Environ. Sci. Technol. 2004, 38, 4782-4790
Spectroscopic Evidence for Fe(II)-Fe(III) Electron Transfer at the Iron Oxide-Water Interface AARON G. B. WILLIAMS AND MICHELLE M. SCHERER* Department of Civil and Environmental Engineering, University of Iowa, 4105 Seamans Center, Iowa City, Iowa 52242
Using the isotope specificity of 57Fe Mo¨ ssbauer spectroscopy, we report spectroscopic observations of Fe(II) reacted with oxide surfaces under conditions typical of natural environments (i.e., wet, anoxic, circumneutral pH, and about 1% Fe(II)). Mo¨ ssbauer spectra of Fe(II) adsorbed to rutile (TiO2) and aluminum oxide (Al2O3) show only Fe(II) species, whereas spectra of Fe(II) reacted with goethite (RFeOOH), hematite (R-Fe2O3), and ferrihydrite (Fe5HO8) demonstrate electron transfer between the adsorbed Fe(II) and the underlying iron(III) oxide. Electron-transfer induces growth of an Fe(III) layer on the oxide surface that is similar to the bulk oxide. The resulting oxide is capable of reducing nitrobenzene (as expected based on previous studies), but interestingly, the oxide is only reactive when aqueous Fe(II) is present. This finding suggests a novel pathway for the biogeochemical cycling of Fe and also raises important questions regarding the mechanism of contaminant reduction by Fe(II) in the presence of oxide surfaces.
Introduction Important elemental cycles, such as carbon and phosphorus, are closely linked to the Fe(III)-Fe(II) redox couple, and more recently, Fe redox cycling has been strongly implicated in the mobility of soil and groundwater contaminants (1-4). A growing body of experimental evidence suggests that in oxygen-limited environments, the Fe(III)-Fe(II) redox cycle is driven, in large part, by microbes that can derive energy from the reduction of structural Fe(III) (5, 6). The widespread occurrence of Fe respiring bacteria, such as Geobacter spp. and Shewanella spp., provides a continuous flux of Fe(II) ions at the surface of minerals that may provide an additional pathway for natural attenuation of groundwater contaminant plumes (1, 7-9). The presence of mineral surfaces has been shown to dramatically enhance rates of contaminant reduction by Fe(II) (10-22). The enhanced reactivity of the adsorbed Fe(II) is commonly attributed to surface complexation of Fe(II) by hydroxo ligands on the mineral surface, which stabilizes the Fe(III) oxidation state and lowers the Fe(III)-Fe(II) redox potential (13, 23-25). The identity of the Fe(II) species responsible for contaminant reduction, however, is unclear. In most cases, the kinetics of contaminant reduction are described as proportional to the concentration Fe(II) adsorbed at the mineral surface (determined from the disappearance of Fe(II) from solution or 0.5 M HCl extraction). As more Fe(II) is adsorbed, rates of contaminant reduction * Corresponding author phone: (319)335-5654; fax: (319)335-5660; e-mail:
[email protected]. 4782
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typically increase, consistent with the notion that formation of a reactive Fe(II) surface complex is responsible for the observed contaminant reduction. Other trends in contaminant reduction by adsorbed Fe(II), however, are not adequately described based solely on amount of Fe(II) adsorbed, and other factors such as structure, location, or Fe(II) surface speciation may be important (12). Of these trends, the effect of mineral composition on rates of contaminant reduction is particularly compelling. The reduction of oxamyl by equivalent concentrations of adsorbed Fe(II) on both goethite and hematite exhibit dramatically different reaction rates, suggesting a compositional or structural influence by the underlying oxide. The composition of Fe bearing minerals has also been invoked to explain reaction rates that varied by several orders of magnitude for the reduction of hexachloroethane by adsorbed Fe(II); however, when reaction rates for Fe(II) sorbed on goethite, hematite, and lepidocrocite were normalized for Fe(II) sorption density, the reaction rates were all within a half an order of magnitude (21). Similar rates of Cr(VI) reduction by Fe(II) adsorbed on SiO2 and goethite have also been observed (20). Rates of contaminant reduction by adsorbed Fe(II) measured as a function of solution chemistry, such as pH and Fe(II) concentration, have also revealed several trends that are inconsistent with the concept of only one adsorbed Fe(II) species with the same reactivity. Typically, rates of reduction increase at higher pH (as more Fe(II) adsorbs); however, rates of U(VI) reduction by Fe(II) adsorbed on hematite were not proportional to the amount of Fe(II) adsorbed and the formation of pH-dependent Fe(II) surface complexes, specifically the formation of nonhydroxylated Fe(II) species at pH 6.0 and a monohydroxylated Fe(II) species at pH 7.5, was used to successfully describe the data (17). In addition, modeling of monochloramine and nitrobenzene reduction by adsorbed Fe(II) have implicated multiple Fe(II) species similar to that suggested in work with U(VI) reduction (14, 15). In addition to pH-dependent Fe(II) surface speciation, high aqueous Fe(II) concentrations or high pH values may also result in the formation of surface precipitates and/or highly reactive secondary minerals (26). For example, the formation of a highly reactive