Environ. Sci. Technol. 2004, 38, 5535-5539
Inhibited Cr(VI) Reduction by Aqueous Fe(II) under Hyperalkaline Conditions Y . T H O M A S H E , * ,† CHIA-CHEN CHEN,‡ AND SAMUEL J. TRAINA§ Environmental Science Graduate Program, The Ohio State University, Columbus, Ohio 43210, Department of Civil Engineering, University of Texas, Austin, Texas 78712, and Sierra Nevada Research Institute, University of California, Merced, California 95344
This study investigated Cr(VI) reduction by dissolved Fe(II) in hyperalkaline pH conditions as found in fluid wastes associated with the U.S. nuclear weapons program. The results show that Cr(VI) reduction by Fe(II) at alkaline pH solutions proceeds very quickly. The amount of Cr(VI) removed from solution and the amount reduced increases with Fe(II):Cr(VI) ratio. However, the Cr(VI) reduction under alkaline pH condition is nonstoichiometric, probably due to Fe(II) precipitation and mixed iron(III)-chromium(III) (oxy)hydroxides blocking Fe(II) surface sites, as well as removing Fe(II) from solution through O2 oxidation. After Cr(VI) was reduced to Cr(III), it precipitated out as mixed FexCr1-xO3(solids) and various Fe(III) precipitates with an overall Cr:Fe ratio of 1:3; all Cr remaining in the solution phase was unreduced Cr(VI). EXAFS data showed that Cr-O and Cr-Cr distances in the precipitates equal to 1.98 and 3.01 Å, respectively, consistent with the spinel-type structure as chromite.
estimated Cr concentration at the time of leakage ranged from 5.09 × 10-2 to 4.13 × 10-1 mol/L (3), and approximately 2685 kg of Cr as Cr(VI) was lost to the vadose zone in the S-SX Hanford tank farms (4). Inasmuch as Cr(VI) is poorly sorbed to mineral surfaces at circumneutral pH and above and toxic to biota, these Cr(VI) discharges posed a potential threat to groundwater resources and the adjacent Columbia River. Zachara et al. (4) examined sediment samples recovered from a HLW plume beneath tank SX-108. These materials exhibited extensive alteration of mineral surfaces near-field to the tank; high concentrations of Na+, NO3-, and 137Cs+; and total Cr levels ranging from 225 to 2098 mg/kg. Chromium K-edge, X-ray absorption near-edge structure (XANES) spectroscopy indicated that 29-75% of the total Cr was reduced to Cr(III) with the remainder as Cr(VI). Chromium reduction was attributed to OH--induced dissolution of Fe(II)-containing minerals (e.g., biotite, magnetite, chlorite, hornblende) and subsequent reaction of dissolved Fe(II) with Cr(VI). However, direct evidence of Cr(VI) reduction by dissolved Fe(II) at high pH was lacking. Homogeneous reduction of Cr(VI) by Fe(II) is welldocumented for the pH range of 1-12 (5-10). Reduction of Cr(VI) under acidic conditions is attributed to the following:
Fe2+ + H2CrO4 + H+ S Fe3+ + H3CrO4
(1)
Fe2+ + H3CrO4 + H+ S Fe3+ + H4CrO4
(2)
Fe2+ + H4CrO4 + H+ S Fe3+ + Cr(OH)2+ + OH- + H2O (6, 11) (3) While the reaction mechanisms are less certain at pH 8, the first step in chromate reduction by Fe(II) may be generalized as
FeOH+ + CrO42- S Fe(OH)3 + Cr(V)
(4)
and
Introduction Chromium is a common subsurface contaminant at the U.S. Department of Energy’s (DOE) site in Hanford, WA. The reduction-oxidation (REDOX) processes used excess Na2CrO7 in the recovery of Pu from spent nuclear fuel. Highlevel radioactive waste (HLW) fluids from the REDOX process and other extraction and purification processes conducted at Hanford were stored in 177 underground tanks during the period of 1944-1990. A total of 149 out of the 177 waste tanks were constructed with single steel walls and are commonly referred to as single shell tanks (SST). Sixty-seven of the SSTs are known or thought to have leaked, allowing from 1 920 000 to 3 456 000 L of HLW to migrate into the underlying vadose zone (1). Residual HLW fluids from REDOX were discharged to the S-SX Hanford tank farms. Several of the S and SX tanks leaked fluids to the subsurface. These fluids contained large quantities of dissolved NaOH, NO3-, NO2-, and Al3+ as well as pH values >13 and were at temperatures well in excess of 50 °C (2). The high temperatures were caused by the decay of fission products. The * Corresponding author present address: Dept. of Geosciences, 4044 Derring Hall, Virginia Tech, Blacksburg, VA 24060; phone: (540)231-1996; fax: (540)231-3386; e-mail:
