Catalyzed Oxidation of Arsenic(III) by Hydrogen Peroxide on the


Environmental Science & Technology 2016 50 (1), 214-221. Abstract | Full Text .... and David L. Sedlak. Environmental Science & Technology 0 (proofing...
0 downloads 0 Views 153KB Size


Environ. Sci. Technol. 2003, 37, 972-978

Catalyzed Oxidation of Arsenic(III) by Hydrogen Peroxide on the Surface of Ferrihydrite: An in Situ ATR-FTIR Study ANDREAS VOEGELIN AND STEPHAN J. HUG* Swiss Federal Institute for Environmental Science and Technology (EAWAG), U ¨ berlandstrasse 133, CH-8600 Du ¨ bendorf, Switzerland

Knowledge of arsenic redox kinetics is crucial for understanding the impact and fate of As in the environment and for optimizing As removal from drinking water. Rapid oxidation of As(III) adsorbed to ferrihydrite (FH) in the presence of hydrogen peroxide (H2O2) might be expected for two reasons. First, the adsorbed As(III) is assumed to be oxidized more readily than the undissociated species in solution. Second, catalyzed decomposition of H2O2 on the FH surface might also lead to As(III) oxidation. Attenuated total reflectionFourier transform infrared (ATR-FTIR) spectroscopy was used to monitor the oxidation of adsorbed As(III) on the FH surface in situ. No As(III) oxidation within minutes to hours was observed prior to H2O2 addition. Initial pseudofirst-order oxidation rate coefficients for adsorbed As(III), determined at H2O2 concentrations between 8.4 µM and 8.4 mM and pH values from 4 to 8, increased with the H2O2 concentration according to the equation log kox (min-1) ) 0.17 + 0.50 log [H2O2] (mol/L), n ) 21, r2 ) 0.87. Only a weak pH dependence of log kox was observed (∼0.04 logarithm unit increase per pH unit). ATRFTIR experiments with As(III) adsorbed onto amorphous aluminum hydroxide showed that Fe was necessary to induce As(III) oxidation by catalytic H2O2 decomposition. Supplementary As(III) oxidation experiments in FH suspensions qualitatively confirmed the findings from the in situ ATRFTIR experiments. Our results indicate that the catalyzed oxidation of As(III) by H2O2 on the surface of iron (hydr)oxides might be a relevant reaction pathway in environmental systems such as surface waters, as well as in engineered systems for As removal from water.

With respect to arsenic adsorption onto mineral phases, Goldberg (3) recently reported that arsenate adsorption on aluminum and iron oxides and clays was maximal at low pH and decreased with increasing pH, whereas arsenite adsorption on clays and aluminum and iron oxides was maximal around pH 8.5. Because the reduction and oxidation processes of arsenic are rather slow, both oxidation states can occur irrespective of the redox conditions (4). Detailed knowledge of arsenic redox transformation kinetics is needed to understand the biogeochemical cycling of arsenic, to assess its impact and remediation in contaminated systems, and to optimize drinking water treatment for arsenic removal. Under oxic conditions, microorganisms are known to rapidly oxidize As(III) (5, 6). Arsenic oxidation, however, can also be promoted inorganically. Manganese oxides are wellknown to effectively oxidize adsorbed As(III) (7-9). As(III) oxidation has also been observed on several clay minerals (10, 11). Whereas some researchers report oxidation of adsorbed As(III) on goethite and amorphous iron oxides (12, 13), others found As(III) adsorbed on iron (hydr)oxides to be stable toward oxidation (7, 14). In aqueous solution, Pettine et al. (15) observed highly pH-dependent rates for As(III) oxidation by hydrogen peroxide. The steep increase in the reaction rate coefficient with pH was explained by the difference in reactivity between the unreactive H3AsO3 and reactive deprotonated species. The observed reaction rates suggest that this process might be relevant in surface waters with elevated H2O2 concentrations (1-10 µM) and alkaline pH’s or treatment systems for contaminated solutions with millimolar H2O2 concentrations (16, 17). The rate of H2O2 decomposition and concomitant As(III) oxidation can be enhanced by the presence of metal cations such as Cu2+ or Fe3+ (18). Iron (hydr)oxide surfaces catalyze the decomposition of H2O2 (19-21) and can enhance the degradation of organic compounds (19, 22). These findings suggest that iron oxide surfaces might catalyze the oxidation of As(III) by H2O2 for two reasons. First, innerspherically adsorbed arsenite might correspond more closely to deprotonated arsenite than to undissociated H3AsO3 and hence might be more reactive toward oxidation by H2O2. Second, iron (hydr)oxide surfaces catalyze the decomposition of H2O2 and possibly also concomitant As(III) oxidation. The objective of the present study, therefore, was to investigate the oxidation of As(III) adsorbed on ferrihydrite in the presence of H2O2. Ferrihydrite (FH) was chosen as a model for natural amorphous iron oxides. Experiments were conducted using attenuated total reflectance-Fourier transform infrared (ATR-FTIR) spectroscopy, which allows for the study of kinetic surface reactions in situ (23-25). Complementary experiments were conducted with FH suspensions to support the findings from spectroscopy.