blue-green phase, consistent with green rust, was observed at high pH values (>8.0) in suspensions of goethite and Fe(II) and resulted in rapid reduction of fluorotribromomethane (11). Rates of reduction for carbon tetrachloride and polyhalogenated methanes were attributed to potential formation of Fe(OH)2(s) (11, 18, 27). The reduction of nitrate has also been associated with the formation of a reactive secondary magnetite phase in suspensions of Fe(II) and lepidocrocite (10). Wet chemical extraction data have also been used to interpret the nature of the reactive Fe(II) surface complex. More specifically, incomplete recovery of adsorbed Fe(II) with weak to moderate extractions has been suggested to be indirect evidence for the formation of strongly bound, potentially less reactive Fe(II) species, such as insoluble or stable iron oxides. Magnetite is often proposed as the most thermodynamically favorable species formation and is thought to limit Fe(II) recovery even under acidic conditions (28-30). The autocatalytic formation of magnetite by electron transfer from adsorbed Fe(II) has been shown to form from adsorption of Fe(II) on both colloidal ferrihydrite and maghemite (31, 32). Recently, significant advances in spectroscopic techniques have opened a new window for identifying surface species 10.1021/es049373g CCC: $27.50
2004 American Chemical Society Published on Web 08/17/2004
and distinguishing surface complexation from surface precipitation (33). Direct spectroscopic measurements of Fe(II) adsorbed on Fe(III) surfaces, however, remain difficult because of the very small quantity of Fe(II) at the interface and interference from the bulk structural Fe(III) (4, 29). Other spectroscopic techniques that have been used to characterize adsorbed surface species, such as electron paramagnetic resonance (EPR) and X-ray absorption spectroscopy (i.e., EXAFS and XANES), are bulk techniques that cannot distinguish the adsorbed Fe from the underlying structural Fe. Surface sensitive techniques, such as X-ray photoelectron spectroscopy (XPS), that might detect only the surface Fe operate at high vacuum and require dry samples that are not representative of hydrated mineral surfaces found in natural environments. To identify the nature of the reactive Fe(II) surface complex formed on commonly occurring oxides, we used the isotope specificity of 57Fe Mo¨ssbauer spectroscopy to enhance and isolate a spectroscopic signal from Fe(II) reacted with five different oxide surfaces. The specific objectives of our work were (i) to develop a technique to spectroscopically observe Fe(II) reacted with oxide surfaces, (ii) to compare spectroscopic signals from Fe(II) adsorbed on different oxide surfaces, and (iii) to improve our understanding of the pathway of contaminant reduction by Fe(II) in the presence of oxide surfaces.
Methods Synthesis of Isotopically Enriched Fe Solutions. Iron metal isotopically enriched for 56Fe and 57Fe was purchased from Chemgas (Boulogne, France). Reported purities were 96% for 57Fe (compared to a natural abundance of 2.12%) and 99.9% for the 56Fe (compared to a natural abundance of 91.8%). The enriched Fe(0) was converted to an aqueous iron salt for use in both oxide synthesis and as an aqueous Fe(II) source. The enriched Fe(0) was dissolved in 0.5 M HCl by boiling the solution using a double-boiler configuration for approximately 30 min. Insoluble Fe(0) was removed from the solution with a magnet, and then the resulting ferrous chloride solution was filtered through an 0.02 µM filter and stored in 0.1 M HCl in an anoxic glovebox to prevent the oxidation of the Fe(II). Oxide Synthesis. Oxidation of the 56Fe(II) acid solution was achieved by addition of a few milliliters of a 30% H2O2 solution. The 56Fe-oxides were synthesized by modifying the methods of Schwertmann to be able to synthesize small batches of oxides (∼300 mg) (34). Briefly, 56Fe-goethite was synthesized from a 56Fe(III) solution mixed with 25 mL of 5 M KOH and diluted to 250 mL with deionized water. A redbrown precipitate formed immediately. The red-brown precipitate was incubated at 70 °C for 3 days during which it transformed from a red-brown ferrihydrite precipitate to a yellow-brown goethite suspension. Hematite was synthesized from a 56Fe(III) solution heated to 90 °C and rapidly precipitated by the addition of 15 mL of 1 M KOH and 5 mL of 1 M NaHCO3, both preheated to 90 °C. The pH of the solution was adjusted to pH 8-9 with 1 M NaOH or 1 M HCl. The pH-adjusted suspension was incubated at 90 °C for 48 h. Ferrihydrite was formed by rapid precipitation of the 56Fe(III) by rapid addition of 1 M NaOH. The final pH was maintained between 7 and 8 until no further base was required. The resulting 56Fe-oxides were washed with deionized water and freeze-dried to form an oxide powder. All oxides were sieved (100 mesh) before use in experiments to provide a more homogeneous particle size distribution. X-ray diffraction patterns verified that 56Fe-goethite and 56Fehematite were pure, crystalline samples. The method for synthesizing ferrihydrite from 56Fe was validated by making multiple batches with naturally abundant Fe (NAFe) and
