Kinetics and Structural Constraints of Chromate Reduction by Green

May 17, 2003 - ... Sciences, Stanford University, Stanford, California 94305-2115 ...... Clay Conference; Kodama , H., Mermut , A. R., Torrance , J. K...
3 downloads 0 Views 175KB Size
Download PDF
Environ. Sci. Technol. 2003, 37, 2750-2757

Kinetics and Structural Constraints of Chromate Reduction by Green Rusts DEBORAH L. BOND AND SCOTT FENDORF* Department of Geological and Environmental Sciences, Stanford University, Stanford, California 94305-2115

Green rusts, ferrous-ferric iron oxides, occur in many anaerobic soils and sediments and are highly reactive, making them important phases impacting the fate and transport of environmental contaminants. Despite their potential importance in environmental settings, reactions involving green rusts remain rather poorly described. Chromate is a widespread contaminant having deleterious impacts on plant and animal health; its fate may in part be controlled by green rust. Here we examine chromate reduction by a series of green rust phases and resolve the reaction kinetics at pH 7. The overall kinetics of the reactions are well described by the expression d[Cr(VI)]/dt ) -k[Cr(VI)]{GR}, and this model was successfully used to predict rates of reaction at varying chromium concentrations. The rates of reduction are controlled by the concentration of ferrous iron, surface area, and chemical structure of the green rust including layer spacing. On a mass basis, green rust (GR) chloride is the most rapid reductant of Cr(VI) followed by GRCO3 and GRSO4, with pseudo-first-order rate coefficients (kobs) (with respect to Cr(VI) concentration) ranging from 1.22 × 10-3 to 3.7 × 10-2 s-1. Chromium(III)substituted magnetite and lepidocrocite were identified as the major oxidation products. The nature of the oxidation products appears to be independent of the anionic class of green rust, but their respective concentrations display a dependence on the initial GR. The mole fraction of Fe(III) in the Crx,Fe1-x(OH)3‚nH2O reaction product ranged from 17% to 68%, leading to a highly stabilized (low solubility) phase.

Introduction Chromium is a pollutant of concern due to its widespread use in industrial applications such as electroplating, metallurgy, and leather tanning. Thus, it is often introduced into the environment at high levels. In addition, the presence of chromium in ultramafic rocks contributes to its abundance in nature. In the environment, chromiumsa redox active elementsexists as Cr(III) or Cr(VI). The release of Cr(VI) to surface and subsurface waters is of concern as Cr(VI) is highly soluble, highly mobile, and a known toxin, mutagen, teratogen, and carcinogen (1). Thus, processes leading to the reduction of chromate transform the metal to a less mobile and less toxic form, Cr(III). Chromium(VI) species are anionic (HCrO4- and CrO42-) in most neutral to alkaline systems. However, in soils and * Corresponding author phone: (650)723-5238; fax: (650)725-2199; e-mail: [email protected]. 2750

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 12, 2003

groundwaters with acidic pH values, these oxyanions have the potential to sorb on iron oxides and (oxy)hydroxides decreasing their aqueous concentration (2-4). Under reducing conditions, Cr(VI) may be converted to Cr(III) which strongly adsorbs to mineral surfaces and forms hydroxide solids of limited solubility, thus minimizing its transport (1, 5). Reduction of Cr(VI) species readily occurs by reaction with aqueous ferrous iron (e.g., refs 6-8), H2S (9); FeS (10), and organic compounds (11-13). In addition, Fe(II) solids contribute to the formation and precipitation of Cr(III) (2, 14). Reduction of Cr(VI) by Fe(II) has been observed in soils (15), aquifers (16), and in permeable reactive barriers (17, 18). In anaerobic systems, abiotic pathways control the reduction of chromate, with Fe(II) exerting the dominant influence in all but acidic environments (19). The rate of Cr(VI) reduction by aqueous ferrous iron is established and proceeds rapidly (6-8). Even under acidic conditions (pH 3), Fe(II) appears to be competitive with organic matter as a reductant of chromate (3). In addition to solids with structural Fe(II) (addressed subsequently), chromate reduction is also possible when aqueous Cr(VI) sorbs on Fe(III) solids such as maghemite (γ-Fe(III)2O3), hematite (R-Fe(III)2O3), goethite (R-Fe(III)OOH), or hydrous ferrous oxides (Fe(OH)3‚nH2O) provided surface-bound Fe(II) is present (20-22). As ferrous-iron rich solutions and solids are effective reductants of chromate in natural systems, green rust minerals in reduced soils (23, 24) may have a pronounced influence on Cr(VI) reduction. Green rusts, Fe(II)-Fe(III) oxyhydroxides within the sjo¨grenite-pyroaurite structural group consisting of brucite-like layers, were initially characterized as corrosion products in steel and iron water distribution pipes (25). Trioctahedral layers of ferrous hydroxide substituted with Fe(III) produce an overall net positive charge that is neutralized by the association of anions and water in the interlayer. The nature of the interlayer anion directly affects the structure and classification of the green rust. The classification of green rust 1 (GR1) refers to the chloride and carbonate species, representing planar anions within the interlayer. In contrast, green rust sulfate falls under the designation GR2 as it is a tetrahedrally coordinated anion occupying the interlayer. Green rusts exist as intermediates in the formation and reduction of ferric iron oxides such as goethite, lepidocrocite, and magnetite (26). In addition, green rusts are a product of microbial degradation by dissimilatory iron reducing bacteria (27, 28) and have been identified in field-scale permeable reactive barriers containing Fe(0) (29, 30). Due to their highly reactive surfaces and reduction potentials, green rusts may have strong influences on the fate of environmental toxins such as the chromate ion. The ability of green rusts to reduce environmentally relevant toxins has been examined in the context of some common redox active elements. The sulfate form of green rust is the most commonly studied, more than likely due to its ease of synthesis. The contaminants evaluated thus far for GRSO4 are nitrate (31, 32), nitrite (33), selenate (34), arsenate (35), halogenated methanes (36), and chromate (37, 38). Chloride green rust has been studied with respect to nitrate (39), nitrite (39), and chromate (38). The least stable form of green rust, GR1(CO3), has minimal representation in the literature, with only reactions involving nitrate (40) and chromate (41). At the present time, there is a paucity of data on the kinetics of green rust redox reactions, yet such parameters are key 10.1021/es026341p CCC: $25.00