Introduction

Experimental Section

The behavior of arsenic in the environment and in water and wastewater treatment systems strongly depends on its oxidation state. Arsenic mainly occurs as the As(III) and As(V) oxyanions arsenite and arsenate, respectively, which largely differ in their chemical behavior. As(III) is considered to have a higher acute toxicity than As(V) (1). In most natural water systems, arsenite [pK1 ) 9.1, I ) 0.1 M (2)] occurs as a neutral undissociated species (H3AsO3). In contrast, arsenate [pK1 ) 2.8, pK2 ) 6.7, pK3 ) 11.8, I ) 0.1 M (2)] mainly occurs as singly (HAsO4-) and doubly (H2AsO42-) protonated anions.

Synthesis of Ferrihydrite. Two-line FH was synthesized following the method of Schwertmann and Cornell (26). Briefly, 0.2 M Fe(NO3)3‚9H2O was hydrolyzed with 1 M KOH to a pH of 8. The suspension was dialyzed at 4 °C in the dark until the conductivity dropped below 1.5 µS/cm. The final suspension contained ∼9.7 g/L FH and had a pH of ∼6.8. X-ray diffraction analysis showed the typical two-line FH diffractogram with no secondary lines from goethite or hematite. The surface area of the freshly synthesized FH as determined from an N2 BET adsorption isotherm at 25 °C was ∼260 m2/g, in agreement with the range given by Schwertmann and Cornell (26). The FH suspension was stored at 4 °C in the dark.

* Corresponding author phone: +41-1-823-5454; fax: +41-1-8235210; e-mail: [email protected] 972

