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Adsorption of Organic Matter at Mineral/Water Interfaces: 3. Implications of Surface Dissolution for Adsorption of Oxalate Stephen B. Johnson,† Tae Hyun Yoon,† Aaron J. Slowey,† and Gordon E. Brown, Jr.*,†,‡ Surface & Aqueous Geochemistry Group, Department of Geological & Environmental Sciences, Stanford University, Stanford, California 94305-2115, and Stanford Synchrotron Radiation Laboratory, SLAC, 2575 Sand Hill Road, MS 69, Menlo Park, California 94025 Received June 10, 2004. In Final Form: August 11, 2004 The adsorption of oxalate on a model aluminum oxide, corundum (R-Al2O3), has been examined over a broad range of oxalate concentrations (0.125-25.0 mM) and pH conditions (2-10). In situ attenuated total reflectance Fourier transform infrared (ATR-FTIR) measurements indicate that at low to intermediate concentrations ([oxalate] e 2.50 mM), oxalate adsorbs to corundum predominantly as a bidentate, mononuclear, inner-sphere complex involving both carboxyl groups. Significant contributions from outerspherically bound oxalate and aqueous Ox2- are additionally observed at higher oxalate concentrations. Consistent with the ATR-FTIR findings, macroscopic adsorption data measured for oxalate concentrations of 0.125-2.50 mM can be generally well modeled with a single bidentate, inner-sphere oxalate complex using the charge distribution multisite complexation (CD-MUSIC) model. However, at intermediate oxalate concentrations (0.50 and 1.25 mM) and pH < 5, the extent of oxalate adsorption measured experimentally is found to fall significantly below that predicted by CD-MUSIC simulations. The latter finding is interpreted in terms of competition for oxalate from dissolved Al(III), the formation of which is promoted by the dissolution-enhancing properties of the adsorbed oxalate anion. In accordance with this expectation, increasing concentrations of dissolved Al(III) in solution are found to significantly decrease the extent of oxalate adsorption on corundum under acidic pH conditions, presumably through promoting the formation of Al(III)-oxalate complexes with reduced affinities for the corundum surface compared with the uncomplexed oxalate anion.

1. Introduction Low molecular weight (LMW) organic anions such as oxalate and citrate are commonly implicated in the enhancement of mineral dissolution processes. The dissolution-promoting properties of LMW organic anions arise from their ability to strongly bind to mineral surfaces in an inner-sphere manner (i.e., via ligand exchange reactions with surface hydroxyl groups, generating direct bonds between surface cations and the organic anions), resulting in the formation of multidentate, mononuclear surface complexes.1-5 Polarization and weakening of bridging surface metal-oxygen bonds in positions trans to each metal-anion bond can result, leaving the complexed metal cations susceptible to liberation from the surface through a rate-determining detachment step. Such ligand-enhanced mineral dissolution processes are of significant environmental importance. For example, under iron deficient conditions, plants commonly release organic anions such as citrate and malate into the rhizosphere in order to dissolve otherwise inaccessible (insoluble) iron * To whom correspondence should be addressed. E-mail: gordon@ pangea.stanford.edu. Phone: +1 650 723-9168. Fax: +1 650 7252199. † Stanford University. ‡ Stanford Synchrotron Radiation Laboratory. (1) Furrer, G.; Stumm, W. Geochim. Cosmochim. Acta 1986, 50, 1847. (2) Zinder, B.; Furrer, G.; Stumm, W. Geochim. Cosmochim. Acta 1986, 50, 1861. (3) Stumm, W.; Furrer, G. The dissolution of oxides and aluminum silicates; examples of surface-coordination-controlled kinetics. In Aquatic Surface Chemistry; Stumm, W., Ed.; Wiley-Interscience: New York, 1987; p 197. (4) Stumm, W. Adv. Chem. Ser. 1995, 244, 1. (5) Stumm, W. Colloids Surf., A 1997, 120, 143.

from oxyhydroxide minerals.6,7 Similarly, ligand-promoted dissolution is of industrial importance in processes such as the removal of rusts during metal treatment and textile cleaning processes.8 In addition to their dissolution-enhancing properties, many LMW organic anions are of importance due to their ability to strongly complex metal cations in solution. For example, a variety of plants are known to release LMW anions such as citrate and malate in response to increased concentrations of dissolved aluminum in soil solutions. The formation of stable Al-organic chelates results and serves to greatly reduce the effective aluminum phytotoxicity.7,9,10 Similarly, LMW anions such as citrate are used as stable metal chelation agents and sequestrants in a wide range of industrial applications, including the production of animal feeds, fertilizers, papers, and cosmetics and the electroplating of metals.11 The high stability of dissolved metal complexes formed by many LMW organic anions is of potential significance for adsorption processes on mineral substrates. For example, a number of workers have demonstrated that the interactions of actinide contaminants such as uranium(VI)12-14 (6) Jones, D. L.; Darrah, P. R.; Kochian, L. V. Plant Soil 1996, 180, 57. (7) Jones, D. L. Plant Soil 1998, 205, 25. (8) Sawada, H.; Murakami, T. Oxalic Acid. In Kirk-Othmer Encyclopedia of Chemical Technology Online; Wiley-Interscience, 2000. (9) Ma, J. F. Plant Cell Physiol. 2000, 41, 383. (10) Ma, J. F.; Ryan, P. R.; Delhaize, E. Trends Plant Sci. 2001, 6, 273. (11) Lopez-Garcia, R. Citric Acid. In Kirk-Othmer Encyclopedia of Chemical Technology Online; Wiley-Interscience, 2002. (12) Waite, T. D.; Davis, J. A.; Payne, T. E.; Waychunas, G. A.; Xu, N. Geochim. Cosmochim. Acta 1994, 58, 5465. (13) Duff, M. C.; Amrhein, C. Soil Sci. Soc. Am. J. 1996, 60, 1393.

10.1021/la048559q CCC: $27.50 © 2004 American Chemical Society Published on Web 11/24/2004

Organic Matter at Mineral/Water Interfaces

and neptunium(V)15 with carbonate can result in the formation of stable solution complexes with significantly reduced affinities for a variety of mineral surfaces. Xu and Ji16 have also demonstrated that strong chelators such as oxalate, malonate, and malate dramatically reduce the extent of aluminum adsorption on kaolinite through the formation of stable Al-organic complexes in solution. Similarly, Fein and Brady17 have predicted that under acidic pH conditions, dissolved aluminum may compete with clay mineral surfaces for oxalate, resulting in the partial partitioning of oxalate into solution. The latter finding is particularly interesting in that oxalate itself can generate substantial concentrations of aluminum in solution through promoting the dissolution of aluminumcontaining minerals. Thus, the potential exists for adsorbed, dissolution-enhancing LMW anions to generate solution conditions that promote their own eventual desorption via preferential complexation of dissolved metal cations. To date, the potential for competition between dissolved metal cations and mineral surfaces for strongly chelating, dissolution-enhancing organic substances and the implications for geochemical (surface complexation) models have received little attention in the literature. The aim of the present study is to examine the effects of dissolution of a model aluminum oxide, corundum (R-Al2O3), on its adsorption of a common LMW organic anion, oxalate. Corundum was selected for study due to its widely investigated and well-understood surface chemistry18-28 and dissolution29-31 properties in aqueous media. Corundum is also an important raw material for the manufacture of advanced ceramic components where, interestingly, it has been investigated for use in conjunction with LMW organic acid dispersants.20,32,33 Oxalate was chosen for study due to its prevalence in the environment (as a result of its exudation by fungi34 and plant roots7,35) and its ability (14) Wazne, M.; Korfiatis, G. P.; Meng, X. Environ. Sci. Technol. 2003, 37, 3619. (15) Bertetti, F. P.; Pabalan, R. T.; Almendarez, M. G. In Adsorption of Metals by Geomedia; Jenne, E. A., Ed.; Academic Press: London, 1998; p 132. (16) Xu, R. K.; Ji, G. L. Soil Sci. 2003, 168, 39. (17) Fein, J. B.; Brady, P. V. Chem. Geol. 1995, 121, 11. (18) Sverjensky, D. A.; Sahai, N. Geochim. Cosmochim. Acta 1996, 60, 3773. (19) Boily, J. F.; Fein, J. B. Geochim. Cosmochim. Acta 1996, 60, 2929. (20) Hidber, P. C.; Graule, T. J.; Gauckler, L. J. J. Am. Ceram. Soc. 1996, 79, 1857. (21) Hidber, P. C.; Graule, T. J.; Gauckler, L. J. J. Eur. Ceram. Soc. 1997, 17, 239. (22) Sahai, N.; Sverjensky, D. A. Geochim. Cosmochim. Acta 1997, 61, 2801. (23) Boily, J. F.; Fein, J. B. Chem. Geol. 1998, 148, 157. (24) Johnson, S. B. The Relationship Between the Surface Chemistry and the Shear Yield Stress of Mineral Suspensions. Ph.D. Thesis, The University of Melbourne, Melbourne, Australia, 1998. (25) Johnson, S. B.; Scales, P. J.; Healy, T. W. Langmuir 1999, 15, 2836. (26) Eng, P. J.; Trainor, T. P.; Brown, G. E., Jr.; Waychunas, G. A.; Newville, M.; Sutton, S. R.; Rivers, M. L. Science 2000, 288, 1029. (27) Trainor, T. P.; Eng, P. J.; Brown, G. E., Jr.; Robinson, I. K.; De Santis, M. Surf. Sci. 2002, 496, 238. (28) Criscenti, L. J.; Sverjensky, D. A. J. Colloid Interface Sci. 2002, 253, 329. (29) Carroll-Webb, S. A.; Walther, J. V. Geochim. Cosmochim. Acta 1988, 52, 2609. (30) Samson, S. D.; Stillings, L. L.; Eggleston, C. M. Geochim. Cosmochim. Acta 2000, 64, 3471. (31) Johnson, S. B.; Yoon, T. H.; Kocar, B. D.; Brown, G. E., Jr. Langmuir 2004, 20 (12), 4996. (32) Sacks, M. D.; Tseng, T. Y. J. Am. Ceram. Soc. 1983, 66, 242. (33) Luther, E. P.; Yanez, J. A.; Franks, G. V.; Lange, F. F.; Pearson, D. S. J. Am. Ceram. Soc. 1995, 78, 1495. (34) Gadd, G. M. Adv. Microb. Physiol. 1999, 41, 47. (35) Ryan, P. R.; Delhaize, E.; Jones, D. L. Annu. Rev. Plant Physiol. Plant Mol. Biol. 2001, 52, 527.