 2003 American Chemical Society Published on Web 05/17/2003

to defining the role of these solids on the fate of toxins. The few reaction rates that have been reported reveal pseudofirst-order kinetics with respect to contaminant concentration (31, 32, 34, 36, 41, 42). In addition, GRCl has been noted as a more facile reductant than GRSO4 (42). Finally, reduction rates for specific compounds appear more rapid when reactions occur prior to green rust precipitation rather than after its formation (33, 34). While the environmental relevance of chromate reduction by green rust is evident, as demonstrated by the thermodynamic favorability of such reactions (19), their interactions have not been resolved. In particular, it is presently unknown how variations in green rust structure will impact chromate reduction. Although initial reports provide an overview of reduction mechanisms by one member of the green rust family (31, 33, 34, 37, 38, 41, 42), a comprehensive examination of structural class has not been conducted. Such information is critical for evaluating the potential deviation in green rust reactivity and specifically on chromate reduction. Accordingly, the objective of this study was to evaluate the reduction of Cr(VI) by three main classes of green rust: GR sulfate, chloride, and carbonate. Rates of reduction were measured in order to assess the significance of the chemical and structural differences within green rusts on chromate reduction. Additionally, we examine variation in the reaction products of chromate reduction by the different forms of green rust.

Materials and Methods Chemical Synthesis Procedures. All chemicals were ACS reagent grade and used without further purification. Green rusts (GR) chloride and sulfate were synthesized per the method of Schwertmann and Fechter (26). Green rust carbonate was synthesized using the method of Benali et al. (43) that accounts for the low solubility of ferrous carbonate (FeCO3). All solutions were made anoxic by boiling and cooling distilled, deionized water under O2-free, N2 gas. All glassware was equilibrated in the N2 glovebox for at least 24 h prior to use. Reactions were performed in a specially designed chamber that was constructed with inlets for nitrogen, air, acid, base, Eh and pH electrodes, and a sampling port. The samples were vacuum filtered through a 0.22-µm Millipore filter under N2, washed with anoxic water, and dried. The solids were then sealed and stored in the glovebox in N2 equilibrated glass serum vials. Batch Reduction. The dried GR samples were sieved through an ASTM standard no. 60 mesh sieve in order to exclude size fractions greater than 250 µm. Each reaction was carried out in a 250 mL glass beaker in an N2-purged glovebox and all were magnetically stirred. MOPS buffer was added in order to maintain pH and minimize adverse interactions upon analysis. Reactions were performed in duplicate with initial and final pH values of 7.0 ( 0.1. Keeping initial Cr(VI) concentrations constant and varying the green rust concentration, the solid-phase green rust was suspended in 100 mL of 5 mM MOPS. Chromate (as K2CrO4) was then introduced to the suspension. Samples were taken at various time points during the reaction sequence. Solid and Solution Phase Analyses. Both filtered and unfiltered samples were analyzed colorimetrically by the 1,5diphenyl carbazide method (44) with a 5 mM MOPS matrix using a Shimadzu 1601 UV-visible spectrophotometer at 540 nm. The two sampling methods were employed in order to ascertain whether chromate was removed from solution via reduction or sorption to the iron oxide surface in its oxidized form. Total solution chromium in filtered samples was measured using a Thermo Jarrell Ash IRIS inductively coupled plasma (ICP) spectrometer. X-ray diffraction patterns were recorded from 5 to 65° 2θ using Cu KR1 radiation (λ ) 1.54056 nm). A Rigaku Geiger-