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 5, 2003

10.1021/es025845k CCC: $25.00

 2003 American Chemical Society Published on Web 02/04/2003

FTIR Measurements. Spectra were recorded on a Biorad FTS 575C instrument equipped with a mercury cadmium telluride (MCT) detector and a nine-reflection diamond ATR unit with KRS-5 optics (SensIR Technologies, Danbury, CT). The diameter of the round probing area of the diamond was 4 mm. Scans from 400 to 4000 cm-1 were taken at 4 cm-1 resolution versus the appropriate background spectrum. Data analysis was performed with Matlab (The MathWorks, Inc). Spectra of As(V) and As(III) Species in Solution. Stock solutions of 40 mM As(V) and As(III) were prepared from Na2HAsO4‚7H2O and NaAsO2, respectively. Aliquots were adjusted to pH values between 3 and 12 with 1 M KOH and 1 M HCl. To collect the infrared spectra, 20-µL samples of these solutions were directly applied onto the ATR crystal. The background spectrum was collected in doubly distilled water. Spectra were averaged from 256 scans. Preparation of Ferrihydrite Layers for ATR-FTIR Spectroscopy. The ATR crystal surface (area ) 12.6 mm2) was coated with 12 µg of FH by spreading 5 µL of a diluted suspension containing 2.4 g/L FH on the crystal. The film was gently dried under a N2 stream and appeared homogeneous upon visual inspection. However, scanning electron microscopy and atomic force microscopy of FH films analogously deposited on a glass surface showed that, on the micrometer scale, the FH was not homogeneously deposited. Film heights ranged between 0 and 4 µm, with patches of up to 10 µm diameter of uncoated glass. The average film height was ∼1 µm. This compares to an average film height of ∼0.8 µm estimated from a FH density of ∼4000 kg/m3 and an average porosity of ∼70% for the deposited film. After being dried, the film was rinsed with 1 mL of a 10 mM KCl solution and dried again under N2. A comparison of spectra of the dried film collected before and after the washing procedure showed that usually less than 5% of the applied FH had been removed by rinsing and that the film was stable in contact with aqueous solutions. Incremental FH addition up to ∼12 µg gave proportional IR absorbances, indicating that the entire FH layer thickness was probed by the evanescent IR light. This is in agreement with the calculated penetration depth of 2.1 µm at 800 cm-1 (with ndiamond ) 2.4, 45° internal reflection angle, and an assumed refractive index of n ) 1.4 for the FH layer). The average absorbance of the FH films between 555 and 565 cm-1 was 0.62, with a relative standard deviation of (12% (n ) 23), indicating fairly good reproducibility. Spectra of As(V) and As(III) Adsorbed on Ferrihydrite. The background spectrum of the FH film in contact with aqueous solution was collected after the addition of 5 µL of 10 mM KCl solution to the dry film. Then, small volumes of As(V) or As(III) stock solutions, i.e., 0.5-5 µL of 1 mM As(V) or As(III), were pipetted onto the FH film, corresponding to 0.5-5 nmol of As or 0.04-0.4 mol of As per kilogram of FH. The As solutions were allowed to equilibrate for 10-15 min. Because of the high initial As concentrations, the As was almost quantitatively adsorbed to the FH film. Infrared spectra were averaged from 32 scans. In Situ Arsenic Oxidation Experiments on the ATR Crystal. All oxidation experiments were conducted under air. One group of experiments under well-controlled pH conditions was performed in a liquid cell consisting of a 50-mL polypropylene beaker with a 4-mm opening at the bottom surrounded by an O-ring seal that was tightly pressed onto the ATR device. Either 1 or 3 nmol of As(III) (0.083 or 0.250 mol/kg, respectively) was quantitatively adsorbed onto the FH film as described above. After equilibration, 20 mL of 10 mM KCl were added over the oxide film and gently stirred. The pH (6.5 and 7.9) was adjusted with 0.02 M HCl and KOH by a computer-controlled titration system, which took between 20 and 45 min to reach a stable pH value. H2O2 (84 mM stock solution) was then added to give final

concentrations of 42 and 168 µM. IR spectra were automatically collected every 2 min. The influence of different modes of As(III) application was investigated in two ways. First, in one experiment (pH 6.5, 42 µM H2O2), 0.250 mol of As(III) per kilogram of FH was spiked into the FH suspension prior to its deposition on the ATR crystal. Second, in two experiments (84 µM H2O2 at pH 5.8, 2100 µM H2O2 at pH 5.7), the initially As(III)-free FH film was equilibrated for 1 h with 50 mL of 1 µM As(III) in 10 mM KCl background, and the amount of As(III) adsorbed was calculated from the decrease of As(III) in solution. The above experimental setup allowed for the close control of the solution pH but led to two experimental difficulties. First, the system was prone to contamination with nanomolar amounts of silicate and phosphate, both of which interfered with the arsenate IR spectrum. Traces of silicate could leach from the glass buret containing the 0.02 M KOH solution and phosphate (originating from the electrode buffer solution) from the electrode diaphragm. Second, because of the large solution volume (20 mL) in contact with the thin FH layer (12 µg), between 35 and 65% of the initially adsorbed arsenic desorbed into the solution within 2-3 h (resulting in total As concentrations of 6 is far lower than 1 µM, and the slight increase in the As oxidation rate with pH runs contrary to the solubility of FH. Our first hypothesis for catalyzed As(III) oxidation on the FH surface was that the adsorbed As(III) resembles more closely the deprotonated than the undissociated As(III) in solution. The observed small pH dependence of the surface oxidation rate coefficient (eq 4) would have been in line with this hypothesis, assuming that the adsorbed As(III) complex did not change between pH 4 and 8. To test this hypothesis, complementary in situ As(III) oxidation experiments were conducted with an amorphous aluminum hydroxide (AH). Spectra from these experiments are presented in Figure 5. As on the FH surface, the ATR-FTIR method allows As(V) and As(III) adsorbed onto the AH to be clearly differentiated (Figure 5). The spectrum of adsorbed As(V) compares well with the spectrum published by Goldberg and Johnston (28). In contrast to the experiments with FH, no oxidation of 0.3 mol/kg As(III) occurred on the AH surface after the addition of 8.4 mM H2O2 (Figure 5). However, after the addition of 0.09 mol of Fe(III) per kilogram of AH and a waiting period of 10 min, subsequent addition of H2O2 again led to As(III) oxidation (Figure 5). As with FH, no oxidation was observed 976

