Formation of Ternary Metal-Oxalate Surface Complexes on α-FeOOH

Sep 19, 2011 - Processes at the aqueous interfaces of metal (hydr)oxide particles ..... of heavy backscattering atoms or new strong multiple scatterin...
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Formation of Ternary Metal-Oxalate Surface Complexes on α-FeOOH Particles Anna A. Simanova,† John S. Loring,†,‡ and Per Persson*,† †

Department of Chemistry, Umea University, Umea 90187, Sweden ‡ Pacific Northwest National Laboratory, Richland, Washington 99352, United States

bS Supporting Information ABSTRACT: Processes at the aqueous interfaces of metal (hydr)oxide particles greatly influence the mobility, bioavailability, and reactivity of metal ions and ligands. Here we investigated the time-dependent reactions of oxalate or Me(C2O4)33‑ (Me = Fe(III), Al(III), Ga(III), Co(III)) with goethite in aqueous suspension at pH 4 using attenuated total reflectance infrared (ATR-IR) and extended X-ray absorption fine structure (EXAFS) spectroscopy. The data indicate four coordination modes for oxalate and Fe(C2O4)33‑ adsorbed at the goethite surface: (1) outer-spherically with a hydration shell similar to aqueous ligand; (2) outer-spherically but hydrogen bonded to a surface site; (3) innerspherically to surface iron; (4) inner-spherically within a ternary type A surface complex. In the presence of oxalate, the two outer-sphere complexes form rapidly, but with time these species are partially consumed and the ternary inner-sphere complex is formed as a result of a dissolutionreadsorption process. We propose that iron in these ternary complexes is more labile than iron that is mostly embedded in the lattice. Thus, ternary complexation may play an important role in iron bioavailabilty in the environment. For goethite reacted with Al(C2O4)33‑ or Ga(C2O4)33‑, these four surface complexes are accompanied by an additional Al(III) or Ga(III) ternary oxalate surface complex.

1. INTRODUCTION Small metal (hydr)oxide particles have a profound impact on the fate of metal ions and anionic ligands in natural environments such as soils and aquifers as well as in industrial processes.1 Due to their frequently large surface areas, these particles act as highly efficient sorbents and are able to concentrate substantial amounts of metals or ligands via a range of sorption processes. Furthermore, sorbing ligands may induce dissolution of the metal (hydr)oxides and thus release of metal ions.26 These coupled sorptiondissolution reactions exert a strong control over the distribution and availability of metal ions in aqueous systems, and this has consequently been an area of intense research.711 The present study concerns cooperative and competitive adsorption reactions among metal ions and anionic ligands. The mutual presence of these species in aqueous particle suspensions may result in surface precipitation, independent adsorption at different surface sites or competition for the same surface sites, or adsorption of metalligand complexes where chemical bonding between the metal ion and the ligand persists at the surface. The latter type of surface species is formed through one of three distinct mechanisms: (1) the adsorption of an intact metalligand aqueous complex by electrostatic forces or hydrogen bonding to surface functional groups,1214 (2) the formation of a type A ternary complex where the metal ion (Me) acts as a bridge between the surface (S) and the ligand (L) (SMeL),13,1519 or (3) the formation of a type B ternary complex where instead the ligand bridges the surface and the metal ion (SLMe).1921 For example, it has been suggested that oxalate and Zn(II) form type r 2011 American Chemical Society

A ternary complexes at the surface of hematite nanoparticles and surface precipitates or type B ternary compexes at the surface of hematite microparticles.15 Spectroscopic studies have shown that ethylenediaminetetraacetic acid (EDTA) coadsorbs with metals such as Pd(II), Pb(II), Zn(II), and Ga(III) either as outer-sphere or type A MeEDTA ternary surface complexes.1214,17 Adsorption mechanisms may also be pH-dependent, as shown by the Cu(II)glyphosate complex that adsorbs as a type A species at high pH and type B ternary complex at low pH on goethite.22 This kind of fundamental knowledge is necessary for proper descriptions of adsorption and dissolution processes and as a basis for development of synthetic routes for metalligand surface layers that find use as electrode materials, catalysts, and chemical sensors. The focus of the current research was the interaction of oxalate (C2O42‑) and Me(C2O4)33‑ complexes (Me = Al(III), Co(III), Fe(III), Ga(III)) with α-FeOOH (goethite) particles, and whether ternary surface complexes are formed in these systems. This is of particular interest in the context of ligand-promoted dissolution of iron(III) (hydr)oxides. Previous studies have shown that oxalate can dissolve goethite with no significant net Fe(III) release into solution because the dissolved metal readsorbs as a labile iron-oxalate surface complex. This dissolution readsorption mechanism is an explanation for the so-called Received: June 22, 2011 Revised: September 8, 2011 Published: September 19, 2011 21191

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Table 1. Compositions of Me(III)-Oxalate Stock Solutionsa

a

[Me]/[oxal]