flex Powder Defractrometer with generator energy of 35 keV and a power of 15 mA was employed. Samples were prepared in the glovebox by admixing with glycerol to form a paste which was smeared onto the sample slide, similar to the approach described by ref 45. The glycerol paste was used to limit oxidation during XRD analysis but produced a broad peak in the background between 15 and 25° 2θ. Surface area was determined using a 5-point BET analysis on a Coulter SA 3100. Samples were dried in the glovebox by vacuum desiccation for 1 week prior to analysis. Exposure to air during transfer to the instrument was minimal as evidenced by the lack of color change even in the most sensitive green rust sample, GR-CO3. X-ray absorption spectra were collected on beamlines 4-1 and 4-3 at the Stanford Synchrotron Radiation Laboratory. The storage ring operates at an energy of 3.0 GeV with currents between 100 and 50 mA. A Si (220) monochromator was used to scan the incident X-ray beam through the K-edge of chromium and iron (the edge for metallic chromium was set to 5989 eV and to 7111 eV for metallic iron). The beam was detuned approximately 70% for chromium and 50% for iron to prevent interferences from higher-order harmonics. Internal energy calibration for each sample was accomplished using an iron foil or Cr(OH)3‚nH2O positioned between the second and third in-line ionization chambers. Incident and transmitted intensities were measured with N2 purged ionization chambers (I0 and I1); fluorescence intensity was measured using a Lytle detector (46) 45° off the sample and orthogonal to the incident beam. Each sample was run for 2-3 scans, and these repetitions were averaged together for data analysis. XAS data analysis was performed using WinXAS 2.0. Spectra were calibrated, and the background was subtracted using a two-polynomial fit. The spectra were then normalized to a unit jump height (atomic cross-section) to allow comparison of spectral features on a per-atom basis (47). Pre-edge features within the Cr spectra were used to deduced, and quantify, the contributions of Cr(VI) and Cr(III). In the case of iron, data analysis for the XANES region was carried using Peak-Fit 4.0 (Jandel Scientific) to determine Fe(II):Fe(III) ratios. The first derivative of the XANES spectrum was truncated to contain only the necessary region around the inflection point. The resulting spectrum was then fit using Gaussian amplitude functions (positive and negative). Using a first derivative Savitzky-Golay method, the spectrum was smoothed (1%) with a Linear D2 baseline of 3%. Peak fitting was based on the nonlinear Marquardt-Levenberg method, and the AutoFit Peaks I Residuals was used to optimize the fits. The higher oxidation state, Fe(III), has a larger effective nuclear charge and thus has a slightly larger binding energy. Raman spectra were obtained using a Kaiser HoloLab Series 5000 Raman microscope with a charge coupled device (CCD) detector and motorized X-Y stage. The sample excitation was carried out with an Invictus diode laser (785 nm) with a low light power (less than 0.7 mW average power) to prevent sample degradation caused by heating and photooxidation. Raman spectra were collected through the 100X objective of an optical microscope in order to enhance the signal-to-background ratio. A minimum collection time for all samples was 15 s, and 15 spectra were averaged over a Raman shift of 100-3500 cm-1 with no visual indication of sample oxidation. Samples were mounted in powder form on a platinum coated mirror, and the sample stage was continuously purged with N2 gas to prevent oxidation.

Results and Discussion Reaction Kinetics. The reduction of Cr(VI) by green rusts SO42-, Cl-, and CO32-, conducted at three different green rust concentrations (0.125, 0.25, and 0.5 g/L), was carried out in order to ascertain the rate expression describing each VOL. 37, NO. 12, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

2751

FIGURE 1. Chromate reduction rates upon reaction with (a) GRCl, (b) GRCO3, and (c) GRSO4 at various suspension densities. Batch reactors were maintained at pH 7 with 5 mM MOPS buffer. Initial Cr(VI) concentration was 192 µM. Solid lines and symbols represent first-order kinetics models and actual data, respectively. Error bars are based on duplicate reactors

FIGURE 2. Effect of (a) green rust concentration, (b) ferrous iron concentration, and (c) surface area on first-order rate coefficients of Cr(VI) reduction by GRCl ([), GRCO3 (9)), and GRSO4 (2). Reactors contained 192 µM Cr(VI) at pH 7. reaction. Each data set was modeled using a pseudo-firstorder rate expression, d[Cr(VI)]/dt ) -kobs[Cr(VI)], and conformance to the expression is noted (Figure 1) suggesting that indeed the reaction is first-order with respect to chromate concentration. The reaction can also be described as having a first-order dependence on green rust concentration on the basis of the linearity observed in the pseudo-firstorder rate coefficients with respect to suspension density 2752

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 12, 2003

(Figure 2a). The lack of a zero y-intercept, however, indicates that the reaction is governed by more complicated mechanisms. Nevertheless, the reduction of Cr(VI) by green rust can be well approximated by an overall second-order reaction, dependent on both the concentration of chromate and green rust

d[Cr(VI)]/dt ) -k[Cr(VI)]{GR}

(1.1)

FIGURE 3. Chromate reduction upon reaction with (a) GRCl, (b) GRCO3, and (c) GRSO4. Batch reactors were maintained at pH 7 with initial GR suspension density ) 0.25 g/L. Solid lines represent overall second-order kinetics model derived from 192 µM reactor applied to varying initial chromate concentrations of 25 µM, 50 µM, 75 µM (for the the case of GRCl), and 125 µM. where {GR} is the solution density of green rust. On the basis of mass concentrations, green rust chloride is the most facile of the three species in chromate reduction; pseudo-first-order rate coefficients (kobs), dependent on chromate concentration, ranged from 1.26 × 10-3 to 3.66 × 10-2 s-1 (for GR suspension densities ranging from 0.125 to 0.5 g/L). The carbonate system followed with a range of kobs from 1.92 × 10-3 to 2.39 × 10-2 s-1. Finally, green rust sulfate showed the slowest rate of reduction with kobs of 6.10 × 10-41.81 × 10-2 s-1. The values obtained for the carbonate system are similar to those obtained in a separate study (41). Using these observed rate coefficients, complete rate expressions for each green rust phase are