confirming ferrihydrite formation with Mo¨ssbauer spectroscopy. Batch Fe(II) Adsorption Experiments. Adsorption experiments were carried out in well-mixed batch reactors containing a mineral suspension buffered at pH 7.4 with 25 mM 4-(2-hydroxyethyl)piperazine-1-ethanesulfonic acid buffer (HEPES). All experiments were conducted inside an anoxic glovebox to avoid exposure to oxygen, which is known to rapidly oxidize Fe(II) in the presence of oxides at circumneutral pH values (35). Care was also taken to avoid localized precipitation of Fe(OH)2(s) due to local high concentrations of OH- or Fe(II) in the reactor. The reactor, before addition of the oxide surface, was allowed to equilibrate after the addition of aqueous 57FeCl2 for a minimum of 2 h. It is important to note that initial experiments where solution pH was controlled by automatic base titration with 50 mM NaOH resulted in the formation of a green rust precipitate that was not visible to the eye. The green rust formed in the presence of each of the five minerals and was an artifact from the locally high pH that developed from the base titration. After equilibration of the 57FeCl2 with the reactor solution, the solution was filtered (0.2 µM) to remove precipitates that may have accumulated from the initial spike of 57FeCl2. Following the filtration of the reactor solution, the aqueous Fe(II) concentration was measured, and a specific oxide mass was added. Reactors contained either 30 mg of the Fe mineral or 500 mg of Al2O3 or TiO2. Each reactor was mixed end-over-end on rotator for 6 h unless otherwise noted in the absence of light. After equilibration, the aqueous Fe(II) concentration was determined by filtering a portion of the oxide slurry. The loss of aqueous Fe(II) was determined by the difference between the amount of Fe(II) prior to the addition of the oxide phase compared to the amount after equilibration. The remainder of the oxide slurry was filtered onto 13 mm filter disks and mounted as a wet paste between two layers of Kapton tape for Mo¨ssbauer analysis. Oxidation of Fe(II) during analysis was not observed as evidenced by the stability of the Fe(II) adsorbed on Al2O3 and TiO2 during consecutive measurements. Mo¨ssbauer spectra were collected in transmission mode with a constant acceleration drive system and a 57Co source. Samples were mounted in either a top-loading Janis cryostat that uses liquid cryogens, liquid nitrogen and liquid helium, or a top-loading Janis exchange-gas cryostat. Both systems maintained the source at room temperature during analysis. Data were calibrated against an R-Fe metal foil collected at room temperature. Spectral fitting was done with the Recoil software package using Voight based spectral lines with a fixed line width (HWHM ) 0.097 mm/s) (University of Ottawa, Ottawa, Canada) (36). Chemical Analyses. Dissolved Fe concentrations were determined on samples filtered (0.22 µm) and immediately acidified with 40 µL of 5 M HCl. Adsorbed Fe(II) was extracted with 0.5 M HCl for 12-18 h. Ferrous iron was measured colorimetrically with the 1,10-phenanthroline method at 510 nm on a UV-Vis spectrophotometer (37). To minimize possible photochemical reduction of the Fe(III)-phenanthroline complex (38), samples for Fe(II) analysis were kept in the dark except for a brief exposure to light (≈1 min) during the addition of the 1,10-phenanthroline and buffer agents. In addition, to account for the background Fe(III) signal from the oxide dissolution, we subtracted the absorbance measured for identical reactors containing the iron(III) oxide, 1,10-phenathroline, and buffer. Samples for nitrobenzene and aniline were filtered though 0.22 µm filters and immediately sealed in autosample vials for analysis. The concentration of nitrobenzene and aniline were determined by HPLC with UV absorbance detection. The eluent consisted of a 60/40 acetonitrile/water mixture and passed through a Supelco LC-18 column at 1 mL/min. VOL. 38, NO. 18, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Experimental Conditions and 4.2 K Mo1 ssbauer Parameters for 57Fe(II) Adsorbed on Al2O3, TiO2, 56Fe-Goethite, 56Fe-Hematite, and 56Fe-Ferrihydrite [57FeII]ads (µg)c
extraction recovery (%)d
surface coverage (%)e
oxide (g/L)
oxide SA (m2/g)a
Al2O3
10
100
nai
180
96(2)
0.8
TiO2
10
9
na
210
nmj
0.9
R-56FeOOH R-56Fe2O3 56Fe HO 5 8
0.6 0.6 0.6
16 20 nm
0.015 0.013 0.016
65 74 250
99 (11) 99 (4) 97 (4)
29 27 6 (est)
oxide
FeII:FeIII b
oxidation state
CS (mm/s)f
QSD (mm/s)g
Fe2+ Fe3+ Fe2+ Fe3+ Fe3+ Fe3+ Fe3+
1.31
2.68
1.32
2.79
0.48 0.49 0.47
-0.1 -0.05 0
H (T)h
area (%)
48.7 53.1 46.9
99+ trace 99+ trace 100 100 100
a Specific surface area was analyzed using five-point BET gas adsorption with N . b Fe(II):Fe(III) ratio based on aqueous Fe(II) initially added 2 compared to the iron oxide mass. c Adsorbed 57Fe(II) based on the difference of wet chemical measurements before and after equilibration with d oxide. Extraction recovery based on amount Fe(II) recovered compared to the amount removed from solution during equilibration. e Surface coverage based on the number of available sites calculated from the oxide surface area and an assumed surface site density of five sites/nm2 compared to the total number of Fe(II) atoms adsorbed based on wet chemical measurements. Surface coverage for ferrihydrite based on an estimated surface area of 300 m2/g. f Center shift. g Quadrupole splitting distribution. h Magnetic splitting; values reported for peak magnetic splitting. i na, not applicable. j nm, not measured.