L2‑ HL H2L MeL+ MeL2 MeL33‑

Al-oxalate

1:3

2.8 0

0

0

5.4

91.8

Ga-oxalate

1:3

0.6 0.1

0

0

1.0

98.4

Fe-oxalate

1:3

1.2 0

0

0

2.3

Co-oxalate

1:3 [as K3Co

96.5 100

(C2O4)3 3 3H2O]

[Me(III)]/[oxalate] ratios of 1:3 and total oxalate concentrations of 30 mM. All formation constants were obtained from literature29 and recalculated from I = 1 to I = 0.1 M using the Davies equation. Hydrolysis constants were obtained from literature30 and recalculated from I = 0 to I = 0.1 M.

synergistic dissolution of iron-bearing minerals in the concurrent presence of oxalate and siderophores,11,2327 and we hypothesize that the readsorbed species is a ternary Fe(III)-oxalate surface complex. The present investigation was designed to complement these previous studies and specifically aimed at providing spectroscopic information on the possible formation and structures of ternary Me(III)-oxalate surface complexes. Herein we have combined infrared and extended X-ray absorption fine structure (EXAFS) spectroscopic methods to characterize surface reaction products. These techniques have been shown to be highly complementary and in many ways ideal to study coadsorption reactions between metal ions and ligands.12,13,1518

2. EXPERIMENTAL SECTION 2.1. Chemicals. All experiments were performed at 25 °C, under nitrogen, and in the dark. The aqueous solutions and mineral suspensions were prepared from deionized (Milli-Q Plus, resistivity of 18.2 MΩ 3 cm) and boiled water. Sodium chloride NaCl (Merck, p.a.) dried at 180 °C was used to adjust the ionic strength to 0.1 M Na(Cl) in all solutions and mineral suspensions. The parentheses around Cl indicate that the chloride concentration varied while the sodium concentration was held constant. All pH adjustments were made with standardized 40 mM HCl or 40 mM NaOH solutions, in 100 or 60 mM NaCl, respectively, in order to keep [Na] constant. Stock solutions containing trioxalatoiron(III), trioxalatoaluminum(III), trioxalatogallium(III), and trioxalatocobalt(III) were prepared at total oxalate concentrations of 30 mM and [Me(III)]/ [oxalate] ratios of 1:3. The solutions containing trioxalatoiron(III) and trioxalatoaluminum(III) complexes were made by dissolving weighed amounts of Na2C2O4 (Aldrich) and FeCl3 3 H2O (Aldrich) or Al2O3 3 6H2O (Aldrich), respectively. The solution containing trioxalatogallium(III) was made by mixing a stock solution of Na2C2O4 and a stock gallium(III) solution. The stock gallium(III) solution (144.36 mM, pH ∼ 1.5) was prepared by dissolving GaCl3 (Aldrich, 99.999%) and was standardized using atomic absorption spectrometry (AAS). Calculated equilibrium compositions of stock Al(III)-, Ga(III)-, and Fe(III)-oxalate solutions are listed in Table 1. A stock solution of trioxalatocobalt(III) was prepared by dissolving a weighed amount of K3Co(C2O4)3 3 3H2O (GFS Chemicals, 99%). The composition of this solution is also listed in Table 1 and can be assumed to consist (within the stated purity of the solid) only of MeL33‑, since ligand exchange and thermal decomposition rates are slow for Co(C2O4)33‑ at 25 °C.28

