HOME HELP FEEDBACK SUBSCRIPTIONS ARCHIVE SEARCH TABLE OF CONTENTS		QUICK SEARCH:   [advanced] Author: Keyword(s): Year:  Vol:  Page:

This Article Abstract Figures Only Full Text (PDF) Free Alert me when this article is cited Alert me if a correction is posted Services Similar articles in this journal Similar articles in ISI Web of Science Alert me to new issues of the journal Download to citation manager Citing Articles Citing Articles via ISI Web of Science (1) Citing Articles via Google Scholar Google Scholar Articles by Luxton, T. P. Articles by Eick, M. J. Search for Related Content PubMed Articles by Luxton, T. P. Articles by Eick, M. J. GeoRef GeoRef Citation Agricola Articles by Luxton, T. P. Articles by Eick, M. J. Related Collections Toxic Trace Metals Heavy Metals Soil Chemistry

Soil Chemistry

Mobilization of Arsenite by Competitive Interaction with Silicic Acid

^a Dep. of Crop and Soil Environmental Sci., Virginia Tech, Blacksburg, VA 24061
^b Dep. of Geoscience, Virginia Tech, Blacksburg, VA

^* Corresponding author (tluxton{at}vt.edu)

	ABSTRACT

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION REFERENCES

 
Due to the acute toxicity of As, mobilization of even a small fraction of As into surface and ground waters used as a source for drinking water represents a substantial risk to human health. Here we evaluate the mobilization of arsenite [As(III)] from the Fe-(hydr)oxide mineral goethite (-FeOOH) through competitive displacement by silicic acid, a naturally occurring and ubiquitous inorganic ligand. The adsorption behaviors of silicic acid and As(III) on goethite were investigated at environmentally relevant pH (3–11). Single ion adsorption and zeta-potential data were collected at silica concentrations characteristic of natural waters (3–30 mg L^–1) and initial solution As(III) concentrations representative of high levels of contamination (3.75–7.5 mg L^–1). Competitive adsorption scenarios with either Si or As(III) sorbed first to the goethite surface, followed by equilibration with the other sorbate, were also examined. No competitive displacement of either oxyanion was observed at total sorbate concentrations less than reactive surface site density, regardless of pH or addition scenario. However, at total sorbate concentrations greater than reactive surface site density, As(III) adsorption was reduced by 10 to 15% over the entire pH range regardless of addition scenario, resulting in aqueous concentrations well in excess of current (10 µg L^–1) drinking water maximum contaminant levels. Surface complexation modeling of single ion adsorption and zeta-potential data using the Charge Distribution Multisite Surface Complexation (CD-MUSIC) model was used to calculate an appropriate set of surface adsorption equilibrium constants for As(III) and silicic acid adsorption, which was used to describe the competitive adsorption scenarios. Comparison of competitive adsorption data and CD-MUSIC model predictions, at total sorbate concentration greater than reactive surface site density of goethite, suggest that silica is competitively displaced by As(III).

Abbreviations: AO, ammonium oxalate • As(V), arsenate • As(III), arsenite • CD MUSIC, Charge Distribution Multisite Surface Complexation • CBD, citrate–bicarbonate–dithionite • DOM, dissolved organic matter • ICP–AES, inductively coupled plasma atomic emission spectrometer • IEP, isoelectric point • MCL, maximum contaminant load • XPS, x-ray photoelectron spectroscopy

	INTRODUCTION

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION REFERENCES

 
ARSENIC is a ubiquitous trace element in the environment, found in the atmosphere, soils and rocks, natural waters, and biomass. Arsenic is both a toxin and a carcinogen, implicated in cardiovascular, pulmonary, immunological, neurological, and endocrine disorders, as well as skin, lung, bladder, and kidney cancers (NRC, 1999, 2001; USEPA, 2001). Long-term exposure to contaminated drinking water has been cited as the most widespread threat to human health posed by As although occupational exposure to airborne As also represents a significant if less widespread health risk (Nordstrom, 2002; Smedley et al., 2002; Smith et al., 2002a). Arsenic contamination of natural waters used as drinking water sources has been well documented worldwide, particularly in areas that rely heavily on ground water as a drinking water supply (Welch et al., 2000; Nordstrom, 2002; Smedley et al., 2002). Sources of As in natural waters include weathering and dissolution of As bearing minerals, and geothermal activity; as well as, anthropogenic inputs such as mineral extraction and processing wastes, coal fly-ash, arsenical pesticides and poultry and swine feed additives, wood preservatives, and glass-making byproducts. However, As contamination of drinking water appears to be most commonly associated with release from (hydr)oxide mineral surfaces (Welch et al., 2000; Stollenwerk, 2003).