GRCl: d[Cr(VI)]/dt ) -5.28 × 10-2 Lg-1 s-1 [Cr(VI)]{GR} (1.2a) GRCO3: d[Cr(VI)]/dt )

-3.34 × 10-2 Lg-1 s-1 [Cr(VI)]{GR} (1.2b)

GRSO4: d[Cr(VI)]/dt )

-1.95 × 10-2 Lg-1 s-1 [Cr(VI)]{GR} (1.2c)

To test the validity of the expressions, systems in which the initial chromate concentration was varied between 25 µM and 125 µM were modeled. The model approximated the data within the analytical error of measurement for the entire concentration range (Figure 3), further supporting the applicability of a second-order expression in describing the rate of chromate reduction by green rust. The reactivity of green rust is influenced by a series of factors that are manifested in the mass normalized rate. Given the dependence of chromate reduction on suspension density, the reactive surface area should have an important influence on the reaction rate. The relative concentration of ferrous iron present in the solid along with the interlayer

structure, as dictated by the anion component, must also be considered. Finally, the relative bond strengths of the solids, as represented by the free energy of formation, may also impact the observed reaction rates. On a surface area basis, GRCO3 with the largest surface area (30.1 m2/g) should react the most rapidly followed by GRCl (S.A. 19.0 m2/g) and GRSO4 (S.A. 3.6 m2/g) (Figure 2c). However, the mass normalized data are inconsistent with this expectation, chloride having the greatest reaction rate followed by carbonate and sulfate. We must therefore explore other contributing factors to the reactivity of green rust to explain the inversion of chloride and carbonate. The proposed chemical formulas for each of the three green rust solids investigated allow the ferrous iron proportion to be evaluated and its impact on reduction rates explored. The ferrous to ferric iron ratios for GRCl, GRCO3, and GRSO4 are 3.1:1, 1.9:1, and 2.0:1, respectively, as determined by fitting the first derivative of the Fe-XANES spectra (Figure 4). GR1-Cl has a formula of [FeII3FeIII(OH)8]+‚[Cl‚nH2O]-, while GR1-CO3 and GR2-SO4 are approximately [FeII4Fe2III(OH)12]2+‚[CO3‚nH2O]2- and [FeII4Fe2III(OH)12]2+‚[SO4‚nH2O]2-, respectively (48), in agreement with the Fe(II)/Fe(III) ratios measured here. Since the ferrous to ferric iron ratio is larger for Cl- than for CO32- or SO42- (3:1 vs 2:1), a greater proportion of reduced iron is available. And although GRCO3 and SO4 have the same Fe(II)/Fe(III) ratio, GRCO3 has more Fe(II) on a per mol basis. Consistent with the fraction of Fe(II)/Fe(III), green rust chloride is the most facile reductant of chromate (Figure 2a). In fact, when the data are normalized to ferrous iron concentration, a sequence consistent with mass normalized data is observed Cl > CO3 > SO4 (Figure 2b). The free energies of formation, a proxy for the lability of the surface groups, for the various solids also support the observed trend in reaction rates on a mass basis. GRCl, having a lower ∆G°f (48), has a greater tendency toward reaction and electron transfer than GRCO3, allowing for a more rapid VOL. 37, NO. 12, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

2753

FIGURE 4. Fe XANES analysis displaying fits of ferrous (dotted line) and ferric (dashed line) contributions to the first derivative spectra (solid line) of (a) GRCl, (b) GRCO3, and (c) GRSO4. Also shown are the linear baseline backgrounds applied to each GR. interaction with chromate for the former. The thermodynamic constants derived for green rusts correspond to the trends noted in reduction rates for mass- and ferrous iron concentration-based analyses. The calculated Gibbs free energy (∆G°f) of anhydrous GRCl, CO3, and SO4 are -2146 ( 5, -3590 ( 10, and -3795 ( 15 kJ/mol, respectively (48). Thus, on a per mass basis, the combined contribution of ferrous iron concentration (relative) and free energy of formation are consistent with the observed reaction rates. Surface area, while not consistent with the observed reaction rates, will nevertheless have a large impact on green rust reactivity. To eliminate this contribution and more fully evaluate other contributing factors to green rust reactivity, surface area normalized data need to be evaluated. We should first, however, consider the heterogeneity in green rust surfaces. Similar to phyllosilicate clay minerals, the reactive surfaces of green rust can be classified as external sites composed of (100) and (010) faces and internal (or interlayer) sites along the (001) plane. Access to the internal sites relies upon exchange of interlayer anions. If one considers only the external sites as reactive, the surface area normalized reduction rates for chromate should all be equal among the green rust phases. Clearly this is not the case. When considering reduction rates normalized to surface area, green rust sulfate has the greatest reaction rate followed by chloride and then carbonate (Figure 2c). Neither ferrous iron concentration nor free energies of formation are able to account for these observations. We, therefore, need to address interlayer reactivity, which will largely be controlled by anion access or exchange. Green rust-SO4, being a class-2 green rust, hosts a tetrahedrally coordinated interlayer anion which results in a larger d-spacing within the compound than that of the planar anions within GRCl and GRCO3. Chromate, with its tetrahedral structure, should therefore exchange more readily with SO42- than Cl- or CO32-, allowing for greater access to the structural Fe(II) (Figure 5). Between GRCl and CO3, the d-spacings are consistent and will not suffice to explain the differences in surface area normalized reaction rate. However, continuation of the above argument would imply less favorable exchange of chromate with the interlayer anions simply by the nature of its coordination, size, and steric hindrances. The available surface sites on GRCl are enriched 2754