Goethite was synthesized as described previously.31 The resulting particles were identified to be goethite by X-ray powder diffraction. The specific surface area was measured to be 104.9 m2/g using the BET N2 adsorption method. An approximately 10 g/L stock suspension that was acidified to a pH of about 4.3 and purged with nitrogen for 12 h in order to remove dissolved CO2 was used for all experiments in this study. 2.2. pH Measurements. The pH was measured with a glass combination pH electrode (Thermo Orion Ross Electrode) containing 0.1 M NaCl in the outer reference cell. The electrode was calibrated prior to each experiment by titrating a volume of 0.1 M NaCl with a standardized HCl solution. Therefore, the pH values in this study are based on proton concentrations rather than activities. In the infrared experiments described below, pH was held constant at 4.00 ( 0.05 using pH STAT methods and standardized HCl or NaOH solutions. 2.3. Infrared (IR) Spectroscopic Experiments. All infrared (IR) spectra were measured with a Bruker IFS-66v/s Fourier transform infrared (FTIR) spectrometer equipped with a DTGS detector and a water-cooled Globar source. The spectrometer was housed in a room thermostatted at 25.0 ( 0.2 °C. In order to avoid any possible photoinduced reactions, a CdTe filter was employed to eliminate the intense visible radiation from the heliumneon laser of the spectrometer. Spectra of species present at the goethite surface were collected using a simultaneous infrared and potentiometric titration (SIPT) method.23,32 The titrations were performed in a titration vessel thermostatted at 25 ( 0.05 °C and equipped with a propeller stirrer and a combination pH electrode. The titration vessel was connected in a closed loop through fluoroelastomer (Chemsure, Gore Industries) and PTFE tubing with a flowthrough attachment inside the spectrometer. The flow-through attachment was made of inert materials and mounted on a singlereflection, 45°, ZnSe attenuated total reflectance (ATR) accessory (FastIR, Harrick Scientific). Prior to each mineral suspension titration, the ZnSe internal reflection element (IRE) was coated with a thin goethite film (overlayer) by evaporating 0.7 mL of a mineral suspension (∼2 g/L) onto the crystal at 75 °C for 2.5 h under a flow of N2. During the titrations, a propeller that was fixed above the ZnSe IRE of the ATR cell facilitated stirring within the flow-through attachment. A volume of ∼10 g/L stock goethite suspension was pumped in a closed loop from the titration vessel to the flow-through ATR cell. Once the background was deemed stable, a background single-beam spectrum (4096 scans at resolution of 4 cm1) was collected at the pHpzc (9.4), and subsequently the suspension was titrated to pH 4. An aliquot of a stock Me-oxalate solution (see Table 1) was added to bring the total oxalate concentration to 1 μmol/m2. The pH was held constant at pH 4.00 ( 0.05 using pH STAT methods. Immediately after addition, spectra were collected as a function of time with varying numbers of scans (between 1024 and 4096). A spectrum of goethite collected at pH 4 in a separate experiment (no oxalate, and the background spectrum was collected at pH of 9.4) was subtracted from all sample spectra to remove the spectral contributions of the mineral. The resulting spectra were baseline corrected over their entire spectral range by subtracting the average of the absorbances between 1900 cm1 and 1950 cm1. Spectra of aqueous C2O42‑ and Me(III)-oxalate stock solutions (see Table 1) were recorded using the same singlereflection, 45°, ZnSe ATR accessory as for the mineral suspension titrations. The background spectrum was 0.1 M NaCl, and 21192

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The Journal of Physical Chemistry C both sample and background spectra were an average of 4096 scans collected at 4 cm1. 2.4. Extended X-ray Absorption Fine Structure (EXAFS) Spectroscopy. A sample was prepared for EXAFS measurements according to the following procedure. A volume of ∼10 g/L stock goethite suspension was placed into a titration vessel equipped with a combination pH-electrode. The suspension was titrated to pH 4, and an aliquot of stock Ga(III)-oxalate stock solution (see Table 1) was added to bring the total oxalate concentration to 1 μmol/m2. The pH was held constant at pH 4.00 ( 0.05 using pH STAT methods. A sample was withdrawn after a 2 h reaction time and centrifuged to concentrate the mineral particles as a wet paste. A 2 mm thick Teflon cell was loaded with the wet paste, sealed with Kapton tape, and loaded into the experimental hutch. The time between centrifugation and beginning of EXAFS data collection was less than 30 min. The Ga K-edge (10.378 keV) EXAFS spectrum was measured at beamline 4-3 at Stanford Synchrotron Radiation Lightsource (SSRL), California, at 3.0 GeV beam energy and 150200 mA electron current. A Si(111) double-crystal monochromator was used, and higher order harmonics were rejected using a mirror. The incoming and transmitted beam was registered using gasfilled ion chambers. The data were collected at room temperature in fluorescence mode, using a Lytle detector filled with Ar (g). A Zn filter of three absorption lengths and three layers of Al foil were used to reduce unwanted fluorescence and scattering contributions to the signal. Internal calibration was performed by simultaneously measuring spectra from an oxidized Ga foil in transmission mode throughout the duration of all scans. Eighteen scans were collected during a total data collection time of approximately 7 h. The EXAFS data were energy-calibrated and averaged with SixPack33 and further analyzed using Viper.34 Standard procedures were used for pre-edge subtraction and data normalization. In order to isolate the EXAFS function, χ, χ ¼ ðμ  μ0 Þ=μ0 where μ is the measured absorption and μ0 is the atomic-like absorption; μ0 needs to be evaluated. This was accomplished by the method of Bayesian smoothing as realized in the program Viper and following the procedures described by Klementev.34 The isolated EXAFS oscillations were k3-weighted and Fourier transformed over the k-range 2.5511.4 Å1 using a Bessel window function. The data were analyzed in R-space by fitting simultaneously the Fourier transform modulus and the imaginary part of the Fourier transform pair. The theoretical phase and amplitude functions for single and multiple scattering paths within an assumed molecular model were calculated by means of the ab initio X-ray absorption fine structure code FEFF7, and used in the refinements of the experimental data. The input structures were those of Ga(C2O4)33‑ from a previous density functional theory geometry optimization35 and goethite where one Fe was replaced by Ga.36 In addition to the quantitative evaluation of EXAFS data, a qualitative analysis of the nature of backscattering atoms in higher coordination shells was conducted using the wavelet transform (WT) method as implemented in the Igor Pro script developed by Funke et al.37 This method complements the conventional Fourier transform (FT) analysis and reveals the energies where the predominant back scattering takes place that give rise to the FT peaks. The WT results are typically visualized