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 5, 2003

Combining the rate of H2O2 decomposition (eq 6) with the rate of As(III) oxidation (eq 3b), the molar ratio, R, of As(III) oxidation to H2O2 decomposition can be estimated as follows

R ) [FH] × [tAs(III)]/dt × (d[H2O2]/dt)-1 ) 0.015 (mol/L)0.5 (mol/kg)-1 × [tAs(III)] × [H2O2]-0.5 (7) With increasing H2O2 concentration and decreasing adsorbed As(III) concentration, the molar ratio of As(III) oxidation to H2O2 decomposition decreases. For [tAs(III)](t)0) ) 0.25 mol/ kg and at H2O2 concentrations above 14 µM, more than one H2O2 molecule is expected to be decomposed per As(III) ion oxidized. This can be rationalized in terms of the reaction mechanisms proposed in Table 1. The slow and rate-limiting reaction of H2O2 with tFe(III) (t denotes surface-bound species) initiates the chain reactions (R2-R7) that decompose H2O2 and oxidize As(III). A steady state between chaininitiating (R1) and chain-terminating reactions (R8 and R9) is quickly established after the addition of H2O2. At high H2O2 concentrations, H2O2 competes with tAs(III) in the reaction with tFe(III)•OH (R3 and R6) and the tAs(IV) intermediate is increasingly oxidized by H2O2 (R5) rather than by O2 (R4), with both processes leading to increased H2O2 degradation in relation to As(III) oxidation. At low H2O2 concentration, however, the arsenic-oxidizing reactions (R3 and R4) might become dominant over the H2O2-consuming chain reactions. The oxidation kinetics were not significantly affected by the addition of 2-propanol (Figure 4), which would have quantitatively scavenged •OH radicals from the solution (39). This indicates that the oxidation occurred with surface-bound hydroxyl radicals or, alternatively, with a higher-valent iron intermediate that is not included in the current scheme [e.g., tFe(III)•OH could be formulated as tFe(IV)]. Note that the reactions listed in Table 1 represent a preliminary working hypothesis that cannot be confirmed from the results from this study alone. Arsenic Oxidation Experiments in Suspension. With ATR-FTIR spectroscopy, the rate of the oxidation of adsorbed As(III) on the FH surface could be determined directly and in situ. In complementary batch experiments with FH suspensions, removal rates of dissolved As(III), which depend on both the rates of As(III) adsorption and oxidation, were determined. Upon addition of 10 or 100 mg/L FH to solutions containing 10 or 60 µM As(III), respectively, rapid adsorption of As(III) was observed within 15 min. At the adsorption equilibrium, the FH was loaded with ∼0.5 mol/kg As(III) and the As(III) concentration in solution was ∼5 µM, corresponding to a distribution coefficient Kd [sorbed arsenite (in moles per kilogram) to arsenite in solution (in moles per