[email protected]. † The Ohio State University. ‡ University of Texas. § University of California. 10.1021/es049809s CCC: $27.50 Published on Web 09/17/2004
2004 American Chemical Society
Fe(OH)20 + CrO42- S Fe(OH)3 +Cr(V) (6, 8, 11, 12) (5) While it is clear that reduction of Cr(VI) by Fe(II) can occur in acid to alkaline solutions, no information is available on this reaction at pH values >12, yet such porewater pH values may have been common during the discharge of HLW fluids from the SX-108 tank. Rai et al. (13) have observed the enhancement of Cr(III) solubility under highly alkaline conditions through the formation of Cr(OH)4- species. However, in systems where the components of a solid solution are homogeneously mixed and fit readily into a crystalline structure without tending to be expelled, they will have lower equilibrium solution activities than pure solid phases. This is because component activities in such a solid solution are less than unity. The Cr3+ and Fe3+ ions meet the crystalline structure criteria, having the same charge and nearly the same ionic radii of 0.63 and 0.64 Å, respectively. Therefore, Cr is less soluble when coprecipitated with Fe (14-16), and the solubility decreases with decreasing mole fraction of Cr(III) in the solid solution(17). Thus, it is not clear what impact high pH fluids would have on the solubility of Cr(III), if Fe(II) is present. The present study examines the effects of NaOH, NaNO3, and Na2CO3 on the reduction of Cr(VI) by dissolved Fe(II), simulating some of the conditions that may have existed within the HLW plume below SX-108. Na2CO3 was added to VOL. 38, NO. 21, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
5535
TABLE 1. Experimental Conditions for Each Treatment -1 (mol L-1) Fe:Cr (mmol L ) treatment ratio Cr(VI) Fe(II) NaNO3 Na2CO3
1 2 3 4 5 6 7 8 9
1:1 2:1 3:1 6:1 12:1 3:1 3:1 3:1
1 1 1 1 1 1 1 1
3 1 2 3 6 12 3 3 3
pH 13.5 in all treatments
1 1 1
0.001 0.01
some samples to simulate reaction of the alkaline solutions with CO2(g) present in the subsurface pores. In all cases, Cr(VI) reduction was rapid but not always complete. The reduction of Cr(VI) and side reactions of Fe(II) led to the formation of iron(III) (hydr)oxides and mixed Fe(III)-Cr(III) precipitates.
Materials and Methods All experimental solutions were prepared using reagent grade chemicals: FeCl2 (Alfa Aesar, 98%), NaOH (5N volumetric solution, J. T. Baker), Na2CrO4 (Aldrich, 98%), NaNO3 (Fisher, crystal, 99.2%), Na2CO3 (J. T. Baker, 99%). Deionized water was boiled while being purged with Ar(g), cooled under Ar(g), and then stored in an airtight plastic container in an Ar(g)-filled glovebox. All reagents were mixed, and the experimental solutions were added to reaction vessels while in the presence of an Ar(g) atmosphere. Samples for Cr(VI) reduction were prepared by adding aliquots of Na2CrO4 and FeCl2 to NaOH solutions in polyethelyne centrifuge tubes in the presence or absence of dissolved NaNO3 and Na2CO3. The concentration of NaOH was 1 mol L-1; total dissolved Cr(VI) was 1 mmol L-1; dissolved Fe(II) ranged from 0 to 12 mmol L-1; NaNO3 was either 0 or 1 mol L-1; and dissolved Na2CO3 was either 0, 1, or 10 mmol L-1 at the initiation of the experiments (see Table 1 for details). Reaction progress was monitored by measuring Fe(II) and Cr(VI) concentration change 1, 6, 24, and 72 h after initiation of reaction. The reacted suspension was centrifuged to separate the solid and solution phase. After separation, each solid phase was dried under Ar atmosphere to avoid exposure to air before XRD and SEM analysis. XAS spectra were collected using wet pastes after centrifugation and mounted to sample holders inside the Ar glovebox and stored in sealed bag under highly moisture conditions. The concentration of aqueous Cr(VI) was measured using a diphenylbarbazide method (18). This method is insensitive to Cr(III). Total Cr and Fe were analyzed using inductively coupled plasma-optical emission spectroscopy (Perkin-Elmer Optima 3000). Solution pH was measured with an accuFet solid-state electrode, calibrated with standards at pH 10 and pH 12. Powder X-ray diffraction (XRD) patterns of solid-phase reaction products were collected using Cu KR radiation and a Philips diffractometer, from 2 to 70° 2θ, with 4-s steps and a step interval of 0.05°. Mineral phases in the precipitate were identified by matching with standards in the ICDD Powder Diffraction File 1993. Transmission electron micrographs (TEM) were collected with a CM12 Philips microscope operated at 60 keV. The samples were prepared in a diluted aqueous suspension and then added as 1-2 drops to Formvar-coated TEM grids and air-dried. Scanning electron microscopy (SEM) and energy dispersive spectroscopy (EDS) data were collected with a JEOL JSM5536
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 38, NO. 21, 2004
FIGURE 1. Fe(II) vs Cr(VI) removal/reduction at pH 13.5. 820 SEL at 15 keV. Chemical analyses were obtained with an Oxford EXL energy dispersive X-ray analyzer. Synchrotron-based X-ray absorption spectroscopy (XAS) measurements were performed at the Stanford Synchrotron Radiation Laboratory (SSRL) on beam lines IV-3 and XI-2, which are equipped with a Si(220) double-crystal monochromator. Fluorescence yield Cr K-edge XAS data were collected using a Canberra Ge 13-element energy-dispersive detector. Higher-order harmonics in the incoming beam were removed with a Pt-coated mirror and by detuning the main beam by 30-50%. Energy was calibrated by collecting the transmission spectrum of a Cr metal foil. The first inflection point of the K-edge of the Cr metal foil was assigned as 5989 eV. XAS spectra were collected over the energy range of 5.76.5 keV. Averaging, normalization, and background subtraction of the raw XAS spectra were performed with the computer program EXAFSPAK (19). Each averaged XAS spectrum was separated into two regions, the X-ray absorption near-edge structure (XANES) spectral region and the extended XAFS (EXAFS) spectral region. XANES spectra (5980-6100 eV) were normalized and compared for qualitative information. EXAFS oscillations were isolated with a spline function and converted from energy to k-space. A spline was fit by using the computer program IFEFFIT (20). The converted EXAFS was then exported back to the EXAFSPAK program for Fourier transform and curve-fitting procedures. The k3-weighted EXAFS spectra of the model compounds sodium chromate and chromium(III) oxide were fit (k-range ) 3-12 Å-1) with phase shift and amplitude functions generated by using the computer code, FEFF8 (21). These fitting results were used to test the theoretical phase shift and amplitude functions that were later used to fit unknown samples. Values for coordination number (CN) and distance to scattering atoms (R) were determined from least-squares fits of the EXAFS and the Fourier-filtered EXAFS of each shell.
Results and Discussion Cr(VI) Reduction under Alkaline Conditions. The reaction of 3 mmol L-1 Fe2+ and 1 mol L-1 NaOH under Ar(g) (in absence of Cr(VI)) yielded greenish precipitates after about 1 d. Upon exposure to air, these immediately became reddishbrown. Subsequent examination with powder XRD detected only goethite in these samples. The addition of Cr(VI) to Fe(II) solutions at pH g 13.5 resulted in virtually instantaneous formation of red-brown precipitates, partial to complete removal of Cr (Figure 1), and complete removal of Fe from solution. TEM micrographs of the precipitates showed them to be micrograined (Figure 2). Some colloids formed loose aggregates and seemed to be composed of even finer units. At higher initial Fe(II):Cr(VI) ratios (6:1 to 12:1), needle-shaped aggregates also formed (Figure 3). Powder XRD patterns (Figure 4) indicated the presence of three different solids, akaganeite (β-FeO(OH)), iron-chromium oxide ((Fe1-xCrx)2O3), and iron (hydr)oxide (FeOOH), whose relative amounts varied with the initial Fe-
FIGURE 4. XRD pattern of precipitates in different Cr(VI):Fe(II) ratios. (a) 3 mmol L-1 Fe(II) in 1 mol L-1 NaOH solution. (b) Fe(II):Cr(VI) ) 1:1. (c) Fe(II):Cr(VI) ) 2:1. (d) Fe(II):Cr(VI) ) 3:1.