The mobility, bioavailability, and toxicity of As are intimately related to its oxidation state. Arsenic most commonly occurs in soils and natural waters as a weakly acidic oxyanion in the form of arsenate [As(V); H₃AsO₄] in oxidized environments or As(III) (H₃AsO₃) under reducing conditions. However, research has shown that As(III) may persist for extended periods of time in oxic environments due to slow kinetics and biological transformations (Inskeep et al., 2002). Both As(V) and As(III) sorb strongly to common soil and aquifer (hydr)oxide minerals, particularly Fe-(hydr)oxides due to their nearly ubiquitous distribution and high abundance, forming inner-sphere surface complexes through ligand exchange reactions with hydroxyl functional groups (Waychunas et al., 1993; Manning et al., 1998). While both As redox species form inner-sphere complexes with the goethite  surface they do exhibit very different adsorption behavior based on the chemical properties of each species. Arsenate adsorption is characterized by maximum adsorption occurring near the first pKa 2.2, while As(III) adsorption is characterized by maximum adsorption occurring near pH 9, the first pKa 9.2 (Manning and Goldberg, 1996; Manning et al., 1998; Grafe et al., 2001; Dixit and Hering, 2003). The solubility of As associated with Fe-(hydr)oxide minerals is controlled through adsorption/desorption and coprecipitation/dissolution reactions at the mineral–water interface that are sensitive to changes or fluctuations in pH and E_H (Sun and Doner, 1996; Manning et al., 1998; Goldberg and Johnston, 2001; Inskeep et al., 2002; Smedley and Kinniburgh, 2002; Zhang et al., 2004). In abiotic sorption experiments As(III) has been shown to have a higher affinity for sorption onto ferrihydrite and goethite surfaces as compared with As(V) in the circumneutral pH range (Dixit and Hering, 2003). However, As(III) has consistently been reported and generally accepted to be more mobile than As(V) in native soils and sediments (Frankenberger, 2002). One proposed mechanism for the mobilization of As as As(III) from Fe-bearing minerals has been attributed to reductive dissolution of oxide surfaces although the detailed mechanisms of As reduction and release remain unclear. Regardless of the mechanism of mobilization, As(III) is considerably more toxic than As(V) (Winship, 1984; Korte and Fernando, 1991).

Previous spectroscopic research has indicated that As(III) binds with (hydr)oxide surfaces through two adsorption mechanisms. The first is a ligand exchange reaction where As(III) forms an inner-sphere surface complex resulting in either a mono- or bidentate surface complex (Sun and Doner, 1996; Manning et al., 1998; Goldberg and Johnston, 2001). In the second, As(III) interacts with the (hydr)oxide surface electrostatically forming an outer-sphere complex (Arai et al., 2001; Goldberg and Johnston, 2001). The presence of an As(III) outer-sphere complex has only been confirmed for As(III) adsorbing onto -Al₂O₃, and amorphous aluminum oxide and then both inner and outer-sphere complexes existed (Arai et al., 2001; Goldberg and Johnston, 2001). However, Goldberg and Johnston (2001) proposed the existence of As(III) outer-sphere complexes from FTIR data for As(III) adsorption onto ferrihydrite.