TABLE 1. Possible Reaction Mechanism for the Catalyzed As(III) Oxidation by H2O2 on the Ferrihydrite Surface reactiona

equationb,c

ref

R1 (i) R2 (c) R3 (c) R4 (c) R5 (c) R6 (c) R7 (c) R8 (t) R9 (t) R10 (t) R11

tFe(III) + H2O2 f tFe(II) + O2•- + 2H+ (10-3) tFe(II) + H2O2 f tFe(III)•OH + OH- (102) tFe(III)•OH + tAs(III) f tAs(IV) + tFe(III) + OH- (109) tAs(IV) + O2 f tAs(V) + O2•- (109) tAs(IV) + H2O2 + tFe(III) f tFe(III)•OH + tAs(V) +OH- (109) tFe(III)•OH + H2O2 f tFe(III) + O2•- + H2O + H+ (107) tFe(III) + O2•- f tFe(II) + O2 (107) tAs(IV) + tFe(III)•OH f tAs(V) + tFe(III) (109) tFe(II) + tFe(III)•OH f 2 tFe(III) + OH- (108) O2•- + HO2•+ H+ f H2O2 + O2 (107) O2•- + H+ T HO2•(pKa 4.8)

40 41 42 42 42 43 41 42 44 45 45

a (i) chain-initiating reaction, (c) chain-propagating reactions consuming H O and oxidizing As(III,IV), (t) chain-terminating reactions. b Symbol 2 2 t denotes surface-bound species in a simplified notation without specifying surface hydroxyl groups and exchange with arsenite and arsenate and summarizes the different possible types of surface complexes. c Values in parentheses give the order of magnitude of the bimolecular rate coefficients (M-1 s-1) for the corresponding reactions in solution (see references in last column).

FIGURE 6. Rate coefficients for As(III) disappearance from solution in the FH suspension experiments as a function of the H2O2 concentration and comparison with the corresponding rate coefficients for the oxidation of adsorbed As(III) on the FH surface (eq 4a). liter)] of ∼105 L/kg. In agreement with the ATR-FTIR experiments, no oxidation of As(III) to As(V) and thus no further decrease of dissolved As(III) was observed in the presence of FH over the time scale from minutes to hours. Following the addition of 84 µM to 8.4 mM H2O2, however, a further decrease of the dissolved As(III) concentration was observed, which followed almost ideal first-order kinetics {d[As(III)diss]/dt ) -k[As(III)diss], with [As(III)diss] denoting the concentration of dissolved As(III) in the solution and k being the rate coefficient for the disappearance of As(III) from solution}. Oxidation of adsorbed As(III) leads to additional adsorption of As(III) from the solution. In the limit where the adsorption of As(III) is much faster than its oxidation on the surface, the rate of As(III) disappearance from solution would be expected to correspond to the rate of As(III) oxidation on the FH surface. In Figure 6, the rate coefficients for As(III) disappearance from solution are compared with the rate coefficients for the surface As(III) oxidation determined in the ATR-FTIR experiments (eq 3b). At H2O2 concentrations above 180 µM, the rate coefficients determined at 100 mg/L FH are within the 95% confidence limits of the rate coefficients for surface oxidation, whereas at 10 mg/L, the rate coefficients determined in the suspensions are lower. This difference can be explained by the kinetics of As(III) adsorption. At high FH concentrations, the equilibration between dissolved and adsorbed As(III) is fast, and the rate of the disappearance of dissolved As(III) is determined by the rate of oxidation at the

surface [dissolved and adsorbed As(III) are in equilibrium during the oxidation]. At low FH concentration, the adsorption equilibrium is reached more slowly and the disappearance rate for dissolved As(III) is influenced by the rate of As(III) adsorption. The two data points in Figure 6 at 84 µM H2O2 and 100 mg/L FH show a relatively low rate coefficient for As(III) disappearance, and at 8.4 µM H2O2 and 10 or 100 mg/L FH, no As(III) disappearance was observed (data not shown). This is most likely because of an initial rapid loss of a fraction of the H2O2 and appears to occur in systems where the ratio of FH to H2O2 is large, as also reported elsewhere (19). Apart from these deviations, however, the results from the suspension experiments are complementary and in line with the rate coefficients for As(III) oxidation on the FH surface that were determined in situ by ATR-FTIR spectroscopy. Environmental Implications. From the rate coefficients of As(III) oxidation on the FH surface (eq 3a) and As(III) oxidation in solution (eq 5), the ratio R between As(III) oxidation on the surface and in solution between pH ∼6 and ∼10 can be estimated as