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 12, 2003

FIGURE 5. Structural characteristics of green rust (a) SO4, (b) CO3, and (c) Cl. GR2-SO4 with its tetrahedrally coordinated interlayer anion and larger d-spacing is better suited to exchange with chromate in the aqueous phase. Planar chloride and carbonate create smaller d-spacings, limiting exchange with tetrahedral chromate thereby decreasing reation rates. in Fe(II) compared to Fe(III), and this alone is able to reconcile the more rapid rate of reduction observed for GRCl than for GRCO3. Comparison of rates between green rust phases has also been made for nitrate, considering GRSO4 and GRCl (42). Nitrate exhibited a decreased reduction rate for GRSO4. Consistent with structural compatibilities, GR-Cl would be more likely to favor exchange of nitrate for chloride since both are planar anions and lead to GR-1 type structures. The charge of the exchanger and its affinity for interlayer sites should also be considered with data normalized to external surface area. Divalent anions will have a greater affinity than monovalent ions for interlayer sites, consistent with the reactivity of nitrate with GRSO4 as compared to GRCl (42). However, the surface area normalized data indicate a more rapid rate for the divalent occupied interlayer (GRSO4) than for the monovalent interlayer (GRCl), thus suggesting that the interlayer structure (spacing) may have an overriding effect on exchange. In general, minerals with layered structures, such as those in the pyroaurite class, exhibit terminal bonds due to

FIGURE 6. (a) Chromium XANES spectra of solid-phase products showing the minimal contribution from Cr(VI) pre-edge peak (dashed line). (b) Chromium EXAFS spectra (k3 weighted) and (c) corresponding Fourier transforms for solid-phase end products of chromate reduction by green rusts. All EXAFS data were fit over k ) 3 to 11 Å-1. Edge- (E) and corner- (C) sharing peaks used to determine mole fractions of Cr(III) and Fe(III) are indicated. uncoordinated oxygen or hydroxyl groups on the (100) and (010) faces creating reactive surfaces. Charge satisfied hydroxyl groups on the (001) plane lead to relatively chemically inert faces. In the case of green rusts, and other layered clay minerals, this should limit the chemical reactivity of interlayer planes. However, isomorphic substitution of Fe(III) for Fe(II) in the green rust octahedral layers disrupts the charge balance of the (001) surface groups allowing for potentially reactive surfaces within the mineral. Furthermore, electron transfer may occur within the interlayer region through bridging ligands. Thus, in addition to reactions at external sites, anion exchange within the green rust interlayers leads to a second mode of reaction, in support of earlier hypothesis for green rust reactivity (37). Solid-Phase Characteristics and Chromium Sequestration. Upon complete reduction of Cr(VI), the same products, magnetite and lepidocrocite, developed for each GR phase as noted within XRD patterns. Magnetite concentrations increase in the order GRCl < GRSO4 < GRCO3, while lepidocrocite concentrations were greatest for GRSO4 and least for GRCl. Raman spectra collected on the solid-phase product of each reaction product confirmed the presence of magnetite and lepidocrocite, with characteristic bands at 338 and 667 cm-1, respectively (49).

The formation of ferric hydroxides, such as lepidocrocite, with coincident production of Cr(III), should lead to a strong partitioning of chromium into the solid phase. Indeed, a large fraction of chromium is removed from solution and even upon acidification (pH ∼ 2) of samples chromium remains sequestered within the solids. Normalized Cr-XANES spectra of the solid products reveal no appreciable contribution from Cr(VI), as noted by a lack of a pre-edge feature characteristic of chromate (Figure 6a). A control was run with no green rust solid phase present, and Cr(VI) was conserved in the system, indicating that green rust is the only contributor to the reduction of chromate and retention of chromium (Figure 2). Chromium reduction by Fe(II) bearing compounds results in mixed Cr/Fe hydroxides that are sparingly soluble (see, for example, thermodynamic consideration by Sass and Rai (50) and structural aspects reported by Hansel et al. (51) and Loyaux-Lawniczak et al. (38)). Furthermore, Cr(III) forms strong surface complexes on iron oxides and oxyhydroxides (2, 52-54). Given the rapid oxidation of ferrous iron and concurrent reduction of chromate, the combined effects of coprecipitation with, and adsorption on, ferric hydroxides limits dissolved chromium levels. As shown previously (51), the contribution of Cr(III) in mixed chromium-iron oxides VOL. 37, NO. 12, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