Competition for sorption sites on Fe-(hydr)oxide minerals by naturally occurring ligands including phosphate, sulfate, carbonate, silicate, and dissolved organic matter (DOM) may also act as a mechanism for As mobilization. Ligands competing for surface sorption sites will influence the stability of adsorbed As and consequently its potential bioavailability (Aggett and Roberts, 1986; Sadiq, 1997; Swedlund and Webster, 1999; Meng et al., 2000; Nickson et al., 2000; Davis et al., 2001; Grafe et al., 2001; Smith et al., 2002b; Waltham and Eick, 2002). Phosphate has a very high affinity for Fe-(hydr)oxides, but because of continuous biological uptake is typically present in unammended soils and natural waters at extremely low concentrations (Geelhoed et al., 1997; Tadanier et al., 2002). Previous research has shown that sulfate, carbonate, and DOM are relatively ineffective competitive ligands for As adsorbed on Fe-(hydr)oxides (Smith et al., 1999; Jain and Loeppert, 2000; Meng et al., 2000; Grafe et al., 2001; Clifford and Ghurye, 2002; Smith et al., 2002b; Zhang et al., 2004). However, under conditions when As solution concentrations exceed Fe (hydr)oxide reactive surface site density, carbonate has been shown to displace adsorbed As (Appelo et al., 2002; Anawar et al., 2004).

Silicic acid, like As, has a high affinity for Fe-(hydr)oxide surfaces, also forming inner-sphere surface complexes through ligand exchange with hydroxyl functional groups (Sigg and Stumm, 1981; Vempati and Loeppert, 1989; Vempati et al., 1990; Hansen et al., 1994b; Doelsch et al., 2001, 2003). Because its high affinity for surface adsorption and ubiquitous occurrence in natural waters and soil solutions, commonly present at 5 to 35 mg L^–1 (0.17–1.24 mM) and not infrequently at concentrations as high as 75 mg L^–1 (2.7 mM) in ground water (Iller, 1979), dissolved silica monomers (H₄SiO₄⁰ and H₃SiO₄^–) represent potentially effective competitive ligand for As sorbed onto Fe-(hydr)oxide surfaces. Furthermore, formation of silica polymers, either in solution, at Si concentrations above 1.1 mM (Alverez and Sparks, 1985; Applin, 1987), or on hydrous oxide surfaces, at substantially lower Si concentrations, may enhance the competitive ligand effect (Vempati et al., 1990; Hansen et al., 1994b; Doelsch et al., 2001, 2003).

The competitive interaction between dissolved silica, both in monomeric and polymeric forms, and As(III) adsorbed on poorly crystalline hydrous Fe(III) oxides has been recently investigated (Swedlund and Webster, 1999; Meng et al., 2000). These studies indicated that silica was able to competitively displace As(III) sorbed on ferrihydrite. However, poorly crystalline hydrous Fe(III)-oxides undergo transformation to more crystalline forms such as goethite, hematite, and magnetite through both abiotic and biotic processes. The kinetics of these transformations depend on temperature, pH, solution composition, and microbial activity (Schwertmann and Cornell, 1996; Ford et al., 1997; Fredrickson et al., 1998; Benner et al., 2002; Ford, 2002). Although the distribution of Fe between amorphous and crystalline forms may vary with fluctuations in solution chemistry, based on Fe content crystalline forms predominate in many natural environments (Roden and Urrutia, 2002). Therefore, the objective of this research is to examine the extent to which silica, as a naturally occurring competitive ligand, may impact the mobility and potential bioavailability of As(III) in environments where sorption on crystalline Fe-(hydr)oxide mineral surfaces is a significant mechanism for contaminant sequestration. Here we evaluate the competitive interaction of Si and As(III) by comparing their adsorption behaviors on goethite individually and in combination at several environmentally relevant concentrations. We also use thermodynamic surface complexation modeling of adsorption data as a tool to suggest potential changes in adsorbed oxyanion surface speciation, to determine the possible presence of polymeric Si surface species on goethite, and to assess the impact of Si polymers on As(III) adsorption.