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Figure 1. Infrared spectra collected as a function of time after addition of aqueous oxalate to a goethite suspension at pH 4 ([C2O42‑]tot = 1 μmol/m2) and the corresponding second derivative spectra (middle). IR spectra of the aqueous C2O42‑ and Fe(C2O4)33‑ species are shown for comparison (bottom).

as contour plots in three dimensions: the wave vector (k), the interatomic distance uncorrected for phase shifts (R), and the WT modulus. The contributions from different backscattering atoms at specific regions in k and R are visualized in these plots as ridges, and the location of the ridges helps to differentiate between light and heavy backscattering atoms. The WT modulus of the adsorption sample was analyzed and compared with that of a reference sample consisting of Ga(C2O4)33‑ (aq). The spectra were analyzed by means of the Morlet-wavelet, and η = 4 and σ = 2 were used in the high-resolution wavelet plots.

3. RESULTS AND DISCUSSION 3.1. Infrared Spectroscopy. Figures 15 show the IR spectra collected as a function of time after addition of oxalate or Me(C2O4)33‑ (Me = Al(III), Co(III), Fe(III), Ga(III)) to a goethite suspension at pH 4 in the dark. The time assigned to a spectrum was the difference between when the spectrum was halfway collected and the reaction-time zero, that is, the time when the addition of oxalate or Me(C2O4)33‑ was made. The IR spectra of C2O42‑ (aq) and Me(C2O4)33‑ (aq) are included for comparison, as they have been shown to be relevant models for outer- and inner-sphere adsorption modes, respectively.7,23,28,29 In order to enhance the resolution of highly overlapped bands, the second derivatives of the spectra were calculated and included in the figures; note that a minimum in the second derivative spectrum corresponds to a maximum in the original IR spectrum. Two-dimensional (2D) correlation analysis was also applied to enhance spectral resolution (see the Supporting Information for details). This method provides additional information on the correlation between band intensity variations as a function of time and is helpful in the assignment of bands belonging to the same surface species and thus in identifying the number of the predominant species.40 Note that the wavenumbers of bands may differ slightly depending on whether the band maximum has been determined from the regular spectra, the second derivative spectra, or the 2D analysis. 21193

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The Journal of Physical Chemistry C 3.1.1. Oxalate Adsorbed on Goethite. Previously, the time dependent features in IR spectra of oxalate adsorbed at the waterboehmite interface have been explained by a slow transition from predominantly outer-sphere to inner-sphere surface coordination.7 The IR spectra of oxalate adsorbed onto goethite presented herein also display a time-dependency (Figure 1), and over time the relative intensities of the bands at 1610, 1558, 1431, 1307, and 1290 cm1 decrease, while the intensities of the bands at 1713, 1693, 1674, 1413, 1394, 1268, and 1254 cm1 increase. It should be noted that the complete collection of bands is only detected in the second derivative spectra and some only appear as shoulders (Figure 1). The 2D correlation analysis presented in the Supporting Information indicates that the bands that are more predominant at the beginning of the experiment (i.e., within the first 2 h based on the intensities in the regular spectra) are due to the presence of two surface species characterized by bands at 1558 and 1307 cm1, and 1610 and 1290 cm1, respectively; the band at 1431 cm1 was not resolved in the 2D analysis. The bands belonging to the former group closely resemble those of the aqueous oxalate dianion at 1569 and 1308 cm1 (Figure 1). Accordingly, we assign these bands to the CO stretching vibrations of an outer-sphere oxalate surface complex. The small shift in the νas CO mode has been explained by a difference in hydrogen bonding to water molecules in solution and at the surface.7 The close similarity to the aqueous species suggests that the D2 symmetry of oxalate is preserved, which implies that it may have an intact first solvation shell and is merely attracted to the surface via electrostatic forces. We tentatively assign the bands at 1610 and 1290 cm1 also to an outer-sphere complex, but in this case the surface coordination is presumably mediated by stronger hydrogen bonding to acidic surface groups acting as proton donors. The rationale behind this assignment is that the band at 1610 cm1 is in close agreement 1 in the spectrum of monowith a νas CO peak at ca. 1610 cm protonated oxalate in aqueous solution (see Figures 15 and Supporting Information Table SI1).39,41,42 In the latter case, the peak most probably originates from the carboxylate group involved in intramolecular hydrogen bonding. Thus, we propose that the presence of this band in the spectrum of adsorbed oxalate is associated with strong hydrogen bonding interactions. This interpretation is in agreement with the fact that νas CO is shifted to substantially higher wavenumbers as compared to aqueous oxalate species (ca. 1560 cm1) but still it is lower than that for the metal-coordinated oxalates (at >1660 cm1). Similar moderate to strong hydrogen bonds have been suggested for other carboxylate surface complexes on goethite in the acidic pH range.43 Since a stronger H-bonding interaction may reduce the symmetry of oxalate, we assign the 1431 cm1 band to this species although it was not detected in the 2D analysis. The surface species growing in with time are characterized by substantially higher νas CO frequencies. These bands appear in a wavenumber region typical for metal-coordinated oxalates (Figure 1) and are thus assigned to inner-sphere surface complexes. The 2D analysis (see the Supporting Information) shows that most likely there are two different surface species and these 1 are characterized by νas CO at 1674 and 1693 cm , respectively. Unfortunately, not all bands are resolved in the 2D plots, and this prevents complete assignment of the bands detected in the second derivative spectra. The presence of two structurally distinct innersphere oxalate surface complexes on goethite has been previously discussed in a study performed under similar experimental conditions (i.e., the same pH and surface coverage).23 It has been