2755

TABLE 1. Mole Fraction of Cr(III) and Apparent Solubilities of Mixed Cr(III)/Fe(III) Hydroxide End Products in the Reduction of Chromate by Green Rust GR

I3 /I2 a

Cr(III) mole fractionb

solubility (Log Ks)c

Cl CO3 SO4

0.552 0.359 0.272

0.829 0.481 0.324

4.14 3.67 3.23

a Intensity of Cr corner-shared shell (I ), normalized to the Cr-Cr(Fe) 3 edge-sharing shell (I2), within the Fourier transformed EXAFS spectra. b Values obtained from the linear relationship (y ) 0.5531x + 0.0931, R2 )0.95) relating the I3/I2 intensity within the FT to the mole fraction of Cr(III) within the Cr/Fe hydroxide as described by Hansel et al. (51). c Solubility is a function of the Cr(III) mole fraction as described by log Ks ) 4.23 - 0.172(1-x)2 - 1.392(1-x)3 + log x (55).

can be extrapolated from the Fourier transform of k3 weighted chromium XAS data (Figure 6b,c). The intensity of the Cr corner sharing shell (3.48-3.56 Å), normalized to the edgesharing Cr-Cr(Fe) shell (3.04-3.30 Å), is proportional to the mole fraction of Cr within the solid (51). Such an analysis demonstrates an increase in the mole fraction of chromium from 32% in products derived from GRSO4 to 83% in those resulting from GRCl (Table 1). This order corresponds to the increasing amount of Fe(II) on a per mole basis in the green rusts. Moreover, the solubility of Cr/Fe-hydroxides varies depending on the mole fraction of Cr relative to Fe within the solid, where the solubility product (Ks) can be estimated from the relation

log Ks ) 4.23 - 0.172(1-x)2 - 1.392(1-x)3 + log x where x represents the mole fraction of Cr(III) (55). The solubilities (log Ks) of the end products thus range from 3.23 to 4.14. The inclusion of even 17% Fe(III), the least enriched of the products, decreases the solubility of the phase nearly 2 orders of magnitude below that even for the sparingly soluble Cr(OH)3‚nH2O phase. Thus, the solid products developed from the reduction of chromate by green rust should be expected to greatly limit dissolved concentrations of chromium. Environmental Implications. Chemical reduction rates of chromate by various reductants, particularly ferrous iron bearing phases, were compared using initial concentrations of [Cr(VI)] ) 100 µM and [Fe(II)](aq) ) 30 µM (Figure 7). Using estimates of 5% green rust within a soil or sediment having 5% total iron (0.27% green rust within the solid phase), the rate of chromate reduction by green rust far exceeds other reductants, even aqueous Fe(II). It is unlikely that the formation of magnetite, another mixed valence iron oxide, contributes appreciably to the reduction of chromate when green rust phases occur. The rate of reduction of Cr(VI) by magnetite at pH 7 is approximately 3.3 × 10-8 M/s (unpublished data). This is nearly 103 times slower than the reaction rates observed by the various forms of green rust. In addition, Cr(III) sorption at the reaction conditions, pH 7, is quite strong, and previous studies have shown a passivating layer on magnetite (52). A more detailed rate expression is available for aqueous ferrous iron reduction of chromate (8), and at pH 7 green rust reactions proceed at rates approximately an order of magnitude faster than the trend for Fe2+(aq) reduction of Cr(VI). Green rusts are one of the most rapid reductants of chromate, abiotic or biotic, when compared to rates compiled for other common reduction pathways (19, 56). Using one of the fastest rates reported for bacterial reduction of chromate, considering Desulfovibrio vulgaris (3 × 10-4 M/h) (56), green rusts react at rates in excess of 50-times those of biological induced reductions. In addition, of the previous 2756

9

ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 37, NO. 12, 2003

FIGURE 7. Rate comparison of contaminant reduction by ferrous iron phases. Initial concentrations of [Fe2+(aq)] ) 30 µM and [X] ) 100 µM where X ) Cr(VI), NH4+, Se(VI), and CCl4. {GR} ) 3.98 g/L based on assumption of 0.27% green rust (see text) in the solid phase and a bulk density of 1.5 g/cm3. Rate expression for Fe2+(aq) is from ref 8; for magnetite from ref 57; for ilmenite from ref 22; for Desulfovibrio vulgaris from ref 56; for Cr(VI) by GRCO3* from ref 41; for Se(VI) from ref 34; for NH4+ by GRCl and GRSO4 from ref 42; and for CCl4 from ref 36. Dotted lines indicate the minimal effect of pH on reaction rate in GR systems (41). kinetic studies involving reduction of NH4+ (40), CCl4 (36), and Se(VI) (34) by GR, chromium reacts the most rapidly. Chromate is, in fact, reduced to Cr(III) by the various green rusts in a process which can be described by a secondorder kinetic expression dependent on chromium and green rust concentration. Chromate reduction rates by GRCl are ∼two times faster than those of GRCO3 and GRSO4 on a per mass basis. Rates of reduction are dependent on the external and internal (interlayer) surfaces. When normalized to external surface area, the structural class (GR1 or GR2) appears to have a controlling factor on the reaction ratesthe Fe(II):Fe(III) ratio has a secondary impact within a structural class. Nevertheless, the reactivity of the external surface has an overriding impact on reduction rate leading to both GRCl and GRCO3 having greater reduction rates (per mass) than GRSO4. The combined influence of external surface area and Fe(II):Fe(III) ratio results in GRCl having the greatest reduction rate. On a practical level, a mass-based analysis is of greater significance in natural settings as systems equilibrate to mass, not surface area. Therefore, field analyses will, more than likely, produce results similar to the per mass reduction rates. Nevertheless, all three anionic classes of green rust appear to be facile reductants of chromate that may impart an important control on the hazard of this ubiquitous toxin.