	MATERIALS AND METHODS

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION REFERENCES

 
Goethite Synthesis and Characterization
Goethite was synthesized by hydrolysis of ferric nitrate [Fe(NO₃)₃] using a method described by Schwertmann and Cornell (1991). All solutions were prepared with distilled doubly deionized water and reagent grade chemicals. Contact with glass surfaces was avoided to prevent silica contamination. The identity, purity, and structure of the goethite sample was confirmed by x-ray diffraction (XRD), differential scanning calorimetry (DSC), thermogravimetric analysis (TGA), and field emission scanning electron microscopy (FESEM). Results form these analyses were consistent with data presented by Schwertmann and Cornell (1991, 1996) and a goethite standard from Bayer (Krefeld, FRG). Field emission scanning electron microscopy images indicated the synthetic goethite crystals were euhedral acicular crystals of uniform shape and size approximately 200 nm in length and 25 nm laterally. The quantity of poorly crystalline or amorphous Fe was determined by the ratio of ammonium oxalate (AO) in the dark to citrate–bicarbonate–dithionite (CBD) extractable Fe (Schwertmann and Cornell, 1991; Loeppert and Inskeep, 1996). The extractable Fe was 0.28% (AO Fe mg/CBD Fe mg), indicating minimal amounts of amorphous or short-ordered crystals. Specific surface area was 73 m² g^–1, as determined by a five point N₂ Brunauer–Emmett–Teller (BET) gas adsorption isotherm. The point of zero charge as determined by the isoelectric point (IEP) was 9.56.

Single Ion Adsorption
Arsenite adsorption isotherms were conducted at 0.05 and 0.10 mM utilizing a pH monitored stirred-batch reactor technique. Adsorption studies were performed in 0.01 M NaNO₃ background electrolyte with a goethite suspension density of 1 g L^–1. All adsorption edges were conducted in duplicate from pH 3 to 11. The pH of the suspension was maintained using a Brinkman 716 Stat-Trino pH stat (Brinkman Instruments, Westbury, NY). Goethite suspensions were placed in Teflon lined reaction vessels, stirred at 300 rpm, kept at a constant temperature (25°C) using a temperature controlled water circulation, and continuously sparged with N₂ to eliminate atmospheric CO₂ interference. A 4-h equilibration time for As(III) was chosen based on previous work (Grafe et al., 2001; Raven et al., 1998; Waltham and Eick, 2002) After the 4-h equilibration time for As(III) adsorption, a 17.00-mL sample was removed from the reaction vessel. Seven milliliters was immediately filtered through a 0.2-µm Fisherbrand (Fisher, Atlanta, GA) membrane for As analysis. Arsenic solution concentrations were analyzed by inductively coupled plasma atomic emission spectrometer (ICP–AES) (SpectroFlame FTMOA85D, Spectro Analytical Instruments, Fitchburg, MA) at the 188.979-nm line. The detection limit for As was 16 µg L^–1 and a limit of quantification of 22 µg L^–1. The remaining 10-mL sample was used immediately for zeta potential determination. After each sampling event, the pH in the reaction vessel was raised by one unit using 0.1 M NaOH and allowed to re-equilibrate for 4 h.

Silicic acid adsorption isotherms were conducted at 0.1, 0.5, and 1.00 mM using a batch technique in individual 250 mL polycarbonate bottles. Adsorption studies were performed in 0.01 M NaNO₃ background electrolyte with a goethite suspension density of 1 g L^–1. All adsorption edges were conducted in duplicate from pH 3 to 11. Individual batch experiments were chosen over the pH monitored stirred batch technique for silicic acid adsorption due to the extended equilibration time period of 40 h (Waltham and Eick, 2002). Before sealing, each bottle was adjusted to the appropriate pH value followed by 5 min of sparging with N₂. Bottles were then sealed and placed in a New Brunswick Scientific (Edison, NJ) Innova 4230 Refrigerated Incubator Shaker at 300 rpm and 25°C. The pH was checked and readjusted three times daily during equilibration and each sample sparged with N₂ for 5 min before resealing. Fluctuations in the pH were minimal (<0.1) after 30 h. After the 40-h equilibration time period the pH was measured and a 17.00-mL sample was removed from the reaction vessel and split into a 7-mL sample for ICP–AES analysis at the 251.611-nm line. The detection limit was 10 µg L^–1 with a limit of quantification of 29.2 µg L^–1. A 10-mL sample was immediately analyzed for zeta potential measurement.