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Figure 2. Infrared spectra of Co(C2O4)33‑ in aqueous solution (bottom) and at the watergoethite interface at pH 4 and [C2O42‑]tot = 1 μmol/m2 (top).

Figure 3. Infrared spectra collected as a function of time after addition of Fe(C2O4)33‑ to a goethite suspension at pH 4 ([C2O42‑]tot = 1 μmol/m2) and the corresponding second derivative spectra (middle). IR spectra of the aqueous C2O42‑ and Fe(C2O4)33‑ species are shown for comparison (bottom).

suggested that one species consists of oxalate coordinated to a structural surface iron and another where oxalate is coordinated to iron in a readsorbed complex formed through a dissolution readsorption process. 3.1.2. MetalOxalate Complexes Adsorbed on Goethite. In order to get more insight to the possible formation and structures of the presumed readsorbed complex, experiments were performed where we added trioxalatometal complexes of Fe(III), Al(III), Ga(III), and Co(III) to goethite suspensions at pH 4 (Figures 25). As Co(III) complexes are characterized by very slow exchange rates of coordinated ligands,28,44 we used Co(C2O4)33‑ (aq) in order to test whether such species can adsorb as intact outer-sphere complexes. This experiment was performed during 5 min only, in order to avoid redox reactions 21194

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Figure 4. Infrared spectra collected as a function of time after addition of Al(C2O4)33‑ to a goethite suspension at pH 4 ([C2O42‑]tot = 1 μmol/m2) and the corresponding second derivative spectra (middle). IR spectra of the aqueous C2O42‑ and Al(C2O4)33‑ species are shown for comparison (bottom).

and the contribution from Co(II). The IR spectrum of Co(C2O4)33‑ adsorbed onto goethite (Figure 2) is very similar to that of the corresponding aqueous species, suggesting that indeed Co(C2O4)33‑ initially adsorbs predominately as an intact outer-sphere surface complex. Thus, the driving force for adsorption is the negative charge of the oxalate complex and the positive surface charge of goethite at pH 4. Unlike Co(C2O4)33‑, adsorbed Fe(C2O4)33‑, Al(C2O4)33‑, and Ga(C2O4)33‑ display IR spectra significantly different from those of the corresponding aqueous species, indicating that the structures of these complexes are not preserved in the adsorbed state (Figures 35). It follows that instead of adsorbing as intact outer-sphere surface complexes, these complexes may be disintegrated, completely or partially, and the oxalate ions may adsorb directly to goethite and/or as part of ternary surface complexes of either type A or type B. The IR spectrum collected after a 2 h reaction of Fe(C2O4)33‑ with goethite closely resembles the spectrum recorded after a 2 day reaction of oxalate with goethite (Figures 1 and 3), suggesting that by the end of both experiments the surface speciation is very similar. This result is consistent with previously published data.23 As shown by the 2D analysis (see the Supporting Information), there seems to be an initial surface species that resembles intact Fe(C2O4)33‑. However, this initial surface species is detected only within the first 10 min of the reaction and then is disintegrated to form the same four surface complexes as in the oxalategoethite system. Hence, the difference in kinetics between the oxalate only and Fe(C2O4)33‑ experiments is likely due to different rates of formation of the ternary Fe(III)oxalate surface complexes, which is influenced by dissolution readsorption reactions in the former case and presumably faster ligand-exchange and adsorption in the latter. The adsorption of Al(C2O4)33‑ onto the goethite particles results in the appearance of additional bands in the regions around 1700 and 1400 cm1 that were not observed in the spectra of adsorbed oxalate or adsorbed Fe(C2O4)33‑ (Figures 1, 3, and 4). In particular, the bands at 1706 and 1726 cm1 are

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Figure 5. Infrared spectra collected as a function of time after addition of aqueous Ga(C2O4)33 to a goethite suspension at pH 4 ([C2O42‑]tot = 1 μmol/m2) and the corresponding second derivative spectra (middle). IR spectra of the aqueous C2O42‑ and Ga(C2O4)33‑ species are shown for comparison (bottom).