log R ) 11.2 - 0.5 log [H2O2] -1.4pH + log [FH] + log Kd (8) The significance of As(III) oxidation on the FH surface increases with decreasing H2O2 concentration ([H2O2] in moles per liter), decreasing pH, increasing FH concentration ([FH] in kilograms per liter), and increasing As(III) adsorption affinity (as expressed by the distribution coefficient Kd in liters per kilogram). In other words, the surface reaction becomes dominant if the reaction in solution is slow and if the fraction of adsorbed As(III) is large. With a rather high pH of 8.5, a H2O2 concentration of 1 mM (found only in engineered systems), a rather low log Kd of 5 [as determined for rather high As(III) concentrations of ∼5 µM], and a very low FH concentration of 1 mg/L (10-6 kg/L), eq 8 yields a ratio R ≈ 1, i.e., As(III) oxidations in solution and on the surface are of similar relevance. However, at neutral to slightly acidic pH, lower H2O2 concentrations, higher Kd, and higher FH (or other amorphous iron hydroxide colloid) concentrations, i.e., under conditions typical for many natural and engineered systems, eq 8 suggests that the surface-catalyzed reaction pathway might be more relevant than the reaction in solution. This applies both to natural systems such as surface waters and to engineered systems such as drinking water treatment. At more acidic pH values, other processes such as catalysis in solution by dissolved Fe(III) might become more important. In environmental systems, the sorption kinetics of As(III), As(V), and other competing anions such as carbonate, silicate, and phosphate will also be important VOL. 37, NO. 5, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

977

factors in the overall oxidation rates and require further attention. The ATR-FTIR technique offers a sensitive tool for studying the surface reactions in these systems in situ.

Acknowledgments Annette Hofmann is gratefully acknowledged for providing synthesized two-line ferrihydrite for this study, Thomas Ruettimann for his help with AFS measurements and suspension experiments, Denis Mavrocordatos for SEM and AFM analysis of iron oxide films, and David Kistler for ICPMS measurements. This study was partly financed by the Alliance for Global Sustainability.

Literature Cited (1) Penrose, W. R. CRC Crit. Rev. Environ. Control 1974, 4, 465482. (2) Martell, A. E.; Smith, R. M.; Motekaitis, R. J. NIST Database 46, Critically Selected Stability Constants of Metal Complexes, Version 6.0; National Institute of Standards and Technology (NIST): Gaithersburg, MD, 2002. (3) Goldberg, S. Soil Sci. Soc. Am. J. 2002, 66, 413-421. (4) Masscheleyn, P. H.; Delaune, R. D.; Patrick, W. H. Environ. Sci. Technol. 1991, 25, 1414-1419. (5) Langner, H. W.; Jackson, C. R.; McDermott, T. R.; Inskeep, W. P. Environ. Sci. Technol. 2001, 35, 3302-3309. (6) Wilkie, J. A.; Hering, J. G. Environ. Sci. Technol. 1998, 32, 657662. (7) Oscarson, D. W.; Huang, P. M.; Defosse, C.; Herbillon, A. Nature 1981, 291, 50-51. (8) Manning, B. A.; Fendorf, S. E.; Bostick, B.; Suarez, D. L. Environ. Sci. Technol. 2002, 36, 976-981. (9) Tournassat, C.; Charlet, L.; Bosbach, D.; Manceau, A. Environ. Sci. Technol. 2002, 36, 493-500. (10) Manning, B. A.; Goldberg, S. Environ. Sci. Technol. 1997, 31, 2005-2011. (11) Lin, Z.; Puls, R. W. Environ. Geol. 2000, 39, 753-759. (12) Sun, X.; Doner, H. E. Soil Sci. 1998, 163, 278-287. (13) De Vitre, R.; Belzile, N.; Tessier, A. Limnol. Oceanogr. 1991, 36, 1480-1485. (14) Manning, B. A.; Fendorf, S. E.; Goldberg, S. Environ. Sci. Technol. 1998, 32, 2383-2388. (15) Pettine, M.; Campanella, L.; Millero, F. J. Geochim. Cosmochim. Acta 1999, 63, 2727-2735. (16) Arienzo, M.; Chiarenzelli, J.; Scrudato, R. Fresenius’ Environ. Bull. 2001, 10, 731-735. (17) Arienzo, M.; Chiarenzelli, J.; Scrudato, R. J. Hazard. Mater. 2001, B87, 187-198. (18) Pettine, M.; Millero, F. J. Mar. Chem. 2000, 70, 223-234. (19) Kwan, W. P.; Voelker, B. M. Environ. Sci. Technol. 2002.