Dual Ion Adsorption
The dual ion adsorption of silicic acid and As(III) on goethite was investigated over the pH range 3 to 11 for two silicic acid and As(III) concentrations [0.10 and 1.00 mM Si; 0.05 and 0.10 mM As(III)]. Both addition scenarios, (i) initial equilibration of Si with goethite followed by the addition of As(III), and (ii) initial equilibration of As(III) with goethite followed by Si addition, were examined. Adsorption studies were conducted using the batch technique described for the Si adsorption studies, with the following modifications. The initial oxyanion was added to the goethite suspension to achieve the appropriate solution concentration [0.10 and 1.00 mM for Si or 0.05 and 0.10 mM for As(III)], the suspension pH was adjusted and then sparged with N₂ for 5 min. Suspensions were equilibrated for the appropriate time period [40 h for Si or 4 h for As(III)]. The pH was periodically checked and readjusted throughout the equilibration period (three times daily). As with the Si adsorption experiments, the reaction vessels were sparged with N₂ for 5 min before resealing. Following equilibration of the initial oxyanion, the second oxyanion was added to the system. The suspensions were then allowed to equilibrate for 50 h. The pH of the dual ion samples was measured and adjusted as necessary during the equilibration period. Fluctuations in the pH were minimal (<0.1) after 35 h. The final pH of the suspensions was recorded, and a 17.00-mL sample was removed from the reaction vessel and split into a 7-mL sample for ICP–AES analysis and a 10-mL sample for immediate zeta potential determination.

Arsenic Speciation
The presence of As(V) due to the oxidation of As(III) by Fe(III) was determined by analyzing for the presence of dissolved Fe(II) colormetrically by 1,10-Phenanthroline, and adsorbed As(V) by x-ray photoelectron spectroscopy (XPS) at pH 3. Since the oxidation of As(III) to As(V) is favored at low pH it was assumed if there was no oxidation of As(III) at pH 3 then the As(III) would be stable at higher pH values. For the Fe(II) analysis a goethite suspension was prepared using the same suspension density and experimental parameters as described above in the single ion adsorption section for As(III). The method was modified by extending the equilibration time to 70 h. A 7-mL sample was removed from the reaction vessel and filtered through a 0.22-µm filter at 30 min, 1, 5, 10, 20, 40, 60, and 70 h for Fe(II) analysis. After filtration the samples were acidified to pH < 2 and stored at 5°C to prevent oxidation of Fe(II). The presence of Fe(II) was determined using the 1,10 Phenanthroline method presented by Loeppert and Inskeep (1996) using a Beckman Coulter DU 640 (Fullerton, CA) spectrophotometer with a visible source set at 510 nm.

After the 70-h equilibration time period a 10-mL sample was filtered and the resulting solid material was allowed to dry under a positive pressure N₂ environment. The subsequent goethite was analyzed by XPS to evaluate the redox status of the adsorbed As on the surface. XPS spectra were collected on a PerkinElmer 5400 ESCA system (Wellesley, MA). The x-ray source was a monochromatic Al K radiation at 1486.7 eV Both wide and narrow scans were used, the former to determine the range and abundance of elements present and the latter to determine the chemical state. A small advantageous C(1s) carbon peak (285.0 eV) was monitored throughout each experiment. The peak was invariant throughout each experiment and corrections for charging were not required. Binding energies for the As(3d) arsenic peak were assigned relative to the advantageous C(1s) carbon peak. Narrow scans for the As speciation were collected from 60 to 40 eV for the As(3d) binding energy. The presence of the Fe(3p) iron peak (55.6 eV) did not interfere with the As(3d) arsenic peak. To determine the chemical state of the adsorbed As, reference spectra of As(III) and As(V) reagent grade sodium salts (NaAsO₂ and NaAsO₃) were collected before analysis to determine peak position. Spectra obtained from the As(3d) binding energy region were collected for 1200 cycles of 0.1 eV/step at a 20 msec/step.

Zeta Potential Measurements
Zeta potential measurements were collected over the pH range 3 to 11 for both As(III) and Si single ion adsorption edges at each sorbate ion concentration, and for selected pH values (3, 5, 7, 9, and 11) for the dual ion adsorption edges. Zeta potentials were calculated from microelectrophoresis measurements collected using a Malvern Zetasizer 3000HSa (Malvern Instruments, Southborogh, MA). Based on preliminary data and particle size, the voltage applied to the capillary cell was set at 100 mV and a Henry function [f(K_a)] of 1.5 was used to calculate zeta potential. The pH of each 10-mL sample was measured before zeta potential measurement to account for any drift in pH.

