Environ. Sci. Technol. 2009, 43, 1086–1090
Investigation of Wu¨stite (FeO) Dissolution: Implications for Reductive Dissolution of Ferric Oxides JE-HUN JANG* AND SUSAN L. BRANTLEY Center for Environmental Kinetics Analysis, Earth and Environmental Systems Institute, 2217 EES Building, Pennsylvania State University, University Park, Pennsylvania 16802
Received April 11, 2008. Revised manuscript received August 22, 2008. Accepted August 25, 2008.
The pH-dependent dissolution flux of FeO (wu¨stite, a ferrous oxide) was measured in this study; flux ) k{H+}n (mol/m2/s), where k ) 10-4.95 and n ) 0.64. This flux was consistent with theoretical predictions based on the rate of water exchange of hexaaquo Fe2+. Interestingly, when compared to published data, the pH-dependent dissolution flux of FeO defined an upper limit for the reductive dissolution fluxes of iron(III) (oxyhydr)oxides, including bacterial dissimilatory iron reduction (DIR). A wide range of dissolution fluxes across several orders of magnitude has been reported for iron(III) (oxyhydr)oxides in the literature and the fluxes were affected by various experimental variables, e.g., pH, ligands, chemical reductants, and bacteria. We concluded that (i) the reductive dissolution fluxes of iron(III) (oxyhydr)oxides, including bacterial DIR, are ultimately bracketed by the detachment rate of reduced Fe(II) from the surface and (ii) the maximum flux can be approached when the mole fraction of reduced Fe(II) at the surface is close to unity.
Introduction Chemical and biological reductive dissolution of iron(III) (oxyhydr)oxides results in the generation of aqueous and solid-associated Fe(II) and regulates the biogeochemical cycle of iron at redox boundaries (1). Under ambient to hydrothermal conditions, Fe(II) can reductively convert various chemical species into less toxic and less mobile forms (e.g., Tc(VII) (2), Cr(VI) (3), U(VI) (4, 5), nitrobenzene (6), etc.), including CO2 f hydrocarbon conversion (7). Thus, in assessing sustainability of contaminated areas and viability of metal-reducing bacteria, an estimate of the extent of reductive dissolution of iron(III) (oxyhydr)oxides is crucial. In addition, iron(III) (oxyhydr)oxides are representative secondary minerals that fill pores and fractures of the aquifer materials under alternating redox conditions (8), so their chemical changes (i.e., reductive dissolution) can affect the physical properties (e.g., hydraulic conductivity, gas permeability, tolerance to shear stress, etc.) of geological structures where we intend to sequester radioactive wastes and CO2. The reductive dissolution has been assumed to be faster than the nonreductive pathway because release of Fe(II) from * Corresponding author current address: Sandia National Laboratories, Carlsbad Programs Group, 4100 National Parks Highway, Carlsbad, NM 88220; e-mail:
[email protected]; phone: (575) 234-0050; fax: (575) 234-0061. 1086
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the lattice is energetically favored over release of Fe(III) (1, 9). Thus, the rate-determining step for the reductive dissolution of iron(III) (oxyhydr)oxides has been considered to be the reduction of surface Fe(III) to Fe(II) (10). For example, use of biogenic organic ligands that can transfer electrons (e.g., quinone) significantly increased the reductive dissolution rates of iron(III) (oxyhydr)oxides (11, 12). Hence, vast experimental studies have measured the reductive dissolution rate as functions of the identity and concentration of surface complexes of Fe(III) reductant, e.g., refs 1, 9, and 13-15. A few investigators have shown that an increase in the ratio of reductant to iron(III) (oxyhydr)oxides over a specific value could not accelerate the reduction rate further (e.g., ascorbate (1) and cysteine (13)). Previous studies including those dealing with dissimilatory iron reduction (DIR) indicated that, during the reductive dissolution, some Fe(III) at the dissolving sites is reduced to Fe(II) prior to the detachment (16). Plus, incrementally released Fe(II) from the solid phase has been detected during the sequential extractions as a function of the “strength” of the extractants (sodium acetate vs 0.5 M HCl extraction in refs 17 and 18), suggesting the presence of Fe(II) embedded within several layers of the solid phase. Thus, a few top layers of iron(III) (oxyhydr)oxides during reductive dissolution may be envisioned as a solid solution made of both Fe(II) and Fe(III) coordinated to crystalline ligands (O2-, OH-). The activity of a component in a binary solid solution decreases dramatically when it becomes a minor constituent of the solid solution (19, 20). Therefore, the solubility of Fe(II) in the presence of the solid solution is lower than the soluble Fe(II) concentration in equilibrium with pure FeO, which results in less thermodynamic driving force for Fe(II) dissolution. Consequently, we hypothesized that the maximum rate of reductive dissolution of iron(III) (oxyhydr)oxides should be the rate of dissolution of FeO (wu ¨ stite). This hypothesis implicitly assumes that, once the surface Fe(III) is completely replaced by Fe(II), the dissolution rate can no longer increase with the reductant concentration. In this study, we envisioned FeO (wu ¨ stite) to be an analogue for iron(III) (oxyhydr)oxides covered entirely with Fe(II) and measured the flux of pHdependent dissolution of FeO to test our hypothesis.
Experimental Section Commercial wu ¨ stite (FeO, Alfa-Aesar), ground in an O2-free glovebox with an agate mortar and pestle, was characterized by X-ray diffraction (XRD), Mo¨ssbauer spectroscopy, and N2BET. Ground wu ¨ stite was packed in N2-filled containers until use. Both XRD and Mo¨ssbauer spectroscopy indicated that the commercial wu ¨stite contains elemental Fe as an impurity (ca. 6% as Fe). In ground FeO, the elemental Fe existed in pellet form (larger than 0.5 mm in diameter) and was easily removed by sieving (Figure 1). The specific surface area was measured as 0.51 m2/g by N2-BET. Dissolution experiments were conducted in batch reactors (Nalgene, polypropylene) at pH 2.0, 3.5, and 5.8 at solid concentrations of 0.522, 0.634, and 0.731 g/L, respectively, for durations up to 10 h. The reactor volume was 250 mL. NaCl was added to deoxygenated deionized water to make 0.1 M NaCl solutions. Dilute NaOH and HCl were used to adjust the pH. The reactors were shaken thoroughly once every hour plus at the time of sampling and pH adjustments. All experiments were completed in an O2-free glovebox. To keep the solid/solution ratio constant throughout the experiments, aliquots were sampled after agitation and then 10.1021/es8010139 CCC: $40.75
2009 American Chemical Society
Published on Web 01/14/2009
FIGURE 1. Characterization of commercial wu¨stite: (a) X-ray diffraction and (b, c) Mo¨ssbauer spectroscopy analyses; (a, b) wu¨stite as received and (c) after sieving to remove elemental Fe. filtered through a 0.2 µm syringe filter. The filtrates were preserved in 0.5 M HCl and stored in the O2-free glovebox until analysis. The dissolution was defined by the concentration of Fe(II) in the filtrates. The Fe(II) concentration was measured by the ferrozine method (21, 22). To test for the presence of aqueous Fe(III), sample splits were reduced using hydroxylamine hydrochloride (NH2OH · HCl) to convert all Fe(III) (23), if any, to Fe(II), followed again by ferrozine analysis. No incremental Fe(II) was observed after addition of NH2OH · HCl, indicating no net oxidation of Fe(II) during the experiments.