The absorbance was monitored at 254 nm for nitrobenzene and 235 nm for aniline as determined by a survey scan to determine optimal wavelengths under the experimental conditions. Nitrobenzene and aniline peaks were determined by comparison to the retention times of standards and quantified based on a standard curve.
Results and Discussion Adsorption of Fe(II). Adsorption of aqueous Fe(II) onto three iron oxides including goethite (R-FeOOH), hematite (RFe2O3), and ferrihydrite (Fe5HO8) as well as two minerals that did not contain structural Fe (TiO2 and Al2O3) was measured at pH 7.4 (Table 1). One pH value was selected to avoid changes in Fe(II) speciation due to complexation with hydroxide ions. Higher concentrations of TiO2 and Al2O3 were necessary to ensure that significant adsorption occurred at pH 7.4. The amount of Fe(II) adsorbed was estimated based on the difference between aqueous measurements of initial Fe(II) and Fe(II) remaining after 6 h. Consistent with previous observations, an initial rapid adsorption of Fe(II) occurred within minutes, followed by much slower adsorption (data not shown) (11, 30, 39, 40). Both goethite and hematite, with similar surface areas, adsorbed about 70 µg of Fe(II) (25% of the initial 100 µM), whereas ferrihydrite, typically with a much larger surface area, adsorbed about 250 µg (90%). The higher concentrations of TiO2 and Al2O3 both resulted in about 200 µg (71%) of adsorbed Fe(II), despite significant differences in surface area of the two oxides. The high concentration of adsorbed Fe(II) on TiO2 compared to Al2O3 is likely due to the lower zero point of charge of the TiO2 surface (12). Spectroscopic Measurements of Adsorbed Fe(II). We used 57Fe Mo¨ssbauer spectroscopy to characterize the adsorbed Fe(II). Mo¨ssbauer spectroscopy detects 57Fe only among all other iron isotopes, which allowed us to spectroscopically select for the adsorbed Fe(II) by adding aqueous Fe(II) as 57Fe(II) and synthesizing the underlying iron(III) oxide from 56Fe. By synthesizing the iron(III) oxide from 56Fe, we effectively turned off the Mo¨ssbauer signal originating from the bulk oxide, as demonstrated for goethite in Figure 1A. The Mo¨ssbauer spectrum of naturally abundant goethite (NAFe-goethite) reveals a sextet (six absorption peaks) with a pattern typical of goethite, whereas the spectrum for 56Fegoethite shows little γ-ray absorption (i.e., no absorption peaks). In both cases, XRD was used to confirm the goethite synthesis (Figure 1B). Mo¨ssbauer spectra of Al2O3, TiO2, 56Fehematite, and 56Fe-ferrihydrite also confirmed negligible γ-ray absorption prior to the introduction of 57Fe(II) (data not shown). Mo¨ssbauer spectra collected at room temperature (RT, 295 K) for 57Fe(II) adsorbed on Al2O3 and TiO2 indicate the 4784
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FIGURE 1. Comparison of goethite synthesized from naturally abundant Fe0 and Fe0 isotopically enriched for 56Fe. (A) Mo1 ssbauer spectra of 56Fe-goethite and NAFe-goethite at 295 K. (B) XRD patterns of 56Fe-goethite and NAFe-goethite. formation of an Fe(II) phase, whereas on goethite, hematite, and ferrihydrite, an Fe(III) phase dominates the spectra (Figure 2A). Ferric and ferrous phases are easily distinguishable with Mo¨ssbauer spectroscopy based on center shifts (CS) derived from spectral fitting. The center shift is measured relative to 57Fe in R-Fe metal (which is arbitrarily assigned a CS value of v ) 0). The shift in velocity, based on the center of the Mo¨ssbauer spectrum, from v ) 0 provides information on the electron density at the Fe nucleus and, with one less electron, Fe(III) shifts about 0.3-0.4 mm/s, which is much lower than the 1.0-1.3 mm/s shift typically observed for Fe(II) (41). The center shift for the spectra on Al2O3 and TiO2 are greater than 1.0 mm/s providing clear indication of an Fe(II) phase, whereas, on the iron oxides, the derived CS values range from 0.35 to 0.44 mm/s, indicating an Fe(III) phase (41). Our first inclination was that, despite our extensive precautions, trace oxygen might be responsible for the observed oxidation of the 57Fe(II) to 57Fe(III) on the iron oxides. Near complete dissolution of the mineral suspensions in 0.5 M HCl, however, resulted in 96-99% recovery of the added Fe(II), eliminating the possibility of net oxidation of