Surface Complexation Modeling
Surface complexation modeling was used as a tool to evaluate the adsorption and zeta potential data collected for Si and As(III) adsorption on goethite. For the current study Si and As(III) were initially assumed to undergo a ligand exchange reaction with the goethite surface to form monomeric inner-sphere surface complexes (Table 1; Reaction 12–19). If the adsorption data could not be described by a ligand exchange mechanism, additional binding mechanisms accounting for either the formation of As(III) outer-sphere complexes, Si solution polymers, or Si Surface polymers were used.

Based on the approach of Hiemstra et al. (1989a, 1996); and Hiemstra and van Riemsdijk (1996) the surface acidity of goethite is controlled by two reactions (Table 1; Reactions 1 and 2). The log K values for both reactions were set equal to the experimentally determined IEP, as determined by electrophoretic mobility measurements (IEP = 9.56), which is a common practice when utilizing the CD-MUSIC model (Hiemstra et al., 1989b; Hiemstra and van Riemsdijk, 1996, 1999, 2000; Geelhoed et al., 1997, 1998; Reitra et al., 1999a, 2000; Tadanier and Eick, 2002). Electrolyte ions were assumed to form outer-sphere ion pairs with surface functional groups. Electrolyte ion pair formation constants have been previously determined by Rietra et al. (2000) and are listed in Table 1 (Reactions 3–6).

Surface site density for singly and triply coordinated surface functional groups for goethite were evaluated using the computational analysis of Barrón and Torrent (1996). Based on the Hiemstra and van Reimsdijk (1996) approach and using a weighted average of functional groups present on the two predominant faces of goethite (Barrón and Torrent, 1996), {110} and {021}, the reactive surface site density was set to 3.52 nm^–2 and 2.7 sites nm^–2 for singly and triply coordinated sites, respectively (Table 2). These values are similar to those used in the literature for CD-MUSIC modeling of goethite (Hiemstra et al., 1996; Hiemstra and van Riemsdijk, 1996, 1999, 2000; Geelhoed et al., 1997, 1998; Reitra et al., 1999a, 1999b, 2000; Tadanier and Eick, 2002).

Zeta potential measurements, used in conjunction with adsorption data, were used to help constrain the 3-plane electrostatic model capacitance factors (C₁ and C₂) and calculate equilibrium constants for single and dual ion adsorption scenarios (Table 2). The electrical potential at the plane of shear was calculated using the electrical potential at the head of the 2-plane provided by the CD-MUSIC model and compared with zeta potential data collected from the electrophoretic mobility measurements. Gouy Chapman Theory (Chapman, 1913; Gouy, 1910) was used to calculate the electrical potential versus distance in the diffuse layer

[1]

where z is the bulk ion charge, e is electron charge (1.602 x 10^–19 C), _x is the electrical potential at a distance x from the 2-plane, k is the Boltzmann constant (1.308 x 10^–23 J K^–1), T is temperature, ₂ is the electric potential at the 2-plane, is the Debye-Huckel screening distance, and x is the distance from the 2-plane. The distance from the 2-plane to the plane of shear was set at 2.6 nm based on modeling results from the current study and previous work modeling phosphorus adsorption on goethite (Tadanier et al., 2002). Additionally, this value is close to the double layer thickness predicted by Gouy Chapman theory (3.07 nm) for a 1:1 electrolyte at 25°C and 0.01 M concentration.

	RESULTS AND DISCUSSION

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION REFERENCES

 
Single Ion Adsorption
Arsenite and silicic acid exhibit contrasting adsorption behaviors on goethite for the solution concentrations investigated. Arsenite adsorption was essentially pH independent, with 98% or more of the As(III) being removed from solution (Fig. 1A ), where as Si adsorption was pH dependent exhibiting an adsorption envelope with maximum adsorption occurring near the first pK_a of Si (9.8) (Fig. 1B). The apparent lack of pH dependence for As(III) adsorption, for the initial solution concentrations investigated, and the strong pH dependence of Si adsorption on goethite have also been reported in the literature (Manning et al., 1998; Swedlund and Webster, 1999; Gustafsson, 2001; Dixit and Hering, 2003). At equimolar concentrations (0.10 mM) the quantity of As(III) adsorbed (1.37 µmol m^–2) on goethite exceeds Si adsorption (1.07 µmol m^–2) by 22% at pH 9.