Results and Discussion Dissolution Fluxes of FeO. Aqueous Fe(II) concentrations ([Fe(II)aq]) increased with time (eq 1 and Figure 2) but remained significantly undersaturated with respect to FeO. The values of the saturation index (SI; eq 2) were calculated using Visual MINTEQ (24) to be -10.5 ( 0.1 at pH 2.0, -8.2 ( 0.2 at pH 3.5, and -4.7 ( 0.4 at pH 5.8 at the end of each experiment. K is an equilibrium constant for reaction 1 (19), and IAP refers to the ion activity product of reaction 1 calculated from the aqueous activities of Fe2+ and H+. FeO(s) + 2H+(aq) S Fe2 +(aq) + H2O(l) (K ) 1011.4)
(1)
IAP SI ) log K
(2)
FIGURE 2. Changes in [Fe(II)aq] (a) at pH 2.0 ( 0.05, (b) 3.5 ( 0.1, and (c) 5.8 ( 0.2 at 25 °C in 0.1 M NaCl media. The initial fluxes (F) were determined as the slopes of the dashed lines and the final fluxes as those of the solid lines, where we assumed DF/Dt ≈ 0. The median and range of initial and final fluxes are plotted as symbols and error bars in Figure 3. Release of aqueous Fe(II) decelerated with time in the batch experiments (Figure 2). Such variations with time are common for batch experiments and generally indicate either that the dissolving surface has high-energy sites that are dissolving fast initially or that the solution is approaching equilibrium with the dissolving phase, causing back reactions to occur (25). Given that all experiments were highly undersaturated (refer to the aforementioned values of SI), FeO dissolution was inferred to slow with time and approached constant flux due to “equilibration” of the crystal surface to dissolving conditions. To estimate dissolution fluxes (F in eq 3), we therefore determined the slopes of the dashed lines (Figure 2) connecting the origin and the first Fe(II) concentration as initial rates, which are the maximum dissolution rates at each pH. We also determined the minimum rates as the slopes of the lines fitting through the last three time points, where we assumed constant fluxes (i.e., ∂F/∂t ≈ 0). Rates (mol/L/s) VOL. 43, NO. 4, 2009 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 3. Comparison of the dissolution fluxes of iron(III) (oxyhydr)oxides under various conditions. The key is given in panel d. The dashed gray line is the best linear fit for FeO dissolution flux data measured in this study (solid square symbols and error bars). Raw data are shown in Figure 2. (a) Proton-promoted dissolution fluxes of hematite and goethite. Data from refs 9, 38, and 39. The solid gray line is the best linear regression for the flux data. For comparison purposes, the same line was drawn in (b)-(d). (b) Nonreductive, ligand-promoted dissolution fluxes of hematite, goethite, lepidocrocite, and HFO. Data from refs 9, 10, 14, and 39-41. (c) Chemical reductive dissolution fluxes of hematite, goethite, lepidocrocite, and HFO. Data from refs 1, 9, 13-15, 26, and 38-43. (d) Biological reduction flux of hematite, goethite, and HFO. Data from refs 12, 42, and 44-49.
were normalized to the surface area (m2/L) to calculate fluxes (mol/m2/s, F in eq 3). These values at each pH defined the range of dissolution flux (black squares and error bars in Figure 3). In eq 3, F is the dissolution flux (mol/m2/s), s is the slope of the lines in Figure 2 (rate in mol/L/s), δ is the concentration of FeO in the reactors (g/L), and A is the specific surface area of FeO measured by N2-BET (m2/g). To keep δ constant during the course of the experiments, the reactors were thoroughly agitated during the sampling. F)
s δA
(3)
The pH dependence of FeO dissolution flux was determined as a linear fit through the measured fluxes (gray dashed lines in Figure 3): log F (mol/m2/s) ) -0.64(pH) - 4.95 (r2 ) 0.994)
(4)