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Figure 1. Al(III) speciation diagram calculated for an aqueous solution consisting of 1 mM Al(III) and 1 mM oxalate at an ionic strength (I) of 0.01 M. Formation constants for Al(III)oxalate complexes were obtained by averaging values reported by Nordstrom and May (ref 86) and Sjoberg and Ohman (ref 87) after correction to I ) 0.01 M using both specific interaction theory (ref 80) and the Davies equation (ref 81). Al(III) hydrolysis constants were obtained from Nordstrom and May (ref 86). Oxalate protonation constants were obtained from Smith and Martell (ref 85). All hydrolysis and protonation constants were also corrected to I ) 0.01 M using the Davies equation (ref 81). For ease of viewing, minor Al(III) hydrolysis products, including AlOH2+, AlOH2+, and AlOH3, are not shown, as these contribute maximum Al(III) mole fractions of less than 0.03 for the system shown.

to dramatically enhance the dissolution kinetics of a wide variety of minerals,2,3,36-38 including aluminum (oxy)hydroxides.1,39,40 Importantly, as is shown in Figure 1, oxalate also forms a variety of stable complexes with dissolved Al(III) in aqueous solution, particularly under acidic pH conditions where the dissolution of aluminum (oxy)hydroxides,40-43 including corundum,29-31 is most significant. In this study, the inter-relationships among oxalate adsorption on corundum, the oxalate-promoted dissolution of corundum, and the complexation of dissolved Al(III) by oxalate have been investigated as a function of pH and oxalate concentration. The adsorption mode(s) of oxalate on corundum have been examined in situ using attenuated total reflectance Fourier transform infrared (ATR-FTIR) spectroscopy. The macroscopic (bulk) adsorption properties of oxalate on corundum were then measured and simulated using the charge distribution multisite complexation (CDMUSIC) model. The dissolution of corundum was also probed over the same range of experimental parameters, and the corresponding solution-based Al(III)-oxalate interactions were discerned by undertaking solution complexation calculations. Finally, the effects of dissolved Al(III) on the adsorption of oxalate have been examined and discussed. The results indicate the importance of (36) Drever, J. I.; Stillings, L. L. Colloids Surf., A 1997, 120, 167. (37) Eick, M. J.; Peak, J. D.; Brady, W. D. Soil Sci. Soc. Am. J. 1999, 63, 1133. (38) Duckworth, O. W.; Martin, S. T. Geochim. Cosmochim. Acta 2001, 65, 4289. (39) Hering, J. G. Adv. Chem. Ser. 1995, 244, 95. (40) Yoon, T. H.; Johnson, S. B.; Musgrave, C. B.; Brown, G. E., Jr. Geochim. Cosmochim. Acta 2004, 68, 4505. (41) Baes, C. F.; Mesmer, R. E. The hydrolysis of cations; WileyInterscience: New York, 1976. (42) Wesolowski, D. J.; Palmer, D. A. Geochim. Cosmochim. Acta 1994, 58, 2947. (43) Bethke, C. M. The Geochemist’s Workbench, Release 4.0; University of Illinois, 2002.

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considering solution-based interactions when examining aqueous-based particulate systems containing dissolutionenhancing anions. 2. Experimental Section 2.1. Materials. High purity (>99.99%) AKP-30 corundum with a Brunauer-Emmett-Teller (BET) surface area of 7 m2/g and a density of 3.97 g/cm3 was obtained from Sumitomo Chemical Co., Japan. It was used as received. Puriss grade (g99%) oxalic acid was purchased from Fluka Chemika. Aluminum nitrate nonahydrate (99.997%) and 99.9995% aluminum chloride hexahydrate were purchased from Sigma Aldrich and Alfa Aesar, respectively. Analytical grade KNO3 (for use in adsorption and dissolution experiments) or KCl (for ATR-FTIR measurements) and Milli-Q grade water (resistivity ) 18.2 MΩ cm at 25 °C) were used in all experiments. Solution and suspension pHs were measured using a Denver Instrument model 215 pH meter equipped with a high-performance sleeve junction pH electrode, which was regularly calibrated using standard pH buffer solutions. Stock suspensions consisting of 66.7 g/L corundum and 0.01 M KNO3 (or KCl) were prepared by briefly ultrasonicating the aqueous mixtures using a Branson model 450 digital sonifier equipped with a 0.5 in. horn. All stock suspensions were tumbled end-over-end for approximately 12 h prior to use. Aqueous corundum-KNO3 (or KCl) and corundum-oxalateKNO3 (or KCl) samples for ATR-FTIR, macroscopic adsorption, and dissolution experiments were prepared by diluting corundum stock suspensions with appropriate quantities and concentrations of oxalate and/or KNO3 (or KCl) solutions in polypropylene centrifuge tubes. In all cases, the final corundum concentration was 50 g/L, the oxalate concentration ranged from 0 to 25.0 mM, and the KNO3 or KCl concentration was 0.01 M. The pH of the suspension sample in each tube was then adjusted to the desired value using small volumes of 1 M HNO3, HCl, or KOH, after which the tube was purged with high purity N2 gas and slowly tumbled end-over-end for the desired reaction time (3-336 h). During this time, the suspension pH was periodically measured and, if necessary, readjusted to the target pH condition. Upon completion of equilibration, the suspension samples were centrifuged at 10 000 rpm for 20 min in a Beckman Avanti 30 benchtop centrifuge. The supernatants were finally decanted and passed through 20 nm syringe filters (Whatman Anotop 25 Plus) in order to remove any remaining corundum particles. All corundum-KCl and corundum-oxalate-KCl pastes for ATRFTIR analysis were measured within several hours of the completion of the “equilibration” period. Filtered supernatants were stored in the dark in polypropylene bottles at 4 °C for a maximum of 1 week prior to ion chromatography and inductively coupled plasma spectroscopy analyses. 2.2. ATR-FTIR Spectroscopy Measurements. A Nicolet NEXUS 470 FTIR spectrometer equipped with a deuterated triglycine sulfate (DTGS) detector was used to collect Fourier transform infrared (FTIR) spectra. All suspension samples were analyzed in situ as wet pastes using a horizontal attenuated total reflectance (HATR) attachment and a germanium crystal in a trough configuration (PIKE Technologies). Five hundred scans (with a spectral resolution of 4 cm-1) were taken and averaged for each sample. Data collection and spectral calculations were performed using OMNIC (version 6.0a, Nicolet Instrument Corp.) software. To avoid complications introduced by the strong absorbance of the nitrate anion in the spectral region of interest, 0.01 M KCl was used in place of 0.01 M KNO3 for all solution and suspension samples analyzed. Both solution and paste samples were applied directly to the germanium crystal, and the sample-holding region was sealed with a lid to hinder evaporation during the ATRFTIR measurements. In the case of suspension pastes, a thin layer of filtered supernatant was placed on top of each sample as an additional precaution against unwanted drying. As is typical for aqueous-based ATR-FTIR measurements,44 all spectra collected were dominated by the strong infrared (44) Hind, A. R.; Bhargava, S. K.; McKinnon, A. Adv. Colloid Interface Sci. 2001, 93, 91.