noticeable, and these are somewhat shifted from those in the spectrum of the aqueous Al(C2O4)33‑ complex at 1697 and 1722 cm1. Moreover, the 1706 and 1726 cm1 bands grow in with time, which is different from the adsorption behavior of intact Fe(C2O4)33‑ discussed above. These observations suggest the formation of a ternary Al(III)-oxalate surface complex with a structure different from that of an intact outer sphere Al(C2O4)33‑, and indicates that Al(C2O4)33‑ is only partially disintegrated at the surface. This conclusion is corroborated by the 2D analysis (see the Supporting Information) that indicated a new surface species characterized by bands at 1279, 1412, and 1729 cm1. The IR spectra of Ga(C2O4)33‑ adsorbed onto goethite are similar but not identical to the spectra of adsorbed oxalate and Fe(C2O4)33‑, indicating that the Ga(III)-oxalate complex is not completely disintegrated and an additional oxalate species is present at the surface (Figures 1, 3, and 5). The small differences are manifested as a blue shift of the band around 1710 cm1 and also different band shapes in the region between 1250 1300 cm1. The 2D correlation analysis resolves more features and indicates a new Ga(III)-oxalate surface complex characterized by bands at approximately 1263, 1700, and 1720 cm1 (see the Supporting Information), which are slightly different from the bands at 1266, 1690, and 1714 cm1 in the spectrum of aqueous Ga(C2O4)33‑ (see Supporting Information Table SI1). Hence, Ga(C2O4)33‑ adsorbed onto goethite behaves similar to the adsorbed Al(C2O4)33‑, suggesting formation of a ternary Ga(III)oxalate surface complex with a structure different from that of intact outer-sphere Ga(C2O4)33‑. Furthermore, in both systems, these presumed ternary surface complexes coexist with the surface complexes that are formed when oxalate or Fe(C2O4)33‑ adsorb onto goethite as a result of partial disintegration of the Ga(III)oxalate complex (see the Supporting Information). In summary, the IR data indicate that Al(III)-, Fe(III)-, and Ga(III)-oxalate form ternary inner-sphere complexes on goethite. Presumably, these are structural analogues, but based on 21195

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Figure 6. Left panel: k3-weighted EXAFS of (a) Ga(C2O4)33‑ reacted with goethite (blue) and the result from the nonlinear least-squares fit (red), (b) aqueous Ga(C2O4)33‑, (c) gallium(III) adsorbed onto goethite at pH 5.72. Right panel: the corresponding Fourier transforms of the EXAFS spectra displayed in the left panel.

IR data alone we are reluctant decide whether the complexes are of type A (surfacemetalligand) or type B (surfaceligand metal). This question is addressed below using EXAFS spectroscopy and Ga(III)-oxalate as the model system. 3.2. EXAFS Spectroscopy of Ga(III)-Oxalate Adsorbed on Goethite. The two primary objectives with the EXAFS investigation of Ga(C2O4)33‑ adsorbed onto goethite (Figure 6) were to determine whether Ga(III) interacts directly with surface functional groups, and if and how oxalate remains coordinated to Ga(III) in the surface complexes. This information is contained in the Fourier transform features between 2 and 3.5 Å, and thus, proper analysis of these contributions is critical. As a first step, we applied a WT method, which has previously been shown to facilitate qualitative assessment of the nature of backscattering atoms.45 In Figure 7, we compare the WTs of EXAFS data from Ga(C2O4)33‑ in aqueous solution and adsorbed onto goethite, and these display some substantial differences. This result alone indicates that adsorption is not caused by an electrostatic interaction between intact anionic Ga(C2O4)33‑ species and the positively charged goethite surface, which is corroborated by comparison of the Fourier transforms in Figure 6. In analogy with Fe(C2O4)33‑ (aq), the pattern in the WT contour plot of Ga(C2O4)33‑ (aq) is caused by a combination of single and multiple scattering paths.46 The feature in the low left-hand corner at low k and R of Figure 7 is mainly caused by singlescattering from the six C atoms of oxalate surrounding Ga(III). As predicted from the ab initio X-ray absorption fine structure calculations of a Ga(C2O4)33‑ structure, the other features originate primarily from multiple scattering or a combination of single and multiple scattering within the GaOCO fragment. This shows that not only heavy atoms but also contributions from strong multiple scattering paths may display WT signals at high k-values. The WT of Ga(C2O4)33‑ reacted with goethite displays signals characteristic of backscattering from oxalate, for instance, the feature at low k and R indicative of carbon backscattering (Figure 7). Moreover, the surface species exhibits a strong feature at k ≈ 7 Å1 and R ≈ 2.8 Å.