FIGURE 2. Mo1 ssbauer spectra of 57Fe(II) reacted with Al2O3, TiO2, 56Fe-goethite, 56Fe-ferrihydrite, and 56Fe-hematite for 6 h. Fe(III) minerals were synthesized from 56Fe, which is not detected by Mo1 ssbauer spectroscopy. Solid circles represent data points and lines connect adjacent data points for a visual aid. (A) Spectra collected at room temperature (RT, 295 K). (B) Spectra collected at liquid helium temperature (LHT, 4.2 K). Experimental conditions: 30 mg of the Fe mineral or 500 mg of Al2O3 or TiO2, 25 mM HEPES buffer at pH 7.4, 25 mM KBr, and an initial aqueous 57Fe(II) concentration of 100 µM (achieved by adding 57FeCl2). Fe(II) by oxygen (Table 1). Recovery of adsorbed Fe(II) after 0.5 M HCl extraction is consistent with previously reported extractions of hematite (28, 29). The complete recovery of Fe(II) observed here and by others, combined with the oxidation of 57Fe(II) to 57Fe(III) observed with Mo¨ssbauer spectroscopy in Figure 2A, suggests that the adsorbed Fe(II) is oxidized by the iron(III) oxide. Fe(II)-Fe(III) Electron Transfer. To test the hypothesis that Fe(III)-Fe(II) electron transfer occurred, we devised an experiment to spectroscopically demonstrate the reduction of structural Fe(III) in the presence of adsorbed Fe(II). We switched the Fe isotopes and synthesized an Fe(III) mineral, specifically ferrihydrite, from 57Fe and exposed it to aqueous Fe(II) made from 56Fe. In this case, Mo¨ssbauer measurements will detect the structural 57Fe(III) rather than the adsorbed 56Fe(II). Before exposure to aqueous 56Fe(II), the high concentration of 57Fe(III) in the ferrihydrite gives a large, broad doublet consistent with what is expected for ferrihydrite (Figure 3A). Spectra were collected at 140 K to avoid magnetic ordering of ferrihydrite (which may obscure a small Fe(II) peak) and maximize sample rigidity. After reacting with 56Fe(II), however, a spectral line at approximately 2.6 mm/s appears in Figure 3B that is uniquely characteristic of the high-energy line position observed in doublets of ferrous
iron compounds (41, 42). Unfortunately, the large background signal from the 57Fe-ferrihydrite obscures the doublet’s lowenergy peak and precludes fitting the spectrum and further characterizing the Fe(II) phase or ratio of Fe(II) to Fe(III). Appearance of the high-energy Fe(II) line, however, clearly indicates reduction of the structural 57Fe(III) in ferrihydrite and, in conjunction with oxidation of 57Fe(II) to 57Fe(III) observed in Figure 2, provides direct evidence for electron transfer between adsorbed Fe(II) and an iron(III) oxide. Our results provide a molecular level explanation for the growing body of macroscopic observations that show markedly different behavior for Fe(II) adsorption compared to other metals (4, 29, 30, 43). Adsorption of metal ions at the mineral-water interface is typically described by surface complexation reactions at low metal concentrations and precipitation of a secondary mineral phase or formation of a solid-solution at high metal concentrations (26, 39, 44). Surface complexation models (SCMs) have been reasonably successful at describing the adsorption of metals on mineral surfaces, with many of the derived surface species confirmed by recently available spectroscopic techniques (45). The incomplete recovery of Fe(II) with mild extractions and formation of new mineral phases observed for Fe(II) in the presence of Fe(III), however, has led many to invoke VOL. 38, NO. 18, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 3. Mo1 ssbauer spectra of 57Fe-ferrihydrite before and after exposure to aqueous 56Fe(II). (A) 57Fe-ferrihydrite before exposure to 56FeII. (B) 57Fe-ferrihydrite after exposure to 100 µm 56FeII for 6 h. Spectra were collected at 140 K. Experimental conditions: 10 mg of 57Fe-ferrihydrite, initial aqueous 56Fe(II) concentration of 100 µM (achieved by adding 56Fe(II)), 25 mM HEPES buffer at pH 7.4, 25 mM KBr. additional processes or reactions to explain their data (11, 19, 29-31, 46-49). Diffusion of Fe(II) into pores, conversion of Fe(II) to more stable or insoluble phases, and electron transfer have all been offered as potential explanations. Here, under conditions typical of natural environments (i.e., wet, anoxic, circumneutral pH, and about 1% Fe(II)), we show that electron transfer occurs. Identification of the Oxidized Fe(III) Phase. A closer look at the room temperature spectra of the oxidized Fe(III) phases in Figure 2A reveals distinct differences between the Fe(III) phases formed on hematite, goethite, and ferrihydrite. On hematite and goethite, some magnetic ordering is observed based on the appearance of a sextet that is well-defined on hematite and only partially resolved on goethite. In contrast, on ferrihydrite, the ferric phase appears as a paramagnetic doublet. Positive identification based on the RT spectra is difficult, and additional spectra were collected at liquid helium temperature (LHT, 4.2 K) to help identify the Fe(III) phases. Lowering the temperature serves several purposes, including slowing down thermal fluctuations of the magnetic structure (which helps observe magnetic order), eliminating dynamic effects, and increasing the rigidity of the Fe environment. As shown in Figure 2B, spectral differences for 57Fe(II) adsorbed on iron(III) oxides as compared to Al O 2 3 and TiO2 are even more striking at LHT than RT. Consistent with the RT spectra, clearly identifiable 57Fe(II) doublets with