Figure 2. The relative abundances of the fully deprotonated (Ox2-), singly protonated (HOx-), and doubly protonated (H2Ox) forms of oxalate in aqueous solution. Data are for an ionic strength (I) of 0.01 M. Acid dissociation constants were corrected from values reported at I ) 0 M (ref 85) to those applicable at I ) 0.01 M using the Davies equation (ref 81). absorbance of water. For each aqueous oxalate sample, the spectral contribution of water was removed by measuring the spectrum of a 0.01 M KCl solution at the equivalent pH and subtracting this from the aqueous oxalate spectrum. For aqueous Al(III)-oxalate samples, spectra of aqueous Al(III) solutions at equivalent pH conditions were instead subtracted. The bulk water signal was removed from the spectrum of each corundumoxalate-KCl paste by measuring and subtracting the spectrum of the filtered supernatant solution. As we have previously reported,31 however, in addition to the spectral contribution from bulk water, the ATR-FTIR spectra of corundum pastes contained additional strong contributions attributable to both physisorbed water and the hydroxylated corundum surface. To isolate the spectrum of oxalate at the corundum-water interface from the latter contributions, for each corundum-oxalate-KCl sample, the spectrum of a corundum-KCl suspension was measured at the same pH and subtracted from the corundum-oxalate-KCl spectrum. As has been previously found,31 this procedure imparts a negative slope to the extracted spectra at wavenumbers less than ca. 1500 cm-1, with the gradient of the slope increasing significantly as a function of increasing oxalate concentration and decreasing wavenumber. For ease of comparison of different spectra, the baseline slope was corrected using 2-3 baseline sections of different gradients. While undertaking all baseline corrections, great care was taken not to enhance or diminish the magnitude of, or shift the position of, any spectral peaks. 2.3. Bulk Adsorption Measurements. Ion chromatography was utilized to measure the residual concentration of oxalate in filtered supernatants obtained after completion of the equilibration procedure described in section 2.1. Samples were diluted to appropriate concentrations, loaded into 0.5 mL filter-capped vials, and analyzed for oxalate using a Dionex ion chromatograph equipped with an AS40 autosampler, a GP50 gradient pump, an LC30 chromatography oven, and an ED40 electrochemical detector. An IonPac AS9-HC analytical column (Dionex Corp.) and a sodium carbonate eluent were used for all measurements. 2.4. Dissolution Measurements. Concentrations of dissolved aluminum in filtered supernatants were measured by inductively coupled plasma (ICP) spectrometry using a TJA IRIS Advantage/ 1000 Radial ICAP spectrometer equipped with a solid-state charge injection device (CID) detector. Dissolution results presented are the average of measurements performed at four wavelengths (λ ) 308.2, 309.2, 394.4, and 396.1 nm).

3. Results and Discussion 3.1. ATR-FTIR Measurements of Aqueous and Adsorbed Oxalate. The aqueous speciation of the oxalate anion is presented as a function of solution pH in Figure 2. As shown, oxalate can exist as fully deprotonated (Ox2-), singly protonated (HOx-), and doubly protonated (H2Ox) species, each of which possesses a distinctly different ATRFTIR spectrum (Figure 3). The FTIR spectrum of Ox2-,


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Figure 3. ATR-FTIR spectra of the fully deprotonated (Ox2-), singly protonated (HOx-), and doubly protonated (H2Ox) forms of oxalate in aqueous solution. Spectra were measured at pH ) 1.11 (H2Ox), 2.08 (HOx-), and 9.88 (Ox2-), respectively, and deconvoluted using a series of spectral subtractions.

which is believed to possess D2d symmetry in aqueous solution,45 consists of two strong absorbances at 1571 and 1308 cm-1, attributable to the asymmetric (νas) and symmetric (νs) stretches of the deprotonated carboxyl groups, respectively. The spectrum of H2Ox also consists of two major peaks at 1742 and 1227 cm-1, which can be assigned to the CdO stretching and C-O stretching/ bending vibrations, respectively, of the protonated carboxyl groups.46 The FTIR spectrum of HOx- is considerably more complex than those of Ox2- and H2Ox, a finding that is understandable given the lower molecular symmetry of the singly protonated oxalate anion. The above results are generally consistent with a number of previous investigations into the aqueous ATR-FTIR properties of oxalate in aqueous solution.38,40,45-50 ATR-FTIR spectra measured for oxalate adsorbed on corundum at pH ) 5.1 after a reaction time of 72 h are shown in Figure 4. For ease of comparison, all spectra have been normalized against the peak positioned at 1700 ( 2 cm-1. At low surface coverages (Γ e 1.4 µmol/m2), the measured spectra contain four major peaks, positioned at 1721, 1700, 1424, and 1294 cm-1. Similar spectra have been reported in previous studies of oxalate adsorption on aluminum (oxy)hydroxide minerals40,46,49 and have been the subject of extensive quantum chemical modeling. They are consistent with oxalate binding in a side-on, mononuclear, bidentate, inner-sphere fashion,40,46 that is, through a ligand exchange process, leading to the formation of a five-membered chelate ring involving both deprotonated oxalate carboxyl groups and a surface Al(III) cation. Importantly, the vibrational frequencies calculated for oxalate bound in alternate inner-sphere modes (i.e., mononuclear monodentate, end-on mononuclear bidentate involving a single carboxyl group, or side-on binuclear bidentate involving both oxalate carboxyl groups) exhibit substantial deviations from the spectra shown at Γ e 1.4 µmol/m2 in Figure 4.40 Under such experimental conditions, the formation of these alternate inner-spherically bound species is therefore unlikely. (45) Hind, A. R.; Bhargava, S. K.; Van Bronswijk, W.; Grocott, S. C.; Eyer, S. L. Appl. Spectrosc. 1998, 52, 683. (46) Axe, K.; Persson, P. Geochim. Cosmochim. Acta 2001, 65, 4481. (47) Hug, S. J.; Sulzberger, B. Langmuir 1994, 10, 3587. (48) Degenhardt, J.; McQuillan, A. J. Chem. Phys. Lett. 1999, 311, 179. (49) Dobson, K. D.; McQuillan, A. J. Spectrochim. Acta, Part A 1999, 55, 1395. (50) Specht, C. H.; Frimmel, F. H. Phys. Chem. Chem. Phys. 2001, 3, 5444.

Figure 4. ATR-FTIR spectra of oxalate adsorbed to corundum at pH ) 5.1 as a function of oxalate surface coverage, Γ (in µmol/m2). The oxalate concentration is also expressed in millimoles/liter in brackets. All spectra were normalized against the major peak present at 1700 ( 2 cm-1.

At higher surface coverages (Γ g 3.6 µmol/m2), Figure 4 shows that several additional peaks gradually emerge as the concentration of oxalate is increased. For example, the peak at 1294 cm-1 for Γ e 1.4 µmol/m2 broadens significantly at Γ g 3.6 µmol/m2 and shifts to higher wavenumber (1309 cm-1 at Γ g 14 µmol/m2). Similarly, a new peak gradually emerges in the frequency region from 1530 to 1630 cm-1, with its maximum intensity gradually decreasing from 1580 cm-1 at Γ ) 14 µmol/m2 to 1576 cm-1 at Γ ) 28 µmol/m2 and then to 1572 cm-1 at Γ ) 71 µmol/m2. Axe and Persson46 and later Yoon et al.40 have reported similar results for the interaction of oxalate with boehmite (γ-AlOOH) surfaces, particularly at high oxalate concentrations. They attributed these new peaks to oxalate adsorbed via an outer-sphere mechanism at the mineral-solution interface. Outer-sphere binding of organic anions typically gives rise to ATR-FTIR spectra that are similar to those of the anion in aqueous solution, but with slight increases in the wavenumber and width of some peaks.31,51-55 These expectations are consistent with the data of Figure 4, in which the peaks in the vicinity of 1571-1580 and 1294-1308 cm-1 lie in close proximity to the νas and νs vibrations, respectively, of the aqueous Ox2- anion. The peaks for the adsorbed species are, however, noticeably broader than those of aqueous Ox2-. In addition, νas is shifted to higher wavenumber (by up to 9 cm-1) in the adsorbed versus the aqueous spectra. These findings are therefore consistent with the outer-sphere adsorption of oxalate at the corundum-water interface, a process that is most significant at Γ g 14 µmol/m2. The gradual decrease in the position of the νas peak from 1580 cm-1 (at Γ ) 14 µmol/m2) to 1572 cm-1 (Γ ) 71 µmol/m2) is presumably due to an increased contribution from nonadsorbed Ox2- at the highest oxalate concentrations investigated. As the spectra in Figure 4 were obtained (51) Persson, P.; Nordin, J.; Rosenqvist, J.; Lovgren, L.; Ohman, L. O.; Sjoberg, S. J. Colloid Interface Sci. 1998, 206, 252. (52) Nordin, J.; Persson, P.; Nordin, A.; Sjoberg, S. Langmuir 1998, 14, 3655. (53) Boily, J. F.; Persson, P.; Sjoberg, S. Geochim. Cosmochim. Acta 2000, 64, 3453. (54) Boily, J. F.; Nilsson, N.; Persson, P.; Sjoberg, S. Langmuir 2000, 16, 5719. (55) Rosenqvist, J.; Axe, K.; Sjoberg, S.; Persson, P. Colloids Surf., A 2003, 220, 91.

