pH (pH 5.6) compared to previous EXAFS-based studies.32 The experimental conditions of the present study more closely mimic conditions that might be found in nearsurface depositional environments of amorphous silica. The range of typical groundwater U concentrations commonly encountered at contaminated near-surface sites is between ∼10−6 and 10−4.5 M,54 while Si concentrations typically range from ∼10−4 to ∼10−2 M.33 Silicate precipitates may take weeks or longer to form under such conditions.55 Therefore, the U coordination environment in the synthetic solids in the present study is a reasonable model for slowly forming U-rich Si solids found in natural systems. In addition to binding strongly with silicate, U in natural systems may partition onto competitive sorbents such as iron

Figure 4. Uranium L3-edge EXAFS spectra and linear combination fitting results for the BZ-VV opal. The BZ-VV opal from Virgin Valley, Nevada was fit with a single component: the synthetic U/Fe/Si solid obtained from reacting 37 μM uranyl with ferrihydrite slurry and 12 mM silicate for 28 days, at pH 5.6. Fit is shown with black dotted lines, and fit residuals are shown in gray. The EXAFS spectra of the synthetic solids and U adsorbed to ferrihydrite are shown for comparison.

gives a reduced X2 of 0.44. Further, the EXAFS spectra for the U/Si solid and the “mixed” opal are consistent in amplitude and similar in phase, although the low U concentration in the OBOR sample limited data quality and spectral comparison (Figure SI4b). From the EXAFS fitting results of the synthetic precipitates (Figures 1 and 2 and Table 1), and the relative amplitudes and phase of the EXAFS spectra, it may be inferred that U is associated with Fe in the BZ-VV opal (Figure 4) but may not be associated with Fe in the OB-OR opal (Figure SI4b).

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DISCUSSION The synthetic U/(Fe)/Si precipitates retained approximately 10 g kg−1 (1% by dry mass) of U(VI), an amount generally higher than typical concentrations of 10 mg kg−1 to 1000 mg kg−1 observed in many U-bearing opals15,16,37 but lower than reported concentrations of U retained in some opals and on natural silica gels.21,33 For natural U-bearing silica solids from the Peny uranium mine in the Massif Central in France, the amount of U was inversely related to the presence of other elements such as Fe and Al;33 that trend was not observed here. In fact, despite the anticorrelation between Al/Fe and U observed by Allard et al.,33 U appears to partition onto the Fe inclusions in this study, consistent with the high affinity of U for both iron (hydr)oxide and amorphous silica. At the experimental pH of 5.6, nearly 100% of the U in the system is adsorbed on ferrihydrite.11,32 However, Milonjić et al.51 and Michard et al.52 also reported that U adsorbs on silica gel in this pH range, suggesting that our results could reflect “cooperative” sorption between Fe oxides and silica during opal formation. Our analysis of the synthetic and natural Fe-bearing solids illustrates that the coordination environment of U is best described as a complex with both the iron (hydr)oxide and silica; it is not simply a mixture of U-bearing amorphous silica and U adsorbed to ferrihydrite (Figure SI4a). It is possible that the U-bearing silica phase is simply unidentified and that the 8641



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reservoir that may be suitable for use in remediation of U contaminated sites. A U-bearing amorphous silica material may be synthesized in weeks, but at the molecular level the U is similar to U in geologic materials that has been retained for millions of years. If U/Si precipitates similar to those in this study prove to be as recalcitrant as opaline material, amorphous Si precipitation may prove to be a useful strategy for the remediation of U-contaminated soils and sediments.