Thus, according to the above, this indicates either the presence of heavy backscattering atoms or new strong multiple scattering paths. Under similar experimental conditions but in the absence of oxalate, previous results have shown that this WT feature originates from GaFe single scattering paths. It is therefore likely that in the present case the signal at k ≈ 7 Å1 and R ≈ 2.8 Å also indicates a close proximity between Ga and Fe.36 Thus, the collective results from the WT analysis are consistent with a structural model where a substantial fraction of Ga(III) at the goethite surface is coordinated to oxalate and also that a substantial fraction of Ga(III) is coordinated inner-spherically with respect to surface functional groups. Guided by the WT results, nonlinear least-squares fitting of the EXAFS data was attempted using models with the main scattering paths from oxalate and from Fe atoms. A model including single-scattering GaO, GaC, and GaFe together with two multiple-scattering paths originating from the first shell and the oxalate structure resulted in a good fit (Figure 6 and Table 2). As expected, Ga(III) is 6-coordinated, presumably in a distorted octahedral geometry, as indicated by the GaO coordination number, the distance, and the comparatively high DebyeWaller factor (Table 2). The GaC coordination number suggests roughly a 1:1 ratio between Ga(III) and oxalate. One important question is whether this ratio reflects the composition of a predominating surface complex or merely the result of a mixture of complexes where some are coordinated to oxalate while others are not. Insight into this question is provided by the obtained fit parameters of the GaFe path where CNGaFe = 1.0 and RGaFe = 3.15 Å suggest the predominance of an edge-sharing coordination geometry between the Ga and Fe octahedra.36 Furthermore, these GaFe parameters are different from those obtained in the absence of oxalate where GaFe distances at ca. 3.05, 3.20, and 3.55 Å were detected.36 This difference is also obvious by qualitative comparison between the EXAFS spectra and the Fourier transforms presented in Figure 6a and c. In k-space, these display substantial differences in the regions 6.58 and 9.511 Å1 and the Fourier transforms show different peak 21196

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Figure 7. High resolution WT modulus displaying the higher coordination shells (η = 4, σ = 2) for aqueous Ga(C2O4)33‑ (left) and Ga(C2O4)33‑ adsorbed onto goethite (right). The brownish-white areas indicate ridges and thus areas of strong contribution to the EXAFS spectra, whereas low amplitude areas are depicted in blue.

Table 2. Results from the Quantitative Modeling of the k3Weighted EXAFS Dataa GaO CN R/Å

σ2

Ga(C2O4)33‑/ 6.1 1.96 0.009

GaCc

GaFed

MS Ie

MS IIf

CN R/Å CN R/Å CN R/Å CN R/Å 1.8 2.67

1

3.15

3.7 3.84 6.1 3.92

Gt Ga(C2O4)33‑ 5.6 1.97 0.0049 5.4 2.75

10.9 4.02

(aq)b a R-factor = 3.21. b Data from literature,35 solution contains 96% Ga(C2O4)33‑. c σ2 fixed to 0.004. d σ2 fixed to 0.006. e GaOC multiple scattering path, CN correlated as 2xCN(GaC) and σ2 fixed to 0.006. f GaOO multiple scattering path, CN = CN(GaO) and σ2 correlated as 2*σ2(GaO).

positions and amplitudes in the range 24 Å. These differences indicate that the concentration of Ga(III) at the surface not coordinated to oxalate is small or even negligible. Hence, a plausible interpretation of the EXAFS results is the existence of one predominating Ga(III)-oxalate ternary surface complex. This complex is formed by interaction between structurally complementary Ga and Fe octahedra and involves one oxalate ligand per Ga. Accordingly, we conclude that Ga(C2O4)33‑ forms type A ternary surface complexes through a process where two oxalate ligands are expelled from the first coordination sphere. An interesting observation is the significant shortening of GaC in the surface complex as compared to the solution complex (Table 2). This suggests a strong Gaoxalate interaction at the surface and may hold a clue to the stability and the dynamics of the ternary Ga(III)-oxalate surface complex. It is also corroborated by the IR data that display consistent blue-shifts by 213 cm1 of the two νas CO of the ternary surface complexes as compared to the solution complexes (see Supporting Information Table SI1), which would indicate a stronger force constant of the CO groups involving the noncoordinated oxygens. This is in accordance with a stronger interaction between Me(III) and oxalate.

4. CONCLUSIONS The collective infrared spectroscopic results of oxalate and Fe(C2O4)33‑ adsorbed at the watergoethite interface showed that at least four spectrally distinct species are formed. We propose