Acknowledgments This research was funded by the Natural and Accelerated Bioremediation Research (NABIR) program, Biological and Environmental Research (BER), U.S. Department of Energy (Grant DE-FG03-00ER63029).

Literature Cited (1) Weckhuysen, B. M.; Wachs, I. E.; Schoonheydt, R. A. Chem. Rev. 1996, 96, 3327-3349. (2) Dzombak, D. A.; Morel, F. M. M. Surface Complexation Modeling: Hydrous Ferric Oxide; Wiley-Interscience: New York, 1990. (3) Eary, L. E.; Rai, D. Soil Sci. Soc. Am. J. 1991, 55, 676-683. (4) Bartlett, T. J.; James, B. R. In Chromium in the Natural and Human Environments; Nriagu, J. O., Niebor, E., Eds.; John Wiley and Sons: New York, 1988; pp 276-304. (5) Eary, L. E.; Rai, D. Environ. Sci. Technol. 1988, 22, 2, 972-977. (6) Fendorf, S. E.; Li, G. Environ. Sci. Technol. 1996, 30, 0, 16141617. (7) Sedlak, D. L.; Chan, P. G. Geochim. Cosmochim. Acta 1997, 61, 2185-2192. (8) Pettine, M.; D’Ottone, L.; Campanella, L.; Millero, F. J.; Passine, R. Geochim. Cosmochim. Acta 1998, 62, 1509-1519. (9) Pettine, M.; Millero, F. J.; Passino, R. Marine Chem. 1994, 46, 335-344. (10) Patterson, R. R.; Fendorf, S. E.; Fendorf, M. Environ. Sci. Technol. 1997, 31, 2039-2044. (11) Ishibashi, Y.; Cerantes, C.; Silver, S. Appl. Environ. Microbiol. 1990, 56, 2268-2270. (12) Llovera, S.; Bonet, R.; Simon-Pujol, M. D.; Congragado, F. Appl. Environ. Microbiol. 1993, 59, 3516-3518. (13) James, B. R.; Bartlett, R. G. J. Environ. Qual. 1983, 12, 169-172. (14) Charlet, L.; Manceau, A. A. J. Colloid. Interface Sci. 1992, 148, 443-458. (15) Masscheleyn, P. H.; Pardue, J. H.; DeLauen, R. D.; Patrick, W. H., Jr. Environ. Sci. Technol. 1992, 26, 6, 1217-1226. (16) Anderson, L. D.; Kent, D. B.; Davis, J. A. Environ. Sci. Technol. 1994, 28, 178-185. (17) Pratt, A. R.; Blowes, D. W.; Ptacek, C. J. Environ. Sci. Technol. 1997, 31, 2492-2498. (18) Powell, R. M.; Puls, R. W.; Hightower, S. K.; Sabatini, D. A. Environ. Sci. Technol. 1995, 29, 1913-1922. (19) Fendorf, S.; Wielinga, B.; Hansel, C. M. Intl. Geol. Rev. 2000, 42, 691-701. (20) Bidoglio, G.; Gibson, P. N.; O’Gorman, M.; Roberts, K. J. Geochim. Cosmochim. Acta 1993, 57, 2389-2394. (21) Eary, L. E.; Rai, D. Am. J. Sci. 1989, 289, 180-213. (22) Peterson, M. L.; Brown, G. E., Jr.; Parks, G. A. Colloids Surf. A - Physiochem. Eng. Aspects 1996, 107, 77-88. (23) Trolard, F.; Ge´nin, J. R.; Abdelmoula, M.; Bourrie, G.; Humbert, B.; Herbillon, A. Geochim. Cosmochim. Acta 1997, 61, 11071111. (24) Ge´nin, J. R.; Refait, P.; Bourrie, G.; Abdelmoula, M.; Trolard, F. Appl. Geochem. 2001, 16, 559-570. (25) Stampfl, P. P. Corros. Sci. 1969, 9, 185-187. (26) Schwertmann, U.; Fechter, H. Clay Miner. 1994, 29, 87-92. (27) Fredrickson, J. K.; Zachara, J. M.; Kennedy, D. W.; Dong, H.; Onstott, T. C.; Hinman, N. W.; Li, S. Geochim. Cosmochim. Acta 1998, 62, 3239-3257. (28) Parmer, N.; Gorby, Y. A.; Beveridge, T. J.; Ferris, F. G. Geomicrobiol. J. 2001, 18, 375-385. (29) Ritter, K.; Odziemkowski, M. S.; Gillham, R. W. J. Contam. Hydrol. 2002, 55, 87-111. (30) Roh, Y.; Lee, S. Y.; Elless, M. P. Environ. Geol. 2000, 40, 184-194. (31) Hansen, H. C. B.; Koch, C. B.; Nancke-Krogh, H.; Borggaard, O. K.; Sorensen, J. Environ. Sci. Technol. 1996, 30, 0, 2053-2056. (32) Hansen, H. C. B.; Koch, C. B. Clay Miner. 1998, 33, 87-101.