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after subtraction of the filtered supernatant for each corundum-oxalate suspension, an excess of nonadsorbed Ox2- must have existed in the vicinity of the corundum surface (presumably in the diffuse electrical double layer) compared with bulk solution. Figure 4 also shows that as the concentration of oxalate is increased, the peaks at 1700 and 1721 cm-1 at Γ e 1.4 µmol/m2 move to slightly higher wavenumber (to 1703 and 1724 cm-1, respectively, at Γ g 3.6 µmol/m2) and are altered significantly in terms of their relative intensities. In addition, the peak present at 1425 cm-1 at Γ e 1.4 µmol/m2 is gradually shifted to 1419 cm-1 at the highest oxalate concentrations investigated. These peak changes could be caused by slight changes in the binding environment of oxalate as the surface coverage increases, with (1) changes in the nature of the adsorption sites and/or (2) interactions with neighboring adsorbed oxalate anions, leading to slight alterations of the relative strengths and frequencies of the characteristic vibrational modes. Alternatively, such spectral changes could potentially be attributed to spectral contributions from dissolved Al(III)-oxalate species that arise as the concentration of oxalate, and therefore the extent of corundum dissolution, increases (see section 3.3). Again, as the spectra in Figure 4 were obtained after subtraction of the filtered supernatant for each corundum-oxalate suspension, this explanation would require that an excess of such Al(III)oxalate species was present in the vicinity of the corundum surface compared with bulk solution. The potential spectral contributions of dissolved Al(III)-oxalate species can be investigated by examining the ATR-FTIR spectra of the filtered suspension supernatants, which are shown in Figure 5a. For clarity, spectra are only shown for supernatants of suspensions where Γ g 3.6 µmol/m2, as at lower surface coverages, oxalate is fully adsorbed at pH ) 5.1 (see section 3.2). In all cases, Figure 5a shows the gradual emergence of peaks in the vicinity of 1571 and 1308 cm-1, indicating the presence of Ox2-. Figure 5a also shows the presence of four additional peaks in the supernatant spectra (in the vicinity of 1720, 1695, 1400-1420, and 1294-1300 cm-1), consistent with the presence of aqueous Al(III)-oxalate species.46 The data of Figure 5a can be compared with those of Figure 5b, in which the spectra of the aqueous AlOx+, AlOx2-, and AlOx33- species are shown. The relative intensities of the peaks in the vicinity of 1720 and 1695 cm-1 show that the Al(III)-oxalate species in the supernatant correlate more strongly with the AlOx2- and AlOx33- species than with the AlOx+ complex. The peak positions of the supernatant Al(III)-oxalate species are also in better agreement with the AlOx2- and AlOx33species than the AlOx+ complex. Such findings are in agreement with thermodynamic considerations, which predict that at pH ) 5.1, AlOx2- and AlOx33- complexes will dominate the aqueous speciation of oxalate under the Al(III) and oxalate concentrations present in the supernatants (see section 3.3). Interestingly, the presence of the additional peak (at 1274 cm-1) characteristic of AlOx33is not observed in the supernatant spectra and is presumably masked as a result of overlap and convolution of peaks from Ox2-, AlOx2-, and AlOx33-. Importantly, as Figure 5a shows that the intensity of the peak in the vicinity of 1720 cm-1 is substantially lower than that at 1695 cm-1 in all supernatants, contributions from aqueous Al(III)-oxalate species cannot account for the changes in the relative intensities of the 1721-1724 and 1700-1703 cm-1 peaks in Figure 4. As a result, we conclude the spectral changes in the 1721-1724, 17001703, and 1419-1425 cm-1 peaks observed in Figure 4

Johnson et al.

Figure 5. (a) ATR-FTIR spectra of selected supernatants after centrifugation and filtration (through 20 nm syringe filters) to separate them from corundum-oxalate-KCl pastes (shown in Figure 4). The total oxalate concentration used in each sample (including that removed via sorption to corundum) is shown in millimoles/liter. For ease of viewing, all spectra were normalized against the absorbance at 1700 cm-1. (b) ATR-FTIR spectra of various Al(III)-oxalate complexes in solution. Solution conditions favoring the formation of each Al(III)-oxalate complex were calculated using averaged formation constants from Nordstrom and May (ref 86) and Sjoberg and Ohman (ref 87), after correction to appropriate ionic strengths using both specific interaction theory (ref 80) and the Davies equation (ref 81). The conditions utilized were as follows: 50 mM oxalate, 200 mM Al(III), pH ) 3.0 for AlOx+; 100 mM oxalate, 50 mM Al(III), pH ) 3.6 for AlOx2-; and 100 mM oxalate, 30 mM Al(III), pH ) 4.9 for AlOx33-. Unwanted spectral responses from aqueous Ox2- and HOx- were subtracted from Al(III)-oxalate spectra where necessary. For ease of viewing, all spectra were normalized against the maximum absorbance in the vicinity of 1700 cm-1.

are instead due to the slight alterations in the average binding environment of inner-spherically adsorbed oxalate, which arise as a function of increasing surface coverage. Figure 6 shows ATR-FTIR spectra measured for oxalate adsorbed on corundum as a function of pH. Γ ) 1.4 µmol/ m2, the reaction time was 72 h, and the spectra have been normalized against the peak present in the vicinity of 1700 cm-1 in all cases. All spectra in Figure 6 show the presence of four major peaks positioned at 1722 ( 1, 1701 ( 1, 1423 ( 1, and 1291 ( 4 cm-1. These spectra closely resemble those obtained at low oxalate concentrations (Γ e 1.4 µmol/m2) in Figure 4 and again indicate that oxalate adsorbs predominantly in an inner-sphere, mononuclear, bidentate manner involving both deprotonated carboxyl groups. It is interesting to note that at pH ) 2.0, the position of the lowest wavenumber peak (1287 cm-1) is considerably lower than that measured at pH ) 3.5 (1291 cm-1) and pH ) 5.0 and 6.5 (1294 ( 1 cm-1). In addition, a small shoulder is clearly visible on the high wavenumber side of the 1287 cm-1 peak at pH ) 2.0, a contribution


Figure 6. ATR-FTIR spectra of oxalate adsorbed to corundum at Γ ) 1.4 µmol/m2 (0.50 mM) as a function of pH. All spectra were normalized against the major peak present at 1700 ( 2 cm-1.

Figure 7. Macroscopic adsorption properties of oxalate on corundum as a function of oxalate concentration and pH. The corundum concentration was 50 g/L, the background electrolyte was 0.01 M KNO3, and the reaction time was 72 h in all cases. The total oxalate concentration was 0.125 mM, b; 0.250 mM, O; 0.500 mM, 2; 1.25 mM, 4; and 2.50 mM, (. Plotted lines represent the CD-MUSIC model fit to experimental data using the model parameters listed in Table 1 and a single innersphere adsorption complex, tAlOx0.50.5- (with ∆z0 ) 0.4 and ∆zβ ) -0.4).

that was similarly observed at higher oxalate concentrations (not shown). These features cannot be due to a contribution from aqueous-based Al(III)-oxalate species, as oxalate is completely adsorbed to the corundum surface at this concentration (see section 3.2). Instead, they are likely to indicate the presence of one or more additional (very minor) oxalate species at low pH. In summary, the ATR-FTIR results presented in Figures 4-6 indicate that oxalate adsorbs to corundum as an innersphere, mononuclear, bidentate complex at all oxalate concentrations investigated. It additionally adsorbs as an outer-sphere complex at high concentrations, most notably at Γ g 14 µmol/m2. The ATR-FTIR spectra of the corundum-oxalate-KCl pastes also contain significant contributions from aqueous Ox2- at the higher oxalate concentrations investigated. Finally, at pH ) 2.0, additional features in the ATR-FTIR spectra of the corundum-oxalate-KCl pastes indicate the likely presence of one or more additional (very minor) adsorbed oxalate species. 3.2. Macroscopic Adsorption of Oxalate on Corundum. The adsorption properties of oxalate on corundum are shown as a function of both oxalate concentration and pH in Figure 7. The background electrolyte was 0.01 M KNO3 and the reaction time was 72 h in all cases. The oxalate concentrations investigated fall at the lower end

Langmuir, Vol. 20, No. 26, 2004 11485

of the concentrations previously examined by ATR-FTIR spectroscopy in Figure 4 (0.125-2.50 mM, corresponding to Γ ) 0.4-7.1 µmol/m2). At the two lowest oxalate concentrations examined (0.125 and 0.25 mM), oxalate is fully adsorbed below pH ) 7.5 and 7.0, respectively. However, its adsorption decreases rapidly under more basic pH conditions, with the measured extent of adsorption being only 10% by pH ) 10. Analogous results have been previously reported for the adsorption of oxalate on (oxy)hydroxide minerals at low oxalate concentrations.17,56,57 At the intermediate oxalate concentrations investigated (0.50 and 1.25 mM), Figure 7 shows that the adsorption of oxalate is again low in the vicinity of pH ) 10 and increases as the pH is reduced. In these cases, however, adsorption reaches a maximum extent in the vicinity of pH ) 5.1-6.0 for a total oxalate concentration of 0.50 mM and at pH ) 4.1-5.4 for an oxalate concentration of 1.25 mM, below which the extent of adsorption is observed to decrease. Violante et al.58 and Rosenqvist et al.55 have previously reported similar results for the adsorption of oxalate on aluminum oxide precipitates and gibbsite, respectively. At the highest oxalate concentration examined (2.50 mM), Figure 7 shows that the extent of oxalate adsorption systematically increases as a function of decreasing pH, rising from 2.6% at pH ) 10 to 39.8% at pH ) 2.1. Interestingly, a maximum in the adsorption versus pH data, as was observed for total oxalate concentrations of 0.50 and 1.25 mM, is clearly not observed at this higher concentration of oxalate. Similar findings, whereby the extent of adsorption increases as a function of decreasing pH, have been previously reported by a number of workers for the binding of oxalate on aluminum-59 and iron-(oxy)hydroxide57,60-62 minerals. For all oxalate concentrations examined, the extent of oxalate adsorption rapidly approaches zero as the pH progresses above the isoelectric point (iep) of corundum (iep ) 9.4 for AKP-30 corundum in 0.01 M KNO324,25,63). Such findings indicate that despite oxalate adsorbing to corundum in a predominantly inner-sphere manner under the conditions shown in Figure 7 (see section 3.1), its adsorption is low under conditions where a net repulsive electrostatic interaction exists between the negatively charged corundum surface and Ox2- anions. As a result, the chemical (nonelectrostatic) contribution to adsorption appears to be small under moderately basic pH conditions. The macroscopic adsorption data shown in Figure 7 can be mathematically linked with the ATR-FTIR results presented in section 3.1 using a suitable surface complexation model. Due to its ability to account for both mononuclear and multinuclear hydroxyl species that are likely to be present on the surface of hydrated corundum particles,26,27,64 we have elected to utilize the CD-MUSIC model of Hiemstra and co-workers.65-68 The CD-MUSIC (56) Lamy, I.; Djafer, M.; Terce, M. Water, Air, Soil Pollut. 1991, 57-58, 457. (57) Filius, J. D.; Hiemstra, T.; Van Riemsdijk, W. H. J. Colloid Interface Sci. 1997, 195, 368. (58) Violante, A.; Colombo, C.; Buondonno, A. Soil Sci. Soc. Am. J. 1991, 55, 65. (59) Horanyi, G. J. Colloid Interface Sci. 2002, 254, 214. (60) Kallay, N.; Matijevic, E. Langmuir 1985, 1, 195. (61) Zhang, Y.; Kallay, N.; Matijevic, E. Langmuir 1985, 1, 201. (62) Mesuere, K.; Fish, W. Environ. Sci. Technol. 1992, 26, 2357. (63) Johnson, S. B.; Russell, A. S.; Scales, P. J. Colloids Surf., A 1998, 141, 119. (64) Hiemstra, T.; Yong, H.; Van Riemsdijk, W. H. Langmuir 1999, 15, 5942. (65) Hiemstra, T.; Van Riemsdijk, W. H.; Bolt, G. H. J. Colloid Interface Sci. 1989, 133, 91.