FIGURE 2. TEM showing precipitates formed at low Fe(II):Cr(VI) ratios (1:1 to 3:1).
FIGURE 5. XRD showing chromite formation from high Fe(II):Cr(VI) ratios at alkaline pH conditions.
TABLE 2. Composition of Precipitates Formed in the Homogeneous System as Measured by SEM/EDS sample 2 (Fe:Cr ) 1:1) 4 (Fe:Cr ) 3:1) 4 (Fe:Cr ) 3:1, aging sample)
FIGURE 3. TEM showing precipitates formed at high Fe(II):Cr(VI) ratios (6:1 to 12:1), EDX showed that needle shaped aggregates are iron oxides. (II) to Cr(VI) ratio. It appears that the relative peak intensities of the FeOOH polymorphs increased with increases in the initial solution concentration of Fe(II) in all of the samples reacted with Cr(VI). In the higher Fe(II):Cr(VI) treatments (6:1 and 12:1), a reflection near that of chromite (FeIICrIII2O4) was detected with powder XRD (Figure 5). Solids with a composition of (Fe0.75Cr0.25(OH)3) have been observed during the homogeneous reduction of Cr(VI) by acidic Fe(II)(aq) (22). In contrast, the heterogeneous reduction of Cr(VI) by SO4-, CO3- and Cl- green rust produced reaction products with Fe(III):Cr(III) ratios of 1:5, 1:1, and 2:1, respectively (17). Patterson et al. (23) found that chromate reduction on the FeS surface and resulted in a solid phase with a composition of Fe0.25Cr0.75(OH)3. Energy-dispersive
O (%)
Cr (%)
Fe (%)
Fe:Cr ratio
53.49 49.36 72.14
3.08 2.34 5.35
10.03 7.21 15.20
3.26 3.08 2.84
74.00 65.09 64.61 58.09
3.49 0.90 3.04 1.29
12.65 2.89 10.96 4.57
3.62 3.21 3.61 3.54
X-ray measurements (SEM based) indicated that the overall Fe:Cr ratio of the precipitates formed in the present study was approximately 3:1 (Table 2), consistent with other studies of homogeneous reduction. Whereas, the extent of Cr sorption by the precipitates increased with the amount of Fe(II) added, differences in ionic strength or concentration of dissolved CO32- (added as Na2CO3) did not affect Cr sorption (data not shown). Colorimetric solution analysis indicated that all soluble Cr was present as Cr(VI). Thus, there was no evidence of soluble Cr(III) hydrolysis products at high pH as have been reported by Rai et al. (13). XANES spectra of the precipitates from these reactions lacked the characteristic pre-edge feature that is diagnostic of Cr(VI), indicating reduction to Cr(III) (Figure 6). Clearly, reaction of Fe(II) with Cr(VI) in 1 mol L-1 NaOH solutions VOL. 38, NO. 21, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
5537
TABLE 3. Cr EXAFS Fitting Parametersa sample
bond
N
R
σ2
Na2CrO4 2 Fe(II):Cr(VI) ) 1:1 4 Fe(II):Cr(VI) ) 3:1
Cr-O Cr-O Cr-Cr/Fe Cr-O Cr-Cr/Fe
4b 4.59 2.70 4.40 2.80
1.66 1.98 3.00 1.98 3.01
0.0031 0.0026 0.0080 0.0026 0.0091
a N is the number of backscattering atoms around the absorbing Cr atom. R is the absorber-backscatterer distance. σ2is the Debye-Waller value. b N values were fixed at known values for model compound Na2CrO4. Estimated errors for N ( 20%, R ( 0.01 based on least-squares fits of EXAFS spectra. Error represents 95% confidence interval.