Selection Criteria for Iron(III) (Oxyhydr)oxide Dissolution Fluxes from the Literature. To compare the measured dissolution flux of FeO to the dissolution of iron(III) (oxyhydr)oxides, we compiled flux data from the literature (gray and open symbols in Figure 3). No further treatment of rate data was made if the rates were given in surface-area-normalized units such as mol/m2/s (i.e., flux, e.g., ref 26). If the data were summarized only as plots of [Fe(II)aq] vs time, the slopes were converted to surface-area-normalized units using reported surface areas and solid concentrations (eq 3). For hydrous ferric oxide (HFO), 600 m2/g (27, 28) was used for the specific surface area if measurements by wet chemical or gas adsorption techniques were not reported. Some papers described solid concentrations as total [Fe(III)] for HFO experiments; in such cases, the molar mass of HFO was assumed to be 89 g of HFO/ mol of Fe as recommended in ref 27. When the particle size (e.g., diameter) and mineralogy were available instead of the 1088
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measured surface area (e.g., ref 9), a maximum estimate of the flux was calculated by assuming that the BET surface area > geometric surface area of particles and by assuming spherical particles to calculate geometric surface area using the following equation: surface area )
6 φF
(5)
where φ is the diameter and F is the density (e.g., 5.26 g/cm3 for hematite, 4.28 g/cm3 for goethite (29)). Papers where we could not make these corrections because of missing parameters were excluded from the data compilation. We noticed that many bioreduction papers report the total Fe(II) determined by extraction with HCl ([Fe(II)extr]), which is larger than [Fe(II)aq] due to Fe(II) associated by sorption or other fixation mechanisms to the solid (6). In some cases, the total reduction rate was determined by measurement of the consumption rate of the reductants (cysteine (13)) or proton (14). The reduction rates determined from [Fe(II)extr] or reductant consumption is expected to be greater than reductive dissolution rates determined from solely [Fe(II)aq] as separated from the suspension by filtration, centrifugation, or dialysis at pH values when the sorption of Fe(II) is significant. Fe(II) sorption increases from 0 to 100% within 2 pH units with the pH of 50% sorption at ∼8.0 (4, 30-33). Most bioreduction experiments we found were performed at neutral pH values less than pKa1 of Fe2+ by more than 2 pH units (Figure 3d), where Ka1 is the first hydrolysis constant of Fe2+ (pKa1 ) 9.4 (34)). Under such conditions, sorbed Fe2+ does not match aqueous Fe2+ significantly (35). Analysis of Compiled Data. A linear regression of the fluxes of proton-promoted dissolution of iron(III) (oxyhydr)oxides (Figure 3, solid gray lines) indicates that the protonpromoted dissolution of iron(III) (oxyhydr)oxides is the slowest flux measured in the literature. The gray dashed line,
the value for the flux constant ko (mol/m2/s) in eq 8 should lie between 10-5.51 and 10-4.91, depending upon the value of n. Therefore, the following expressions are consistent with the hypothesis by Casey et al. (36): log F ) -0.4(pH) - 5.51
(9)
log F ) -0.7(pH) - 4.91
(10)
or
FIGURE 4. Dissolution fluxes of FeO measured in this study (symbols) at pH 2.0, 3.5, and 5.8 were compared with the theoretical prediction of Casey et al. (36). The gray dashed line is the best linear fit of the data. Black lines stand for the theoretical estimates of the dissolution flux for FeO based on the work of Casey et al. (36), where the logarithms of the dissolution rates of NiO, CoO, MnO, and ZnO were correlated to the logarithm of the first-order rate coefficient for solvent exchange (kexch) around metal cation at pH 2.0. Using kexch for Fe(H2O)62+ at pH 2.0 in ref 36, the flux of FeO dissolution at pH 2.0 was estimated using linear regression (eq 6), and the value for ko in eq 8 was determined.