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Johnson et al.

model describes the protonation of the surface charge groups of corundum in terms of the reactions

tAlOH0.5- + H+ T tAlOH20.5+ tAl2O- + H+ T tAl2OH0 tAl2OH0 + H+ T tAl2OH2+ tAl3O0.5- + H+ T tAl3OH0.5+

log K1,2 ) 9.9 (1) log K2,1 ) 11.9 (2) log K2,2 ) 0 (3) log K3,1 ) 5.9

(4)

where the given formation constants (log Ki,j) are values calculated by Trainor et al.69 for corundum surface groups comprising i Al(III) and j H moieties, based on the empirical bond valence model of Hiemstra et al.68 Examination of reactions 2 and 3 shows that over the pH range investigated in the present study (2-10), binuclear surface groups are predominantly present in the neutral tAl2OH0 form. As a result, they will not contribute significantly to the corundum surface charge. In addition, the ATR-FTIR results previously presented in section 3.1 indicated that oxalate adsorbs to corundum predominantly in an inner-sphere, mononuclear manner. As a result, binuclear surface groups are not expected to play a significant role in the adsorption of oxalate and have therefore not been considered further in our CDMUSIC simulations. By contrast, examination of reactions 1 and 4 shows that both mononuclear (tAlOH0.5-) and trinuclear (tAl3O0.5-) surface groups will contribute to the surface charging characteristics of corundum in the pH range of 2-10. As oxalate adsorbs to corundum predominantly in an inner-sphere, mononuclear manner (see section 3.1), however, only mononuclear surface groups are expected to interact significantly with oxalate over the range of oxalate concentrations examined in Figure 7. As a result, for the purposes of simplifying CD-MUSIC simulations, both mononuclear and trinuclear groups are assumed to participate in protonation reactions, but only mononuclear groups are assumed to take part in surface complexation reactions with oxalate. Previous transmission electron microscopy (TEM) analyses24 of AKP-30 corundum have shown it to consist of submicron (diameter ) 0.3 µm) sized particles of ill-defined morphology. As a result, the dominant crystal planes comprising the hydrated corundum surface are not obvious, and so the likely concentrations of mononuclear and trinuclear sites could not be directly estimated from crystallographic considerations. Instead, the concentrations of mononuclear and trinuclear groups were used as fitting parameters in CD-MUSIC model simulations. The reasonableness of the relative tAlOH0.5- and tAl3O0.5concentrations obtained was assessed by comparison with previous iep measurements of AKP-30 corundum in 0.01 M KNO3.24,25,63 In addition to reactions 1-4, a complete description of the charging properties of the corundum surface requires (66) Hiemstra, T.; De Wit, J. C. M.; Van Riemsdijk, W. H. J. Colloid Interface Sci. 1989, 133, 105. (67) Hiemstra, T.; Van Riemsdijk, W. H. J. Colloid Interface Sci. 1996, 179, 488. (68) Hiemstra, T.; Venema, P.; Van Riemsdijk, W. H. J. Colloid Interface Sci. 1996, 184, 680. (69) Trainor, T. P.; Templeton, A. S.; Brown, G. E., Jr.; Parks, G. A. Langmuir 2002, 18, 5782.

consideration of the role of the background electrolyte (KNO3) according to the ion-pair formation reactions

tAlOH0.5- + K+ T tAlOH0.5-- - -K+ log KCat ) 0.1 (5) tAl3O0.5- + K+ T tAl3O0.5-- - -K+

log KCat ) 0.1 (6)

tAlOH20.5+ + NO3- T tAlOH20.5+- - -NO3log KAn ) 0.1 (7) tAl3OH0.5+ + NO3- T tAl3OH0.5+- - -NO3log KAn ) 0.1 (8) where log KCat and log KAn are ion-pair formation constants taken from Hiemstra et al.64 for the interaction of indifferent electrolyte ions with aluminum hydroxide minerals. CD-MUSIC model simulations were performed using the optimization program FITEQL 4.070 in the format proposed by Tadanier and Eick,71 using the FITEQL source code modifications undertaken by Gustafsson.72 This treatment required formatting of the electrostatic coefficients in the B matrix to reflect the noninteger charges of the corundum surface groups. Fits to experimental data were evaluated with use of the relationship

V(Y) )

WSOS ) DF

(

∑

)

Ycalc - Yexp sexp npnII - nu

2

(9)

where WSOS is the weighted sum of squares, DF is the degrees of freedom, Ycalc and Yexp are the calculated and experimental data, sexp is the error associated with the experimental data, np is the number of data points, nII is the number of components for which both total and free concentrations are known, and nu is the number of adjustable parameters. Relative and absolute errors in the adsorbed oxalate concentrations were defined as 2% and 1 × 10-6 M, respectively. Errors in pH were left at the default FITEQL values.70 CD-MUSIC model calculations were performed according to the three-plane model, which divides the electrical double layer into surface (or 0), inner-Helmholtz (or β), and outer-Helmholtz (or d) planes with charge densities of σ0, σβ, and σd, respectively. The charge of inner-sphere complexes is typically distributed between the 0- and β-planes,67 while indifferent electrolyte ions are assumed to bind in the d-plane. Unfortunately, the latter assumption led to poor convergence characteristics using FITEQL 4.0 over the wide range of oxalate concentrations and pH conditions examined in this study (even after implementation of the recommended FITEQL code changes71,72). As a result, ion-pair formation reactions were modeled by placing the indifferent electrolyte ions in the β-plane. The effect of this change is not expected to be substantial in light of the low background electrolyte concentration (0.01 M) utilized in all adsorption experiments. Interestingly, (70) Herbelin, A.; Westall, J. FITEQL: A computer program for determination of chemical equilibrium constants from experimental data, Version 4.0; Report 99-01; Department of Chemistry, Oregon State University: Corvallis, OR, 1999. (71) Tadanier, C. J.; Eick, M. J. Soil Sci. Soc. Am. J. 2002, 66, 1505. (72) Gustafsson, J. P. http://www.lwr.kth.se/forskningsprojekt/mow/ fiteql.htm, 2004.


Langmuir, Vol. 20, No. 26, 2004 11487

Table 1. Parameters Used in Simulations of Oxalate Adsorption on Corundum Using the CD-MUSIC Model parameter [tAlOH1/2-] [tAl3O1/2-] CT log KHOx- for Ox2- + H+ T HOxlog KH2Ox for Ox2- + 2H+ T H2Ox

value

source

Site Densities 6.7 sites/nm2 1.7 sites/nm2

optimal model fits

Specific Capacitance 0.92 F/m2 from Boily and Fein (refs 23 and 74) Solution Reactions 4.09 from Smith and Martell (ref 85), corrected to I ) 0.01 M using the Davies equation (ref 81) 5.25

log K3,1 for tAl3O1/2- + H+ T tAl3OH1/2+

Surface Protonation Reactions 9.9 from Trainor et al. (ref 69), based on the semiempirical model of Hiemstra et al. (ref 68) 5.9

log KCat (see reactions 5 and 6) log KAn (see reactions 7 and 8)