(33) Hansen, H. C. B.; Hansen, B.; Borggaard, O. K.; Sorensen, J. Geochim. Cosmochim. Acta 1994, 58, 2599-2608. (34) Myneni, S. C. B.; Tokunaga, T. K.; Brown, G. E., Jr. Science 1997, 278, 1106-1109. (35) Randall, S. R.; Sherman, D. M.; Ragnarsdottir, K. V. Geochim. Cosmochim. Acta 1999, 65, 1015-1023. (36) Erbs, M.; Christian, H.; Hansen, B.; Olsen, C. E. Environ. Sci. Technol. 1999, 33, 307-311. (37) Loyaux-Lawniczak, S.; Refait, P.; Lecomte, P.; Ehrhardt, J.-J.; Ge´nin, J. M. R. Hydrol. Earth Syst. Sci. 1999, 3, 593-599. (38) Loyaux-Lawniczak, S.; Refait, P.; Ehrhardt, J. J.; Lecomte, P.; Ge´nin, J. R. Environ. Sci. Technol. 2000, 34, 438-443. (39) Hansen, H. C. B.; Koch, C. B.; Erbs, M.; Dickow, J.; Guldberg, S. Abstr. Papers Am. Chem. Soc. 2000, 220, U327-U327. (40) Hansen, H. C. B.; Koch, C. B. Clays for Our Future, Proceedings of the 11th International Clay Conference. In Clays for Our Future, Proceedings of the 11th International Clay Conference; Kodama, H., Mermut, A. R., Torrance, J. K., Eds.; pp 295-302. (41) Williams, A. G. B.; Scherer, M. M. Environ. Sci. Technol. 2001, 35, 3488-3494. (42) Hansen, H. C. B.; Guldberg, S.; Erbs, M.; Koch, C. B. Appl. Clay Sci. 2001, 18, 81-91. (43) Benali, O.; Abdelmoula, M.; Refait, P.; Ge´nin, J. R. Geochim. Cosmochim. Acta 2001, 65, 1715-1726. (44) Bartlett, R.; James, B. J. Environ. Qual. 1979, 8, 31-35. (45) Hansen, H. C. B.; Poulsen, I. F. Clays Clay Miner. 1999, 47, 312318. (46) Lytle, F. W.; Greegor, R. B.; Sandstone, D. R.; Marques, E. C.; Wong, J.; Spiro, C. L.; Huffman, G. P.; Huggines, F. E. Nucl. Instrum. Methods Phys. Res. Sect. A 1984, 226, 542-548. (47) Sto¨hr, J. In Emission and Scattering Techniques: Studies of Inorganic Molecules, Solids, and Surfaces; Day, P., Ed.; Reidel: Boston, 1981; pp 213-250. (48) Ge´nin, J. R.; Bourrie, G.; Trolard, F.; Abdelmoula, M.; Jaffrezic, A.; Refait, P.; Maitre, V.; Humbert, B.; Herbillon, A. Environ. Sci. Technol. 1998, 32, 1058-1068. (49) Cornell, R. M.; Schwertmann, U. The Iron Oxides: Structure, Properties, Reactions, Occurrence and Uses; VCH Publishers: Weinheim, 1996. (50) Sass, B. M.; Rai, D. Inorg. Chem. 1987, 26, 6, 2228-2232. (51) Hansel, C. M.; Wielinga, B. W.; Fendorf, S. Geochim. Cosmochim. Acta 2003, 67, 401-412. (52) Kendelewicz, T.; Liu, P.; Doyle, C. S.; Brown, G. E. J.; Nelson, E. J.; Chambers, S. A. Surf. Sci. 1999, 424, 219-231. (53) Charlet, L.; Manceau, A. J. Colloid. Interface Sci. 1992, 148, 443458. (54) Leckie, J. O.; Appleton, A. R.; Ball, N. B.; Hayes, K. F.; Honeyman, B. D. Adsorptive removal of trace elements from fly ash pond effluents onto iron oxyhydroxide; Electric Power Research Institute: Palo Alto, CA, 1984. (55) Rai, D.; Zachara, J.; Eary, L.; Girvin, D. C.; Moore, C. A.; Resch, C. T.; Sass, B. M.; Schmidt, R. L. (56) Lovley, D. R.; Phillips, E. J. P. Appl. Environ. Microbiol. 1994, 60, 726-728. (57) Peterson, M. L.; White, A. F.; Brown, G. E., Jr.; Parks, G. A. Environ. Sci. Technol. 1997, 31, 1573-1576.

Received for review November 18, 2002. Revised manuscript received March 1, 2003. Accepted April 7, 2003. ES026341P

VOL. 37, NO. 12, 2003 / ENVIRONMENTAL SCIENCE & TECHNOLOGY

9

2757