that these oxalate surface species are two outer-sphere and two inner-sphere surface complexes. One of the outer sphere species adsorbs electrostatically with an intact solvation shell, whereas the other is believed to be partially desolvated and to interact with surface functional groups via H-bonding. Moreover, we propose that the inner-sphere complexes consist of one type where oxalate is directly coordinated to structural surface Fe ions, whereas the other is a ternary Fe(III)-oxalate surface complex formed via a dissolutionreadsorption mechanism. When Al(C2O4)33‑ and Ga(C2O4)33‑ adsorb onto goethite, they partially decompose (i.e., some of the oxalate ions are removed from the first coordination sphere) and form ternary surface complexes, which are most probably structural analogues to the ternary Fe(III)-oxalate species. The existence of ternary surface complexes was further corroborated by EXAFS spectroscopy of a Ga(C2O4)33‑/goethite sample. The EXAFS data also revealed that these ternary complexes are of type A where Me(III) acts as a bridge between the ligand and the surface. It is likely that the same structure is formed during the dissolution readsorption process in the presence of oxalate. A plausible reaction mechanism is that oxalate promotes dissolution of surface iron at one high-energy crystalline face or at defect sites. Subsequently, the dissolved species readsorb as a type A ternary surface complex at surface sites of lower energy. Iron in these readsorbed compexes is highly mobile and readily accessible to other complexing agents stronger that oxalate, for example, siderophores. An explanation to the difference in lability between this ternary surface complex and the inner sphere surface complex can be found in the number of bonds between the central Fe and oxygens at the surface. In the inner sphere surface complex, the oxalate ion occupies one edge of the Fe octahedron while surface structural oxygens shared with other structural irons occupy the remaining four positions. Thus, in order to remove Fe from this complex, four surface FeO bonds need to be broken and a vacancy is created. In contrast, the proposed ternary surface complex only shares an edge with a structural surface Fe octahedron, and oxalate and presumably water or OH occupy the other positions. Hence, to desorb this complex, only two surface FeO bonds need to be broken, and this difference may account for the differences in lability. In addition, the stronger Me(III)oxalate interaction in the ternary surface complexes, which was indicated by the IR and Ga K-edge EXAFS results, could potentially also contribute to the increased lability of these surface complexes. 21197

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The Journal of Physical Chemistry C The formation of ternary surface complexes may be not unique for oxalate, but may be representative for other chelating carboxylic acids. The proposed dissolutionreadsorption mechanism suggests that iron mobility and bioavailability is greater than can be estimated based solely on solution phase measurements and that ternary surface complexation must be taken into account when predicting iron bioavailability.

’ ASSOCIATED CONTENT

bS

Supporting Information. Full description of the material. This material is available free of charge via the Internet at http://pubs.acs.org.

’ AUTHOR INFORMATION Corresponding Author

*Telephone: +46-90-786 55 73. E-mail: [email protected].

’ ACKNOWLEDGMENT We thank the staff of Stanford Synchrotron Radiation Lightsource (SSRL), particularly Dr. Matthew Latimer and Dr. John Bargar, for their help and advice. SSRL is operated by the U.S. Department of Energy, Office of Basic Energy Sciences. We also acknowledge the National Institutes of Health, National Center for Research Resources, Biomedical Technology Program, and the Department of Energy Office of Biological and Environmental Research, which support the SSRL Structural Molecular Biology Program whose instrumentation was used for the measurements. The Kempe Foundation is gratefully acknowledged for funding the FT-IR spectrometer. The Swedish Research Council provided financial support for this project. One of us (P.P.) acknowledges financial support from the Wenner-Gren Foundations and the Blaustein Visiting Professorship Fund of the School of Earth Sciences, Stanford University. ’ REFERENCES (1) Brown, G. E.; Henrich, V. E.; Casey, W. H.; Clark, D. L.; Eggleston, C.; Felmy, A.; Goodman, D. W.; Gratzel, M.; Maciel, G.; McCarthy, M. I.; Nealson, K. H.; Sverjensky, D. A.; Toney, M. F.; Zachara, J. M. Chem. Rev. 1999, 99, 77–174. (2) Kraemer, S. M. Aquat. Sci. 2004, 66, 3–18. (3) Schwertmann, U. Plant Soil 1991, 130, 1–25. (4) Drever, J. I.; Stillings, L. L. Colloids Surf., A 1997, 120, 167–181. (5) Stumm, W. Colloids Surf., A 1997, 120, 143–166. (6) Zinder, B.; Furrer, G.; Stumm, W. Geochim. Cosmochim. Acta 1986, 50, 1861–1869. (7) Axe, K.; Persson, P. Geochim. Cosmochim. Acta 2001, 65, 4481–4492. (8) Bhandari, N.; Hausner, D. B.; Kubicki, J. D.; Strongin, D. R. Langmuir 2010, 26, 16246–16253. (9) Johnson, S. B.; Yoon, T. H.; Slowey, A. J.; Brown, G. E. Langmuir 2004, 20, 11480–11492. (10) Noren, K.; Loring, J. S.; Bargar, J. R.; Persson, P. J. Phys. Chem. C 2009, 113, 7762–7771. (11) Reichard, P. U.; Kretzschmar, R.; Kraemer, S. M. Geochim. Cosmochim. Acta 2007, 71, 5635–5650. (12) Bargar, J. R.; Persson, P.; Brown, G. E. Geochim. Cosmochim. Acta 1999, 63, 2957–2969. (13) Kaplun, M.; Nordin, A.; Persson, P. Langmuir 2008, 24, 483–489. (14) Schlegel, M. L.; Manceau, A.; Charlet, L. J. Phys. IV 1997, 7, 823–824.

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