Ion-Pair Formation Reactions 0.1 from Hiemstra et al. (ref 64) 0.1

log K1,2 for tAlOH1/2- + H+ T tAlOH21/2+

Oxalate Adsorption Reactions log K for tAlOH0.5- + H+ + 0.5Ox2- T tAlOx0.50.510.22 optimal model fit

placement of background electrolyte ions in the β-plane is consistent with the recent finding by Bargar et al.73 that some so-called indifferent electrolyte ions, such as nitrate, participate in the formation of ternary surface complexes with metal ions such as uranyl and thus do not behave as indifferent electrolytes in such systems. The total capacitance of the Stern layer region, CT, was modeled using an outer-Helmholtz capacitance, C2, of 5.00 F m-2 and an inner-Helmholtz capacitance, C1, of 1.13 F m-2, such that CT ()(1/C1 + 1/C2)-1) was equal to 0.92 F m-2. The latter value has previously been determined by Boily and Fein23,74 to be appropriate for use with corundum particles dispersed in 0.01 M NaNO3 solutions. It is also within the range of Stern layer capacitances used by Hiemstra et al.64 in their MUSIC modeling of aluminum hydroxide materials. A summary of the parameters used in CD-MUSIC model simulations of oxalate adsorption on corundum is given in Table 1. Optimal CD-MUSIC model fits to adsorption data for oxalate on corundum are shown in Figure 7 and were obtained using a single inner-sphere complex according to the surface complexation reaction

tAlOH0.5- + H+ + 0.5Ox2- T tAlOx0.50.5log K ) 10.22 (10) Stoichiometrically, eq 10 indicates that the adsorption of one oxalate anion involves an inner-sphere (ligand exchange) reaction with two surface hydroxyl groups. To maintain consistency with the ATR-FTIR results presented in section 3.1, we assume that the two surface hydroxyls involved in the ligand exchange reaction with each oxalate anion are bound to the same surface aluminum group (with each interacting with a different oxalate carboxyl group), such that the adsorbed oxalate complex is mononuclear in nature. However, the stoichiometry of eq 10 is equivalent for mononuclear and binuclear oxalate complexes, and so an unequivocal distinction between these complexes is reliant upon spectroscopic evidence such as that presented in section 3.1 and elsewhere.40,46 (73) Bargar, J. R.; Fuller, C. C.; Davis, J. A. Abstracts of Papers, 223rd ACS National Meeting, Orlando, FL, April 7-11, 2002, GEOC120. (74) Boily, J. F.; Fein, J. B. Chem. Geol. 2000, 168, 239.

The optimal CD-MUSIC model fit shown in Figure 7 was obtained by distributing the charge of the adsorbed oxalate complex between the 0- and β-planes such that ∆z0 ) 0.4 and ∆zβ ) -0.4, where ∆z0 and ∆zβ are the charge changes in the 0- and β-planes upon ligand exchange of a singly coordinated surface hydroxyl for an oxalate carboxyl group at the corundum surface (see eq 10). Such values of ∆z0 and ∆zβ imply that 40% of the charge of the inner-spherically adsorbed oxalate anion is located in the β-plane. The reasonableness of this value can be assessed using the relationship

R)z+

∑sici

(11)

where R and z are the residual charge and the valence () -2) of a given oxalate oxygen atom, respectively, and si and ci are the bond valence and covalency of each bond i to a given oxalate oxygen atom. Equation 11 can be approximately solved for the surface bonding and nonbonding (solution-facing) oxygens in the adsorbed oxalate anion by utilizing a combination of the bond valences calculated by Brown75 for oxalic acid dihydrate, a knowledge of the bond valence contributions from surface Al(III) cations, and the electronegativitybased ionic character tables of Pauling.76 For the surface bonding oxygen in each oxalate carboxyl group, the bond valence of the C-O bond is 1.34 valence units (vu),75 with the electronegativity difference of 0.89-1.07 between the C and O atoms77 yielding a bond with 75-82% covalent character. By definition, this oxalate oxygen also forms a bond with a surface Al(III) with a bond valence of either 0.50 vu (for 6-coordinated aluminum) or 0.60 vu (for 5-coordinated aluminum40) and with 36-44% covalent character (based on an electronegativity difference of 1.83-2.03 between Al(III) and O77). From eq 11, the surface bonding oxygen in each oxalate carboxyl group therefore possesses a residual charge in the range of -0.64 to -0.82. By contrast, for the nonbonding (solution-facing) oxygen in each oxalate carboxyl group, the bond valence of the CdO bond is 1.66.75 Unfortunately, the ionic character relationship of Pauling76 is applicable only to single bonds. (75) Brown, I. D. The Chemical Bond in Inorganic Chemistry: The Bond Valence Model; Oxford University Press: Oxford, 2002. (76) Pauling, L. The Nature of the Chemical Bond, 3rd ed.; Cornell University Press: Ithaca, NY, 1960. (77) Huheey, J. E.; Keiter, E. A.; Keiter, E. A. Inorganic Chemistry: Principles of Structure and Reactivity, 4th ed.; HarperCollins College Publishers: New York, 1993.

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However, the covalent character of the CdO bond is expected to be greater than that of the C-O bond (which is 75-82% covalent), and we therefore define the possible covalency of the CdO bond as 75-100%. Disregarding solvation effects for simplicity, the resulting residual charge on each nonbonding oxygen is in the range of -0.34 to -0.76. A similar treatment yields a possible residual charge on the central C of each carboxyl group in the range of 0.24-0.75. Assuming both that the residual charges on the surface bonding and nonbonding oxygens reside in the 0- and β-planes, respectively, and that the residual charge on the central C atom is distributed evenly between the 0- and β-planes, then 30-46% of the charge of each adsorbed oxalate carboxyl group is predicted to reside in the β-plane. While obviously only a rough approximation, this simple bond valence treatment does provide a rationale for the significant proportion of the charge of the inner-spherically adsorbed oxalate complex that is required to reside in the β-plane in order to obtain reasonable CD-MUSIC model fits to oxalate adsorption data. The optimal CD-MUSIC model fit also yielded tAlOH0.5- and tAl3O0.5- concentrations of 6.7 and 1.7 sites/ nm2, respectively, with the total site density of 8.4 sites/ nm2 (excluding Al2OH0 groups) lying within the range of previous crystallographic estimates78 and also within the range of site densities considered in other surface complexation studies of adsorption on corundum.23,31,79 The relative concentrations of tAlOH0.5- and tAl3O0.5- give a point of zero charge for the corundum surface of 9.7, in reasonable agreement with the experimental isoelectric point of 9.4 obtained by Johnson and co-workers for AKP30 corundum in 0.01 M KNO3.24,25,63 At the two lowest (0.125 and 0.250 mM) and the highest (2.50 mM) oxalate concentrations investigated, the average value of V(Y) ()9.7) obtained for the CD-MUSIC model fits shown in Figure 7 is in accordance with the FITEQL guidelines laid out by Herbelin and Westall70 (where a satisfactory model fit is indicated by values for V(Y) in the range of 0.1-20). At pH g 5, reasonable model fits are also obtained at oxalate concentrations of 0.50 and 1.25 mM (average V(Y) ) 12.4). At lower pH conditions, however, the model fit clearly and systematically overestimates the extent of oxalate adsorption measured at 0.50 and 1.25 mM. The latter phenomenon is believed to be largely attributable to competitive solution complexation of oxalate by dissolved Al(III) and requires consideration of the dissolution properties of corundum in the presence of adsorbed oxalate (see section 3.3). The CD-MUSIC model fits shown in Figure 7 were obtained using an ionic strength of 10 mM and solution protonation constants (for aqueous oxalate species and H2O) that were corrected to that ionic strength using both specific interaction theory80 and the Davies equation.81 Obviously, however, in addition to the 10 mM KNO3 background electrolyte, solution-based (nonadsorbed) oxalate species can also contribute to the overall ionic strength. For example, considering the most extreme case of the highest oxalate concentration examined (2.50 mM) and pH ) 10 (where the extent of oxalate adsorption is very low (see Figure 7) and oxalate is present predomi(78) Koretsky, C. M.; Sverjensky, D. A.; Sahai, N. Am. J. Sci. 1998, 298, 349. (79) Hayes, K. F.; Redden, G.; Ela, W.; Leckie, J. O. J. Colloid Interface Sci. 1991, 142, 448. (80) Pettit, L. D. Ionic strength corrections using specific interaction theory, 1.2 ed.; IUPAC, 2002. (81) Perrin, D. D.; Dempsey, B.; Serjeant, E. P. pKa prediction for organic acids and bases; Chapman and Hall: London, 1981.

Johnson et al.