CS values near 1.3 mm/s are observed on Al2O3 and TiO2 at LHT. A spectrum of a frozen aqueous solution containing KBr and FeCl2 confirmed that frozen Fe(II) in the sample pore water is not interfering with the adsorbed Fe(II) signal. Although it is expected that Fe(II) spectral areas will increase at lower temperatures (50), a similar increase in spectral area (almost 6-fold) has also been observed for adsorbed Fe(II) in a clay-interlayer and interpreted as being due to ordering of water molecules (51). The presence of a doublet at LHT, rather than a magnetically ordered sextet or octet, is also unusual for Fe phases and suggests that the Fe(II) phase lacks the high degree of Fe-Fe interactions that would be expected if a secondary surface precipitate had formed, such as Fe(OH)2(s) (41, 42). Both of these observations, (i.e., lack of rigidity and lower Fe-Fe interactions) seem consistent with the notion of an Fe(II) species adsorbed at the surface. It should be noted, however, that the experiments are at fairly low concentrations of Fe4786
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FIGURE 4. Effect of temperature on Mo1 ssbauer spectra of 57Fe(II) reacted with 56Fe-ferrihydrite. Solid circles represent data points and lines connect adjacent data points for a visual aid. Experimental conditions: 30 mg of the 56Fe-ferrihydrite, 25 mM HEPES buffer at pH 7.4, 25 mM KBr, and an initial aqueous 57Fe(II) concentration of 100 µM (achieved by adding 57FeCl2). (II) (∼100 µM) and low surface coverage (∼1%). It seems plausible that, at higher concentrations and surface coverage, some surface precipitation may occur. In contrast to Al2O3 and TiO2, the LHT spectra from iron bearing oxides show exclusive 57Fe(III) signatures with CS values near 0.48 mm/s (Figure 2B). Spectral identification of the oxidized Fe(III) phase indicates growth of an Fe(III) layer on the Fe mineral surfaces that is physically similar but not identical to the underlying bulk mineral. Spectral identification is based on a combination of CS and hyperfine fields (H) derived from spectral fitting (Table 1). At LHT, the hyperfine field for the oxidized Fe(III) phase on hematite (H ) 54.17 mm/s) and goethite (H ) 48.7 mm/s) are slightly smaller than those expected for bulk hematite (Hbulk ) 53.5 mm/s) and goethite (Hbulk ) 50.6 mm/s) but consistent with the small decrease in the hyperfine field of surface atoms (52, 53). At RT, there is a slight relaxation in the hyperfine fields of about 10%, which suggests that the 57Fe(III) exists as surface layer with reduced magnetic splitting due to fewer neighboring magnetic metal ions (54). In addition, the quadrupole splitting distributions (QSD) on both hematite and goethite deviate from those expected for the bulk minerals (52). A similar observation was made for coatings of 57Fe-hematite deposited on 56Fe-hematite, but the deviation was significantly less (55). Presumably, the perturbed QSDs reflect the unique coordination environment of the oxidized Fe(III) phase at the oxide-water interface, but further interpretation requires a better understanding of QSDs that may only be obtainable with more detailed electronic structure calculations (56). Unlike hematite and goethite, ferrihydrite exhibits a doublet at RT and a broad sextet at LHT (Figure 2B). The broadness of the spectral lines, even at LHT, suggests a range of different Fe environments, which is typical of the more highly disordered ferrihydrite (57). At 4.2 K the outer line half widths are about 1.5 mm/s, consistent with the reported range of 0.9-1.7 mm/s used as an identifying characteristic of ferrihydrite at 4.2 K (57). The smaller hyperfine field is also consistent with that expected for ferrihydrite; however, the broadness of the peaks makes it difficult to positively identify the spectrum. To gain more information on the oxidized
FIGURE 5. Reduction of nitrobenzene by Fe(II) “adsorbed” on goethite in the absence and presence of aqueous Fe(II). Experimental conditions: 30 mg of the goethite, 25 mM HEPES buffer at pH 7.4, 25 mM KBr, and an initial aqueous Fe(II) concentration of 1 mM (achieved by adding FeCl2). Control experiments confirmed that no ArNO2 was removed from solution in the presence of goethite or aqueous Fe(II) alone. phase on ferrihydrite, additional spectra were collected over a wider temperature range (Figure 4). As the temperature is lowered from RT to 30 K, thermal fluctuations of the magnetic structure decrease and magnetic order starts to appear in the form of an unresolved sextet, although the ordering extends over a wide range of temperatures consistent with ferrihydrite (57). The appearance of the magnetic order at less than 40 K is also consistent with the characteristic ordering temperature of ferrihydrite; however, it is important to recognize that particle size can affect ordering temperatures (52, 57, 58). The four lines of