FIGURE 6. Cr XANES spectra of precipitates formed in the homogeneous system. (a) Fe(II):Cr(VI) ) 1:1. (b) Fe(II):Cr(VI) ) 2:1. (c) Fe(II):Cr(VI) ) 3:1. (d) Cr(VI) model compound. led to the reduction of Cr(VI) to Cr(III) and association of Cr(III) with the resulting solid phase(s). Manceau and Charlet (24) measured a Cr(VI)-O distance of 1.62 Å and a Cr(III)-O distance of 1.99 Å in studies of Cr(III) oxidation by manganese oxides. EXAFS studies of Cr(III) sorption on hydrous ferric oxide revealed that Cr(III) was surrounded by three metal (Fe or Cr) shells at 3.00-3.05, 3.40-3.46, and 3.94-4.03Å (24). When coprecipitated with Fe(III), only two metal shells were detected around Cr(III) at 2.99 and 3.40 Å, indicating that Cr substituted for Fe in an R-(Fe,Cr)OOH framework. In fitting Cr(VI)- and Cr(III)containing model compounds, Peterson et al. (25) reported a Cr(VI)-O distance of 1.63 ( 0.03 Å and a Cr(III)-O distance of 1.99 Å. The reaction of Cr(VI) with magnetite yielded a Cr-O distance of 1.98 ( 0.01Å. Second shell distances between Cr and Cr or Fe atoms were 3.01 ( 0.01 Å. Hansel et al. (26) observed Cr(III)-O distances of 2.01 Å, second shell Cr-Cr(Fe) distances of 3.04-3.06 Å, and third shell Cr-Cr(Fe) distances of 3.26-3.30 Å for Cr(VI) reacted with green rust. Analysis of EXAFS spectra collected in this study (Figure 7 and Table 3) indicated that precipitates formed in systems with initial Fe(II):Cr(VI) ratios of 1:1 and 3:1 had essentially the same local structures. The first shell cor-
responded to a Cr-O bond at 1.99 Å, and the second shell corresponded to a Cr-Cr/Fe bond at 3.01 Å. In the sodium chromate model compound, the Cr-O bond distance was 1.66 Å. Stoichiometry and Kinetics. It is well-known that reduction of Cr(VI) by Fe(II) involves three one-electron-transfer steps:
Fe(II) + Cr(VI) S Fe(III) + Cr(V)
(6)
Fe(II) + Cr(V) S Fe(III) + Cr(IV)
(7)
Fe(II) + Cr(IV) S Fe(III) + Cr(III)
(8)
and
The overall stoichiometry of these reactions is 3:1, with 3 mol of Fe(II) required to reduce 1 mol of Cr(VI). Unlike the stoichiometric reduction observed in acidic to neutral solutions, Fe(II) reduction of Cr(VI) in this study did not exhibit an apparent 3:1 stoichiometry. Only 50% of the potential Cr(VI) reduction was observed at initial dissolved Fe(II) to Cr(VI) ratios of 1:1 to 3:1. Complete Cr(VI) reduction and removal of Cr from solution only occurred at initial Fe(II) to Cr(VI) ratios of 6:1 and 12:1. Eary and Rai (10) reported similar results at pH > 10. It is possible that incomplete reduction of Cr(VI) by Fe(II) was due to the removal of Fe(II) from solution as a solid. As noted above, greenish Fe precipates formed in the Cr-free control samples with complete removal of dissolved Fe after 1 d of reaction; however, the identity of these solids is not known. Exposure
FIGURE 7. Cr EXAFS spectra. Solid line is the data, and dotted line is the nonlinear least-squares fit. (a) Cr EXAFS; (b) Fourier transform. 5538
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 38, NO. 21, 2004
to air prior to XRD analyses resulted in a change in color, from green to yellow. Thus, the actual nature of initial precipitates in the absence (and perhaps in the presence of Cr(VI)) is not known. Fe(II)-containing solids can facilitate the heterogeneous reduction of Cr(VI), but the surfaces are quickly passivated by Cr(III) and Fe(III) reaction products, rendering some of the solid-phase Fe(II) unreactive (2729). Thus, if Fe(II) solids were present, surface passivation could have led to the apparent nonstoichiometric reduction of Cr(VI) by Fe(II). Oxidation of dissolved Fe(II) by trace amounts of