(hydr)oxides (e.g., ferrihydrite, goethite, hematite, or magnetite). In this study, ferrihydrite slurry was added to the uranyl/Si solution to mimic the presence of iron oxide inclusions reported by Zielenski.15 When ferrihydrite was present, U in the resulting solid was coordinated (at least partially) by the iron (hydr)oxide inclusions, as illustrated by Fe backscatterers at ∼3.49 Å. The Fe was bridged through two equatorial oxygens, indicating a mononuclear, bidentate coordination geometry (and consistent with one Fe atom at this distance). In addition to the Fe atoms in the coordination environment, Si also plays an important role, as Si is observed at two distances, ∼3.1 and ∼3.9 Åidentical to the U−Si interactions noted in the absence of Fe. The shorter U−Si distance is indicative of edge-sharing coordination geometry (U bridged through two equatorial oxygen atoms per Si) and a more distant (Si) double corner-sharing coordination geometry (bridged through ∼1 equatorial oxygen atom per Si). The U−Fe interatomic distance indicates that the U is adsorbed to ferrihydrite, as described by previous investigators.11,56 Since EXAFS spectra provide a measurement of the average coordination environment of all of the U atoms in the solid, and the coordination number for the U−Fe distance at ∼3.49 Å is ∼1, most, if not all, of the U in the solid appears to be adsorbed to the ferrihydrite in a mononuclear, bidentate coordination geometry. However, the U−Si interatomic distances also indicate that the U acts as a bridging cation between the ferrihydrite and silicate in both corner-sharing and edge-sharing coordination geometries. That is, the entire U−Si complex (similar to the structure of the noferrihydrite U/Si precipitate in this study, but without a U−U pair) is apparently adsorbed to the ferrihydrite surface (Figure 3). Therefore, the iron (hydr)oxide is a competitive sorbent for uranyl in the synthetic system, and uranyl appears to be retained on iron (hydr)oxide inclusions in the silicate matrix. The absence of a U−U pair in the U/Fe/Si sample is likely due to competitive adsorption of U by the iron (hydr)oxide surface. The BZ-VV opal sample examined in this study has a U L3edge EXAFS spectrum closely matches that of the U/Fe/Si synthetic precipitate, suggesting that the U in the BZ-VV sample is associated with Fe inclusions (Figure 4). Zielenski15 postulated that this opal deposit formed from a solution with pH 7.0−8.5, as a gel precipitate from a Si-supersaturated solution. Uranyl undergoes near complete adsorption to iron oxides in this pH range. In contrast to BZ-VV, the OB-OR “mixed” opal has an EXAFS spectrum that resembles that of the U/Si synthetic precipitate, despite the presence of iron-rich inclusions within the opaline material (Figure 4). This suggests that the U in this opal is associated primarily with Si. Indeed, the clear (hyalite) portions of the OB-OR opal have greater fluorescence under an ultraviolet lamp, as compared to the fire opal portions. Uranium concentrations measured using SIMS suggest that U concentrations are low (∼3 mg kg−1 or less) and variable throughout the opaline material but that U concentrations are not especially enriched in the fire opal portion. The lack of association between U and Fe in the sample is possibly due to a lower pH of formation, as uranyl adsorbs minimally to iron (hydr)oxides below pH ∼3.5.11 Opal formation from an acidic, Si-saturated hydrothermal fluid would explain the molecular-scale U coordination environment inferred from the EXAFS spectrum of the OB-OR “mixed” opal. The retention of U(VI) within opal and amorphous silicate materials over million-year time scales, indicates that U incorporation into amorphous silica represents a stable storage

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ASSOCIATED CONTENT

S Supporting Information *

Supporting information contains sample descriptions, X-ray fluorescence maps of natural samples, X-ray powder diffractograms, normalization information for EXAFS spectra, “shell-byshell” fitting of EXAFS spectra from structural models discussed in the main text, a comparison of the EXAFS spectrum of uranyl adsorbed on ferrihydrite with the EXAFS spectrum of the synthetic U/Fe/Si solid, and a comparison between the EXAFS spectrum of the synthetic U/Si solid and that of a U EXAFS spectrum taken from the Opal Butte opal. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Support for M.M. was provided partially by the Robert and Marvel Kirby Stanford Graduate Fellowship. Additionally, this research was supported by the U.S. Department of Energy Office of Biological and Environmental Research, through the Subsurface Biogeochemical Research program (grant number DE-SC0006772) and the SLAC Science Focus Area Research Program (FWP #10094). This research was also supported by grants from the National Science Foundation (EAR-0921134 and EAR-1321511) to K.M. Use of the Stanford Synchrotron Radiation Lightsource, SLAC National Accelerator Laboratory, is supported by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences under Contract No. DE-AC02-76SF00515. The authors would like to thank Moses Gonzalez, Perach Nuriel, Matt Cobble, and Guangchao Li for their tireless assistance. The authors also appreciate the technical and safety support provided by L. Amoroso, D. Day, A. Gooch, B. Kocar, R. Marks, A. Mehta, D. Menke, C. Miller, C. Morris, D. Murray, C. Patty, R. Russ, and S. Webb.

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