evidence (i.e., CS and QSD at RT, H at LHT, a broad sextet at LHT, and a low ordering temperature) however provide a compelling argument to support the growth of ferrihydrite on ferrihydrite via oxidation of adsorbed Fe(II) by structural Fe(III). Growth of an Fe(III) layer on ferrihydrite, hematite, and goethite that is similar to the bulk oxide is not consistent with the prevailing notion that magnetite is the most likely end product from Fe(III)-Fe(II) interfacial electron transfer (10, 29-32). These results also differ from the previously observed catalytic conversion of ferrihydrite to goethite in the presence of iron reducing bacteria and low Fe(II) concentrations (59, 60). It is important to note that by using 57Fe Mo ¨ ssbauer spectroscopy and oxides synthesized from 56Fe, we are characterizing only the Fe that adsorbs and becomes oxidized, whereas many of the previous studies used techniques that average structural changes over the bulk or near-surface regions. In addition, the end product will most likely be strongly influenced by experimental conditions including pH, mixing conditions (flow or batch), water chemistry (e.g., Fe(II) concentration, oxide sorbent, carbonate and phosphate concentrations, etc.), and equilibration time. For this study, we chose reasonably simple conditions, but we are currently evaluating the end products formed under a variety of experimental conditions. Fate of the Electron. At this point in the study we can reasonably identify the oxidized ferric phases, but we can only speculate on what happens to the electrons transferred from the adsorbed Fe(II) to the Fe(III) mineral. On the basis of Figure 3B, it is evident that electrons are transferred to the 57Fe-ferrihydrite to create ferrous atoms in an environment
rigid enough to give a measurable Mo¨ssbauer effect. Presumably, the electrons transfer to Fe(III) atoms in a manner similar to the Fe(III)-Fe(II) charge transfer observed in mixed-valence Fe minerals, such as magnetite (61, 62). Recently, a similar electron-hopping mechanism was used to construct an ab initio model to describe the transport of electrons through hematite (63). Results from the model suggest that electrons injected into hematite remain localized enough to create momentarily distinct Fe(II) atoms rather than delocalizing into the conduction band. Our results support the formation of a distinct, solidbound Fe(II) species, but it is unclear whether the electrons remain at the surface or migrate into the bulk mineral, which may have important implications for contaminant fate in the presence of Fe(II) and oxides. Radiolytic reduction of hematite has been shown to yield substantially less Fe(II) then expected and led the authors to suggest that some charge migration into the interior of the mineral must take place in addition to the expected reductive dissolution (64). We, as well as others, have measured low Fe(II) recoveries after mild chemical extractions, which is consistent with charge trapping in the bulk of the mineral (29, 60). Charge migration is typically thought to proceed through neighbor to neighbor interactions, possibly through bridging oxygen bonds, but electron transfer between spatially separated redox species is also possible on materials with semiconductor capabilities, as was recently demonstrated for the oxidation of galena by Fe(III) (65). Implications for Contaminant Reduction by “Adsorbed” Fe(II). Fe(II)-Fe(III) electron transfer at Fe mineral surfaces raises important questions regarding the mechanism of contaminant reduction by Fe(II) in the presence of oxide surfaces. The most compelling of which is what is responsible for the observed contaminant reduction? If the initially adsorbed Fe(II) becomes oxidized to Fe(III), is the newly formed Fe(II) species [i.e., tFe(III)sOsFe(II) f tFe(II)s OsFe(III)] located at the surface or in the bulk of the mineral? If it is located in the bulk of the mineral, can contaminant reduction still occur at the oxide surface? Clearly there are numerous studies that have demonstrated contaminant reduction in the presence of Fe(II) and oxide surfaces, yet VOL. 38, NO. 18, 2004 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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there have been some observations that exposure to high concentrations of Fe(II) or exposure for long times results in significantly slower rates of contaminant reduction (3, 30). Indeed, some have even speculated that “armoring” of the surface might be occurring due to Fe(II)-Fe(III) electron transfer (3, 30). To address whether the oxide can reduce an oxidized contaminant after Fe(III)-Fe(II) electron transfer has occurred, we measured the reduction of nitrobenzene (ArNO2) after exposing goethite to aqueous Fe(II). We wanted to eliminate the presence of any Fe(II) other than the Fe(II) formed from reduction of structural Fe(III). To do