Figure 8. Concentration of Al(III) dissolved from corundum after a reaction time period of 72 h as a function of oxalate concentration and solution pH. The corundum concentration was 50 g/L and the background electrolyte was 0.01 M KNO3 in all cases. The total oxalate concentration was 0.125 mM, b; 0.250 mM, O; 0.500 mM, 2; 1.25 mM, 4; and 2.50 mM, (.

nantly as Ox2-), a total ionic strength (including the 10 mM KNO3 background electrolyte and K+ arising from suspension pH adjustment) of approximately 18 mM will result. To examine the significance of this increased ionic strength on the model fits shown in Figure 7, CD-MUSIC simulations were repeated using an ionic strength of 18 mM. In undertaking this exercise, all solution protonation constants were corrected to an ionic strength of 18 mM, while the surface site densities, specific capacitances, surface protonation constants, ion-pair formation constants, and oxalate surface complexation constant were maintained at the same values as in Table 1. The charge distribution was similarly unchanged, with ∆z0 ) 0.4 and ∆zβ ) -0.4. The resulting goodness of fit of the CD-MUSIC model simulations to the adsorption data of Figure 7 was found to be very similar to that originally obtained at an ionic strength of 10 mM. For example, at the two lowest (0.125 and 0.250 mM) and the highest (2.50 mM) oxalate concentrations modeled using an ionic strength of 18 mM, an average V(Y) of 9.6 was obtained (compared with V(Y) ) 9.7 at an ionic strength of 10 mM). Similarly, at pH g 5 and oxalate concentrations of 0.50 and 1.25 mM, an average V(Y) of 12.2 was obtained for an ionic strength of 18 mM (compared with V(Y) ) 12.4 at an ionic strength of 10 mM). Such findings clearly indicate that ionic strength contributions from nonadsorbed oxalate do not significantly affect the CD-MUSIC model simulations undertaken in this study. 3.3. Dissolution of Corundum by Adsorbed Oxalate. The dissolution properties of corundum are shown as a function of both oxalate concentration and pH in Figure 8. The background electrolyte concentration was 0.01 M KNO3 and the reaction time was 72 h in all cases. At the lowest oxalate concentration examined (0.125 mM), the concentrations of dissolved Al(III) in solution are very similar to those previously reported as a function of pH for corundum in the absence of LMW organic adsorbates.31 The very low concentrations of dissolved Al(III) found in the intermediate pH range of 5-9 are characteristic of aluminum-(oxy)hydroxide materials and are consistent with the slow rate of dissolution29 and/or the low solubility of corundum43 in this pH domain. The significantly increased concentrations of Al(III) found under both more acidic (pH < 5) and basic (pH > 9) conditions are consistent with an increased attack on the corundum surface by dissolution-enhancing proton and hydroxyl species, respectively, and are also typical of aluminum-(oxy)hydroxide minerals.41-43 Unlike other aluminum-(oxy)hydroxides,


however, the dissolution rate of corundum displays a marked insensitivity to pH at pH values less than 4,29,31 as is shown by the almost constant concentrations of dissolved Al(III) found for the lowest oxalate concentrations examined (see Figure 8). This phenomenon is believed to be due to the slow regeneration rate of dissolution-active sites at low pH, which is, in turn, driven by the low rate of water exchange at surface Al(III) sites under acidic conditions.30 The concentrations of dissolved Al(III) in solution are similar over the entire pH range examined for the two lowest oxalate concentrations (0.125 and 0.250 mM) shown in Figure 8. Combined with the close similarity of these data with Al(III) concentrations found as a function of pH in the absence of oxalate,31 this finding indicates that at low surface coverages, the adsorption of oxalate does not have a significant effect on the dissolution of corundum. Axe and Persson46 have similarly reported that a low surface coverage (Γ ) 0.57 µmol/m2) of oxalate on boehmite (γ-AlOOH) did not lead to any detectable enhancement of the dissolution of boehmite. Such findings suggest that a minimum, critical concentration of adsorbed oxalate is required to initiate oxalate-induced dissolution of aluminum-(oxy)hydroxide minerals. At the three highest oxalate concentrations examined (0.50, 1.25, and 2.50 mM), Figure 8 shows that the concentration of dissolved Al(III) in solution gradually increases compared with lower oxalate concentrations under acidic conditions. This effect is most marked at low pH at an oxalate concentration of 0.50 mM but leads to significantly enhanced concentrations of dissolved Al(III) being detected under all acidic pH conditions at [oxalate] ) 1.25 and 2.50 mM. Three mechanisms are likely to be responsible for the oxalate-induced enhancement of corundum dissolution. (1) The adsorption of oxalate in an inner-sphere, mononuclear, bidentate fashion polarizes bridging Al-O bonds in the corundum surface, thus weakening them and leaving them susceptible to detachment from the surface. This mechanism of oxalatepromoted dissolution of (oxy)hydroxide minerals has been discussed in detail in a number of previous studies.1-4 (2) The adsorption of oxalate reduces the magnitude of (and can even reverse) the positive charge at the corundumwater interface under acidic pH conditions.24 As a result, the concentration of H+ cations in the vicinity of the interface is expected to rise as a function of increasing oxalate concentration, increasing the rate of protolytic attack on the corundum surface. An enhanced protonpromoted rate of dissolution will result. (3) The presence of oxalate will increase the rate of regeneration of dissolution-active sites under acidic pH conditions. Samson et al.30 have recently shown that the slow rate of dissolution-active site regeneration on corundum under acidic pH conditions can be correlated with the low water exchange rate of the aqueous Al3+ species. Strongly coordinating electron-donating ligands such as F- and OH- dramatically enhance the rate of water exchange around Al3+.82,83 The presence of oxalate similarly enhances the water exchange rate around Al3+ 84 and, by (82) Phillips, B. L.; Casey, W. H.; Crawford, S. N. Geochim. Cosmochim. Acta 1997, 61, 3041. (83) Phillips, B. L.; Tossell, J. A.; Casey, W. H. Environ. Sci. Technol. 1998, 32, 2865. (84) Phillips, B. L.; Crawford, S. N.; Casey, W. H. Geochim. Cosmochim. Acta 1997, 61, 4965. (85) Smith, R. M.; Martell, A. E. Critical Stability Constants, Vol. 6; Plenum Press: New York, 1989. (86) Nordstrom, D. K.; May, H. M. Aqueous Equilibrium Data for Mononuclear Aluminum Species. In The Environmental Chemistry of Aluminum; Sposito, G., Ed.; CRC Press: Boca Raton, FL, 1996.

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analogy, is therefore expected to increase the rate of dissolution-active site regeneration at the corundum surface. The data obtained at pH < 3.5 in Figure 8 provide support for this hypothesis in showing that the extent of corundum dissolution becomes increasingly pH-dependent as the oxalate concentration is raised (i.e., the rate-limiting effect of dissolution-active site regeneration is diminished as a function of increasing oxalate concentration). Oxalate strongly complexes dissolved Al(III) in solution under acidic pH conditions. This is illustrated in Figure 9a-e, in which the speciation of dissolved Al(III) is shown at total oxalate concentrations of 0.125, 0.250, 0.500, 1.25, and 2.50 mM, respectively. The concentrations of dissolved Al(III) are taken from Figure 8, while the concentrations of solution-based (nonadsorbed) oxalate are derived from Figure 7. At total oxalate concentrations of 0.125 and 0.250 mM, no oxalate is present in solution at pH < 7, and Figure 9a,b shows that the dissolved Al(III) is present predominantly in the form of Al3+ and AlOH2+. By contrast, at a total oxalate concentration of 0.500 mM, a small concentration of oxalate is present at pH < 5 and interacts with dissolved Al(III) to form AlOx+. A significant concentration of free Al3+ is also present in solution under these pH conditions (see Figure 9c). At a total oxalate concentration of 1.25 mM, free (nonadsorbed) oxalate is present in still larger concentrations over the entire range of pH conditions investigated. At pH < 4.5, oxalate is fully complexed by dissolved Al(III), forming AlOx+ and AlOx2- species (see Figure 9d). Significant concentrations of free Al3+ are also present under such low pH conditions. By contrast, at pH g 4.5, some uncomplexed Ox2- is also present, as is the trinuclear species AlOx33-. Above pH ) 7, however, no Al(III)-oxalate species are formed in significant concentrations, and oxalate is present predominantly as the Ox2- anion. At the highest oxalate concentration examined (2.50 mM), significant quantities of nonadsorbed oxalate and dissolved Al(III) are present over the range of acidic pH conditions investigated (see Figure 9e). In this case, a number of different Al(III)-oxalate complexes again exist, including AlOx+ (at pH e 3.5), AlOx2- (at pH e 7.0), and AlOx33- (between pH ) 2.5 and pH ) 7.5). Interestingly, unlike at lower total oxalate concentrations, uncomplexed oxalate (in the form of HOx- or Ox2-) is also present across the entire range of pH conditions examined, and a significant concentration of Al3+ is found only at the lowest pH investigated (pH ) 2.0). The speciation data of Figure 9a-e provide a likely explanation for the seemingly anomalous adsorption results obtained at oxalate concentrations of 0.500 and 1.25 mM and pH values less than 5 (see Figure 7). At the lowest oxalate concentrations (0.125 and 0.250 mM), oxalate is strongly adsorbed to the corundum surface and has little effect on the corundum dissolution behavior. By contrast, at intermediate oxalate concentrations of 0.500 and 1.25 mM, adsorbed oxalate promotes the dissolution of corundum. Upon extracting Al(III) cations from the corundum surface at pH < 5, desorbed oxalate will rapidly form one or more Al(III)-oxalate complexes that will be either be electrostatically repelled from the positively charged surface (in the case of AlOx+) or have only a weak electrostatic attraction for the surface (in the case of AlOx2-). Such solution Al(III)-oxalate complexes are therefore anticipated to interact weakly (if at all) with the corundum surface, an expectation that has been borne out by the recent ATR-FTIR studies of Axe and Persson,46 which demonstrated that AlOx+ does not adsorb to (87) Sjoberg, S.; Ohman, L. O. J. Chem. Soc., Dalton Trans. 1985, 2665.