described by eq 4, is the linear fit for measured FeO dissolution flux in this study. Ligand-promoted (nonreductive) dissolution fluxes of iron(III) (oxyhydr)oxides are considerably faster than protonpromoted fluxes (Figure 3b), consistent with the hypothesis that ligands accelerate detachment of Fe(III) from the lattice. These proton- and ligand-promoted dissolutions (Figure 3a,b) are, however, all slower than the proton-promoted dissolution of FeO. Both abiotic and biotic reductive dissolution fluxes (Figure 3c,d) are well bracketed between the fluxes of protonpromoted dissolution of iron(III) (oxyhydr)oxides and that of FeO. This observation suggests that the proton-promoted dissolution flux of FeO is the upper limit of reductive dissolution of various iron(III) (oxyhydr)oxides. Theoretical Validation of Dissolution Flux for FeO Measured in This Study. Casey et al. (36) demonstrated that log flux (log F) of dissolution of divalent transition metal oxides (NiO, CoO, MnO, and ZnO) varies linearly with log kexch, the first-order rate coefficient describing water exchange from the solvent into the inner coordination sphere of the hydrated ion (e.g., aqueous Mn(H2O)62+). Casey et al. (36) proposed the following equation to describe log F vs log kexch dependence for divalent transition metals at pH 2.0: log F (mol/cm2/s) ) 1.93 log kexch - 22.72 (r2 ) 0.989) (6) The value of kexch for Fe(H2O)62+ at pH 2.0 is 106.5 s-1 (36), allowing the theoretical estimation of the dissolution flux of FeO at pH 2.0: Flux of FeO dissolution at pH 2.0 ) 10-6.31 mol/m2/s (7) The proton-promoted dissolution fluxes of divalent metal oxides are proportional to {H+}n (36): F ) ko{H+}n ) ko × 10-(pH)n
(8)
where {H+} is the activity of the proton in solution, ko is the dissolution flux at pH 0, and n can take on nonintegral values from 0.4 to 0.7. Using the flux value at pH 2.0 (eqs 7 and 8),
The measured flux of dissolution of FeO (eq 4) was within the predicted range defined by eqs 9 and 10 (Figure 4). Surface Model for Iron(III) (Oxyhydr)oxides during Reductive Dissolution. Our results indicated that the reductive dissolution flux for iron(III) (oxyhydr)oxides cannot exceed the flux of pH-dependent dissolution of wu ¨ stite. During the reductive dissolution of iron(III) (oxyhydr)oxides, the surface cannot simply be an ideal truncation of the bulk. As indicated by sequential extractions (17, 18), it is highly possible that the “reductively-dissolving” surface of iron(III) (oxyhydr)oxides resembles a solid solution of Fe(II) and Fe(III). Mixed-valence iron precipitates are sometimes amorphous and control the solubility of iron under alternating redox conditions (37). However, the mole fraction of Fe(II) at the interface cannot achieve unity during the reductive dissolution, since reduction and detachment of Fe(II) from the surface occur in sequence; i.e., Fe(III) is continuously being exposed due to the detachment and reduction is the rate-determining step (25). The concentration of Fe(II) in equilibrium with a solid solution of Fe(II) and Fe(III) oxides is generally lower than the concentration in equilibrium with pure FeO (19, 20), which represents less thermodynamic driving force for Fe(II) detachment. Hence, the reductive dissolution flux of Fe(II) from iron(III) (oxyhydr)oxides should be lower than the dissolution flux of Fe(II) from FeO, as illustrated in Figure 3. Consistently, it is not surprising to observe that proton- and ligand-promoted dissolution of iron(III) (oxyhydr)oxides is much slower than the reductive pathway (Figure 3) since no interfacial reduction is involved (i.e., the mole fraction of Fe(II) at the surface is null).
Acknowledgments We acknowledge support from National Science Foundation Grant No. CHE-0431328 and EAR-0311898. SLB also acknowledges support from the NASA Astrobiology Institute Coop. Agreement NCC2-1057 and from U.S. Department of Energy, Office of Biological and Environmental Research (OBER). We acknowledge Professor Gary L. Catchen in Department of Mechanical and Nuclear Engineering at Pennsylvania State University (PSU) for allowing Je-Hun Jang to use his instrumentation for Mössbauer spectroscopy analysis. We also acknowledge two anonymous reviewers for constructive comments. X-ray diffraction and BET analysis were performed by Material Characterization Laboratory (MCL) at PSU. We are also grateful to Professor David A. Dzombak at Carnegie Mellon University, who served as Associate Editor, for his fair and timely handling of this manuscript.
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