this, we divided a single batch of goethite exposed to aqueous Fe(II) for 3 h into two equal batches. One batch was filtered and resuspended in a buffer-electrolyte solution without aqueous Fe(II), and the other batch was left in the buffer-electrolyte solution with aqueous Fe(II) remaining. As expected based on previous studies (15), in the presence of goethite and aqueous Fe(II), ArNO2 was rapidly reduced with stoichiometric production of aniline (Figure 5). In the absence of aqueous Fe(II), however, no ArNO2 reduction was observed. To see whether it was indeed the presence of aqueous Fe(II) that determined whether ArNO2 reduction occurred (and not an artifact of filtering), we added aqueous Fe(II) back to the batch without any aqueous Fe(II) after about 1 h. Almost immediate reduction of ArNO2 occurred at a rate similar to that observed in the reactor with aqueous Fe(II) present initially. The negligible ArNO2 reduction observed in the absence of aqueous Fe(II) suggests that the newly formed Fe(II) species resulting from Fe(III)-Fe(II) electron transfer is not accessible to reduce ArNO2 at the oxide surface. Most experiments measuring contaminant reduction by adsorbed Fe(II) have been done in the presence of aqueous Fe(II), and it has been previously demonstrated that aqueous Fe(II) is necessary to sustain reduction by regenerating the adsorbed Fe(II) pool (11, 18). Our results, however, seem to suggest that aqueous Fe(II) is required for any ArNO2 reduction to occur, not just to sustain reduction after the adsorbed Fe(II) pool has been consumed. It is not clear, however, how the aqueous Fe(II) participates in reduction of ArNO2. On the basis of control experiments, we know that aqueous Fe(II) in the absence of goethite (at pH 7.4) does not reduce ArNO2; therefore, it seems reasonable to conclude that the aqueous Fe(II) must first associate with the surface in some manner. At this point, we can only speculate that there may be some type of competition or kinetic limitations that control whether the adsorbed Fe(II) injects an electron into the oxide or reduces an oxidized contaminant at the oxide surface. Measuring the rates of reduction of contaminants with different rates as a function of Fe(II) concentration may help provide additional insight. Direct spectroscopic data has shown that Fe(II)-Fe(III) electron transfer occurs at the iron oxide-water interface under conditions typical of many natural environments. Electron transfer between adsorbed Fe(II) and Fe(III) oxides is not surprising given the evidence for Fe(III)-Fe(II) electron transfer in mixed-valent Fe minerals, Fe colloid chemistry, and phosphate-catalyzed Fe(II)-Fe(III) electron transfer in mammalian and bacterial ferritin (64, 66-68). In addition, electron transfer has been suggested as a potential explanation for isotopic fractionations observed during dissimilatory Fe(III) reduction and accompanying Fe(II) sorption (69, 70). Molecular modeling of mixed-valent Fe minerals, such as magnetite, suggest that electron transfer is facilitated by the large overlap of the Fe(III)-Fe(II) octahedral d-orbitals found in an edge sharing coordination environment (62). Many divalent metals form bidentate surface complexes that have a similar edge sharing coordination environment when they adsorb to iron oxides (71). If the environment of an Fe(II) ion adsorbed on an Fe(III) mineral is similar to that of adjacent 4788
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Fe(III)-Fe(II) octahedra in mixed-valent Fe minerals, then interfacial electron transfer seems highly plausible. It has also been shown that electron transfer induces growth of an Fe(III) layer on the oxide surface that is similar to the bulk oxide. These findings have important implications for the fate of environmental pollutants as well as geochemical cycling of trace elements, corrosion control, and catalytic industrial processes and require a paradigm shift in our conceptualization of Fe(II) surface-mediated reactions. More specifically, our results suggest that the assumption of static, passive surfaces inherent in most of our current models describing adsorption of redox active ions at mineral surfaces is not appropriate and that more explicit inclusion of the dynamics of surface reactions must be included.
Acknowledgments We thank Denis G. Rancourt (University of Ottawa, Department of Physics) for helpful discussions regarding interpretation of the Mo¨ssbauer spectra. We also thank the Central Microscopy Research Facility at The University of Iowa for help with Mo¨ssbauer measurements. This work was supported by grants from the National Science Foundation (BES9983719) and the American Chemical Society Petroleum Research Fund (38774-AC5S).
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Received for review April 26, 2004. Revised manuscript received June 29, 2004. Accepted July 5, 2004. ES049373G