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Figure 9. Speciation of dissolved Al(III) in solution based on experimental solution concentrations of Al(III) (from Figure 8) and oxalate (derived from Figure 7). The total oxalate concentration (including adsorbed oxalate; see Figure 7) is (a) 0.125 mM, (b) 0.250 mM, (c) 0.500 mM, (d) 1.25 mM, and (e) 2.50 mM. Formation constants for Al-oxalate complexes were obtained were obtained by averaging values reported by Nordstrom and May (ref 86) and Sjoberg and Ohman (ref 87) after correction to I ) 0.01 M using both specific interaction theory (ref 80) and the Davies equation (ref 81). Al(III) hydrolysis constants were obtained from Nordstrom and May (ref 86). Oxalate protonation constants were obtained from Smith and Martell (ref 85). All hydrolysis and protonation constants were corrected to I ) 0.01 M using the Davies equation (ref 81). Major species present in solution were Al3+, b; Ox2-, O; HOx-, 2; AlOx+, 4; AlOx2-, [; AlOx33-, ]; and AlOH2+, 3. Other species (e.g., AlHOx2+, Al3OH3Ox3, Al2OH2Ox4,4- AlOH2+, AlOH3, AlOH4-) were also considered in all speciation calculations but were not found to be present in significant concentrations over the range of pH conditions examined.

boehmite (γ-AlOOH) to a significant extent (at least at pH ) 5.1). As a result, dissolved Al(III) effectively competes with the corundum surface for oxalate at pH < 5 through formation of Al(III)-oxalate solution complexes with a lower affinity for the corundum surface than the uncomplexed Ox2- anion. A similar competition between dissolved Al(III) and aluminum-(oxy)hydroxides has previously been proposed by Fein and Brady.17 At the highest oxalate concentration examined (2.50 mM), the presence of uncomplexed Ox2- and HOx- in solution and the absence of dissolved Al3+ at all but the lowest pH investigated indicate that sufficient oxalate exists to both fully complex the corundum surface and satisfy the complexation requirements of dissolved Al(III). As a result, no competition for oxalate between the corundum surface and dissolved Al(III) is observed, and the adsorption of oxalate

progressively increases as a function of decreasing pH over the entire pH range investigated (as predicted by CD-MUSIC modeling; see Figure 7). At a total oxalate concentration of 1.25 mM, Figure 9d shows the presence of a significant concentration of AlOx33- between pH ) 4.5 and 7. Similarly, at a total oxalate concentration of 2.50 mM, Figure 9e shows the presence of the AlOx33- species at pH ) 2.5-7.5. Such species may be expected to possess a strong electrostatic attraction to the positively charged corundum surface and may therefore be anticipated to enhance the apparent adsorption of oxalate in those pH domains. However, past electrokinetic studies24 performed after a short equilibration time (i.e., where the concentration of dissolved Al(III) is very low) have demonstrated that the effective corundum interfacial potential (due to the charged


Figure 10. Macroscopic adsorption properties of oxalate on corundum as a function of dissolved Al(III) concentration at pH ) 2.5. The total oxalate concentration was 1.25 mM, the corundum concentration was 50 g/L, the background electrolyte was 0.01 M KNO3, and the reaction time was 72 h in all cases. Dissolved Al(III) was added in the form of Al(NO3)3, although a small concentration of Al(III) also resulted from dissolution of the corundum surface. Actual total Al(III) concentrations were quantified using ICP spectrometry.

corundum surface and any tightly bound ionic species, including Ox2-) is negative above pH ≈ 5 at [oxalate] ) 1.25 mM and above pH ≈ 3 at [oxalate] ) 2.50 mM. As a result, an electrostatic repulsion is expected to exist between the interface and AlOx33- under those pH conditions, and so the adsorption of AlOx33- is not anticipated occur to a significant extent. The predicted effects of dissolved Al(III) in decreasing the adsorption of oxalate on corundum at intermediate oxalate concentrations are confirmed more directly in Figure 10, in which the extent of oxalate adsorption is shown as a function of solution Al(III) concentration at pH ) 2.5. In all cases, the total oxalate concentration was 1.25 mM, the reaction time was 72 h, and solution-based Al(III) was added in the form of Al(NO3)3. Figure 10 demonstrates that as the concentration of dissolved Al(III) increases, the extent of oxalate adsorption is substantially reduced. The extent of the adsorption decrease is most significant at the lowest concentrations of dissolved Al(III) but continues to be observed up to the highest (20.0 mM) Al(III) concentration examined. Under the latter condition, the extent of oxalate adsorption is only 41% of that observed at the lowest Al(III) concentration investigated ()0.5 mM). Such results clearly demonstrate that dissolved Al(III) in solution will significantly reduce the adsorption of oxalate on corundum under acidic conditions. 3.4. Time-Dependent Studies of Oxalate Adsorption and Oxalate-Promoted Dissolution of Corundum. The adsorption of oxalate on corundum at a total oxalate concentration of 1.25 mM is shown as a function of reaction time and pH in Figure 11. The background electrolyte concentration was 0.01 M KNO3 in all cases. The corresponding corundum dissolution behavior over the same range of experimental parameters is shown in Figure 12. At pH > 5, Figure 11 shows that the extent of oxalate adsorption is almost identical at the two longest reaction times examined (72 and 336 h), indicating that an equilibrium (or pseudoequilibrium) adsorption condition has been attained. By contrast, the oxalate adsorption data obtained at the shortest reaction time (3 h) at pH > 5 consistently fall below those obtained at longer times, indicating that an equilibrium (or pseudoequilibrium) condition has not yet been attained. Figure 12 shows that at pH > 5, the concentrations of dissolved Al(III) present after 72 and 336 h are very

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Figure 11. Macroscopic adsorption properties of oxalate on corundum as a function of equilibration time and pH. The total oxalate concentration was 1.25 mM, the corundum concentration was 50 g/L, and the background electrolyte was 0.01 M KNO3 in all cases. The reaction time was 3 h, b; 72 h, O; and 336 h, 2.

Figure 12. Concentration of Al(III) dissolved from corundum at a total oxalate concentration of 1.25 mM as a function of equilibration time and solution pH. The corundum concentration was 50 g/L and the background electrolyte was 0.01 M KNO3 in all cases. The reaction time was 3 h, b; 72 h, O; and 336 h, 2.

similar, indicating that any effect of dissolved Al(III) on the corresponding adsorption data of Figure 11 should be similar for these two systems. At pH < 5, however, the concentration of dissolved Al(III) is observed to be significantly higher after 336 h than after 72 h, due to the longer time over which oxalate-induced dissolution of the corundum surface has occurred. This effect is manifested in Figure 11, where at pH < 5, the concentration of oxalate adsorbed after 336 h is lower than that adsorbed after 72 h, consistent with the greater competition for oxalate by dissolved Al(III) at the longest equilibration time. Interestingly, Figure 11 also shows that at pH < 3.5 and the shortest equilibration time of 3 h, the extent of oxalate adsorption on corundum is higher than that obtained after either 72 or 336 h. The latter result indicates that despite the short time allowed for oxalate adsorption to take place, the much less significant competition for oxalate by Al(III) (due to the substantially lower dissolved Al(III) concentrations obtained after 3 h; see Figure 12) leads to a higher extent of oxalate adsorption being attained than at longer reaction times. 4. Conclusions The adsorption of a simple low molecular weight organic anion, oxalate, on a model aluminum oxide colloid, corundum (R-Al2O3), and its effects on the dissolution of corundum have been investigated over a range of different

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oxalate concentrations (0.125-25.0 mM, corresponding to surface concentrations, Γ, of 0.4-71 µmol/m2 when fully adsorbed) and pH conditions (2-10). At low to intermediate oxalate concentrations (Γ e 7.1 µmol/m2), in situ ATRFTIR spectroscopic measurements indicate that oxalate is adsorbed to corundum predominantly as a bidentate, mononuclear, inner-sphere complex involving both deprotonated carboxyl groups. Some very minor contributions from outer-spherically adsorbed oxalate are also observed at Γ g 3.6 µmol/m2, with these becoming more significant at higher concentrations (Γ g 14 µmol/m2). Significant contributions from nonadsorbed Ox2- are also evident at high total oxalate concentrations. In accordance with the ATR-FTIR data, macroscopic data for oxalate adsorption on corundum can be generally well fit using the CD-MUSIC model and a single bidentate, inner-sphere complex over a broad range of oxalate concentrations (0.125-2.50 mM). At pH < 5, however, the concentration of adsorbed oxalate measured at intermediate (0.500 and 1.25 mM) oxalate concentrations is found to fall significantly below CD-MUSIC model predictions. The latter behavior is due to competition for oxalate from solution-based Al(III), the formation of which is driven by oxalate-promoted dissolution of the corundum

Johnson et al.

surface. The formation of dissolved Al(III)-oxalate complexes such as AlOx+, AlOx2-, and AlOx33- can result, with these species having a lower affinity for the corundum surface than uncomplexed Ox2-. The addition of dissolved Al(III) to corundum-oxalate suspensions at pH ) 2.5 confirms the competitive nature of Al(III), with the extent of oxalate adsorption decreasing as a function of rising solution-based Al(III) concentration. At pH < 3.5, the extent of oxalate adsorption is also found to decrease as a function of increasing reaction time. The latter behavior can be attributed to the increased concentrations of dissolved Al(III) generated at the longest reaction times examined. Acknowledgment. The authors thank Dr. Scott Fendorf for the use of FTIR, ion chromatography, and inductively coupled plasma spectrometry equipment in his laboratory and Ben Kocar and Guangchao Li for their help in performing IC and ICP measurements, respectively. This work was supported by EPA-STAR Grant EPAR827634-01-1 (G.E.B.) and NSF Grant CHE-0089215 (G.E.B.). LA048559Q

Water Interfaces - American

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