dissolved O2 and subsequent precipitation of Fe(III) solids could also have inhibited Cr(VI) reduction. Even though these experiments were conducted under an Ar(g) atmosphere, the presence of trace quantities of O2 cannot be ruled out. In acidic solutions the oxidation of dissolved Fe(II) by O2 is a slow reaction, but Fe(II) oxidation has a half-life of a few seconds at pH values greater than 7.5 (30). Indeed, at pH > 10.0 and in alkaline phosphate solutions, the competitive oxidation of aqueous Fe(II) with dissolved oxygen resulted in an apparent nonstoichiometric reduction of aqueous Cr(VI) (10). At high pH the oxidizability of Fe(II) is great and O2 more efficient in oxidizing Fe(II) than Cr(VI) (31). The overall rate for oxidation of dissolved Fe(II) by O2 is 7-14 times slower than Fe(II) oxidation by Cr(VI) at pH 6, near equal at pH 7, and 10-20 times faster at pH 8 (6). Lin (32) showed that, at pH > 8.0 and in the presence of dissolved O2, Cr(VI) reduction by Fe(II) was greatly suppressed due to the rapid oxidation of Fe(II). Thus, even in the presence of dissolved Cr(VI), the presence of minor amounts of dissolved O2 could have resulted in the formation of Fe(III) solids, decreasing the amount of aqueous Fe(II). The reduction of Cr(VI) by Fe(II) is a relatively fast reaction in acidic to moderately basic solutions, taking from tens of seconds in a homogeneous system (7, 8, 10) to several hours/ days in heterogeneous systems (12) to reach completion. The reduction kinetics in heterogeneous systems are related to the kinetics of Fe(II) release to solution. Buerge and Hug (33) described the pH-dependent kinetics of Cr(VI) reduction by ferrous iron with a rate law based on Fe(II) speciation:
-d[Cr(VI)]/dt ) (k1[Fe2+] + k2[FeOH+] + k3[Fe(OH)20])[Cr(VI)] (9) where Fe2+, FeOH+, and Fe(OH)20 are the dominant species of Fe(II) in acidic to slightly basic solutions. The rapid reaction rates observed in the present study can be explained by increases in the concentrations of FeOH+, Fe(OH)20, and Fe(OH)3-. Eary and Rai (12) found that Fe(OH)20 is oxidized more rapidly than either Fe2+ or FeOH+. The increase in the rate of ferrous ion oxygenation at pH greater than 10.0 relative to oxidation by chromate coincides with the pH above which Fe(OH)3- becomes the dominant ferrous species (8, 12). Schlautman and Han(9) also observed that adding strong base can cause short-term, locally high pH conditions that may lead to enhanced Cr(VI) reduction kinetics in both the presence and absence of dissolved O2 conditions. While the present study did not formally measure the reaction rates, it was still apparent that Cr(VI) reduction by Fe(II) at pH g 13.5 was extremely rapid. The results of this work show that homogeneous reduction of Cr(VI) by dissolved Fe(II) is possible at pH values g 13.5, leading to the rapid removal of Cr from solution and incorporation of Cr(III) into solid-phase reaction products. The formation of iron precipitates limits the availability of Fe(II) for Cr(VI) reduction. The combination of XANES, EXAFS, and XRD data indicated that mixed Cr(III)-Fe(III) solids formed from this reaction, with structures similar to that of chromite. Further work is required to determine the
efficacy of this reaction in the presence of solid-phase forms of Fe(II) in minerals such as biotite and magnetite.
Acknowledgments The authors thank Dr. J. Bigham of The Ohio State University for his help with the XRD analysis. The authors also thank Dr. S. Chao and three anonymous reviewers whose comments significantly improved the manuscript. The study was funded by EMSP, Department of Energy (Grant DE-FG07-99ER15010).