978

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 5, 2003

(20) Lin, S.-S.; Gurol, M. D. Environ. Sci. Technol. 1998, 32, 14171423. (21) Watts, R. J.; Foget, M. K.; Kong, S.-H.; Teel, A. L. J. Hazard. Mater. 1999, B69, 229-243. (22) Watts, R. J.; Udell, M. D.; Kong, S.-H.; Leung, S. W. Environ. Eng. Sci. 1999, 16, 93-103. (23) Hug, S. J.; Sulzberger, B. Langmuir 1994, 10, 3587-3597. (24) Kesselman-Truttmann, J. M.; Hug, S. J. Environ. Sci. Technol. 1999, 33, 3171-3176. (25) McQuillan, A. J. Adv. Mater. 2001, 13, 12-13. (26) Schwertmann, U.; Cornell, R. M. Iron Oxides in the Laboratory; VCH Verlagsgesellschaft: Weinheim, Germany, 1991. (27) Yamamoto, M.; Urata, K.; Murashige, K.; Yamamoto, Y. Spectroc. Acta B: Atom. Spectrosc. 1981, 36, 671-677. (28) Goldberg, S.; Johnston, C. T. J. Colloid Interface Sci. 2001, 234, 204-216. (29) Myneni, S. C. B.; Traina, S. J.; Waychunas, G. A.; Logan, T. J. Geochim. Cosmochim. Acta 1998, 62, 3285-3300. (30) Roddick-Lanzilotta, A. J.; McQuillan, A. J.; Craw, D. Appl. Geochem. 2002, 17, 445-454. (31) Manceau, A. Geochim. Cosmochim. Acta 1995, 59, 3647-3653. (32) Randall, S. R.; Sherman, D. M.; Ragnarsdottir, K. V. Geochim. Cosmochim. Acta 2001, 65, 1015-1023. (33) O’Reilly, S. E.; Strawn, D. G.; Sparks, D. L. Soil Sci. Soc. Am. J. 2001, 65, 67-77. (34) Waychunas, G. A.; Rea, B. A.; Fuller, C. C.; Davis, J. A. Geochim. Cosmochim. Acta 1993, 57, 2251-2269. (35) Su, C.; Suarez, D. L. Clays Clay Miner. 1997, 45, 814-825. (36) Villalobos, M.; Leckie, J. O. J. Colloid Interface Sci. 2001, 235, 15-32. (37) Wijnja, H.; Schulthess, C. P. Soil Sci. Soc. Am. J. 2001, 65, 324330. (38) Cornell, R. M.; Schwertmann, U. The Iron Oxides; VCH Verlagsgesellschaft: Weinheim, Germany, 1996. (39) Hug, S. J.; Canonica, L.; Wegelin, M.; Gechter, D.; von Gunten, U. Environ. Sci. Technol. 2001, 35, 2114-2121. (40) Pignatello, J. J. Environ. Sci. Technol. 1992, 26, 944-951. (41) Rush, J. D.; Bielski, B. H. J. J. Phys. Chem. 1985, 89, 5062-5066. (42) Klaning, U. K.; Bielski, B., H. J.; Sehested, K. Inorg. Chem. 1989, 28, 2717-2724. (43) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J. Phys. Chem. Ref. Data 1988, 17, 513-886. (44) Stuglik, Z.; Zagorski, Z. P. Radiat. Phys. Chem. 1981, 17, 229233. (45) Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L.; Ross, A. B. J. Phys. Chem. Ref. Data 1985, 14, 1041-1100.

Received for review June 3, 2002. Revised manuscript received December 5, 2002. Accepted December 17, 2002. ES025845K