Literature Cited (1) Hanlon, B. M. Waste Tank Summary Report for Month Ending January 31, 1996; Pacific Northwest Laboratory: Richland, WA, 1996. (2) Zachara, J. M.; Ainsworth, C. C; Brown, G. E. Geochim. Cosmochim. Acta 2004, 68, 13-30. (3) Qafoku, N. P.; Ainsworth, C. C.; Szecsody, J. E. Reactive Transport of Al-rich, Alkaling and Saline High Level Waste Simulants in the Hanford Sediments; Pacific Northwest Laboratory: Richland, WA, 2002. (4) Jones, T. E.; Watrous, R. A.; Maclean, G. T. Inventory Estimates for Single-Shell Tank Leaks in S and SX Tank Farms; Pacific Northwest Laboratory: Richland, WA, 2000. (5) Buerge, I. J.; Hug, S. J. Environ. Sci. Technol. 1997, 31, 14261432. (6) Pettine, M.; D’Ottone, L.; Campanella, L.; Millero, F. J.; Passino, R. Geochim. Cosmochim. Acta 1998, 62, 1509-1519. (7) Fendorf, S. E.; Li, G. Environ. Sci. Technol. 1996, 30, 1614-1617. (8) Sedlak, D. L.; Chan, P. G. Geochim. Cosmochim. Acta 1997, 61, 2185-2192. (9) Schlautman, M. A.; Han, I. Water Res. 2001, 35, 1534-1546. (10) Eary, L. E.; Rai, D. Environ. Sci. Technol. 1988, 22, 972-7. (11) Epsenson, J. H. J. Am. Chem. Soc. 1970, 92, 1880-1883. (12) Eary, L. E.; Rai, D. Am. J. Sci. 1989, 289, 180-213. (13) Rai, D.; Hess, N. J.; Rao, L. J. Solution Chem. 2002, 31, 343-367. (14) Powell, R. M.; Puls, R. W.; Hightower, S. K.; Sabatini, D. A. Environ. Sci. Technol. 1995, 29, 1913-1922. (15) Sass, B. M.; Rai, D. Inorg. Chem. 1987, 26, 2228-2232. (16) Seaman, J. C.; Bertsch, P. M.; Schwallie, L. Environ. Sci. Technol. 1999, 33, 938-944. (17) Bond, D. L.; Fendorf, S. Environ. Sci. Technol. 2003, 37, 27502757. (18) Fishman, M. J.; Friedman, L. C. Techniques of Water Resources Investigations of the United States Geological Survey; United States Geological Survey: Denver, 1989. (19) George, G. N.; Pickering, I. J. EXAFSPAK; Stanford Synchrotron Radiation Laboratory: Stanford, CA, 1995. (20) Newville, M. IFEFFIT Program; University of Chicago: Chicago, IL, 1997. (21) Ankudinov, A.; Ravel, B.; Rehr, J. J. FEFF8; University of Washington: Seattle, WA, 1999. (22) Buerge, I. J.; Hug, S. J. Environ. Sci. Technol. 1999, 33, 42854291. (23) Patterson, R. R.; Fendorf, S.; Fendorf, M. Environ. Sci. Technol. 1997, 31, 2039-2044. (24) Manceau, A.; Charlet, L. J. Colloid Interface Sci. 1992, 148, 425442. (25) Peterson, M. L.; Brown, G. E., Jr.; Parks, G. A. Colloids Surf., A 1996, 107, 77-88. (26) Hansel, C. M.; Wielinga, B. W.; Fendorf, S. Geochim. Cosmochim. Acta 2003, 67, 401-412. (27) Kendelewicz, T.; Liu, P.; Doyle, C. S.; Brown, G. E. Surface Sci. 2000, 469, 144-163. (28) Chambers, S. A. Molecular level processes governing the interaction contaminants with iron and manganese oxides; Pacific Northwest National Laboratory: Richland, WA, 1996. (29) Williams, A. B.; Scherer, M. M. Environ. Sci. Technol. 2001, 35, 3488-3494. (30) Kieber, R. J.; Helz, G. R. Environ. Sci. Technol. 1992, 26, 307312. (31) Lewis, D. G. Adv. GeoEcol. 1997, 30, 345-372. (32) Lin, C.-J. Water, Air, Soil Pollut. 2002, 139, 137-158. (33) Buerge, I. J.; Hug, S. J. Environ. Sci. Technol. 1998, 32, 20922099.
Received for review February 7, 2004. Revised manuscript received July 18, 2004. Accepted August 3, 2004. ES049809S VOL. 38, NO. 21, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
5539