View larger version (22K): [in this window] [in a new window]   Fig. 1. Oxyanion adsorption on goethite as a function of pH. (A) Arsenite adsorption for two initial solution concentrations (0.05 and 0.10 mM). (B) Silicic acid adsorption as a function of pH for three initial solution concentrations (0.10, 0.50, and 1.00 mM). Symbols represent experimental data, lines represent predicted Si and As(III) adsorption for the optimized Charge Distribution Multisite Surface Complexation (CD MUSIC) model fit to experimental data (Log K values listed in Table 1). Suspension density 1 g L–1, background electrolyte 0.01 M NaNO3.

View larger version (31K): [in this window] [in a new window]   Fig. 2. Zeta potential at the plane of shear, based on electrophoretic mobility measurements, as a function of pH (A) zeta potential for As(III) adsorption at two initial solution concentrations (0.05 and 0.10 mM). (B) zeta potential for Si adsorption at three initial solution concentrations (0.10, 0.20, 0.50, and 1.00 mM). Symbols represent experimental data, lines represent the Charge Distribution Multisite Surface Complexation (CD MUSIC) model predicted zeta potentials for Si and As(III) adsorption, assuming the plane of shear is located 2.6 nm from the head of the 2-plane. Suspension density 1 g L–1, background electrolyte 0.01 M NaNO3.

Single-Ion Surface Complexation Modeling
The CD-MUSIC model adequately described oxyanion adsorption data and zeta potential measurements with monomeric inner-sphere surface complexes (Fig. 1 and 2; Table 1). For the sorbate ion concentrations examined, model results indicate As(III) primarily forms a bidentate surface complex on goethite, while the surface speciation of Si shifts from a bidentate to monodentate complex as the quantity of adsorbed Si increases. The As(III) results are in good agreement with those presented by Sun and Doner (1996) who used FTIR to investigate As(III) adsorption on goethite at similar solution concentrations, and Manning et al. (1998) who used x-ray absorption fine structure (XAFS) spectroscopy to investigate As(III) binding on goethite. While no conclusive evidence exists for the speciation of Si adsorbed to goethite, a shift from a bidentate to a monodentate complex with increasing ion concentration is consistent with a decrease in the number of available surface sites for complexation.

In addition to inner-sphere complex formation, outer-sphere surface complexes were considered for As(III) single ion adsorption based on recently published data (Arai et al., 2001; Goldberg and Johnston, 2001). However, CD-MUSIC model predictions of zeta potential measurements were appreciably diminished by their inclusion (Fig. 2A). Furthermore, the value of equilibrium constants calculated for As(III) outer-sphere complexes were unrealistically large due to the need for As(III) to deprotonate at pH values well below the first pK_a value of 9.22.

The CD-MUSIC model more accurately predicted Si adsorption at low Si concentrations (0.10 and 0.50 mM) as compared with the higher loadings (1.00 mM) (Fig. 1B). This may be due in part to the potential polymerization of Si at the goethite surface (Hansen et al., 1994b). While the highest initial solution concentration for Si (1.00 mM) is well below the concentration required for precipitation of amorphous silica in solution, the goethite surface may act to catalyze Si precipitation/polymerization at lower total Si solution concentrations. Vempati et al. (1990) confirmed the existence of Si polymers on ferrihydrite when Si solution concentration equaled or exceeded 3.5 mM, while Hansen et al. (1994b) and Doelsch et al. (2001, 2003) reported the formation of Si-O-Si bond stretching for silicic acid adsorbed on ferrihydrite at initial Si solution concentrations of 0.97 mM using FTIR. Previously the formation of polymeric silica species on ferrihydrite has been reported for ratios of Si (mmol)/Fe (hydroxide) surface area (m²) as low as 0.04 (Vempati and Loeppert, 1989; Vempati et al., 1990; Hansen et al., 1994b; Doelsch et al., 2001, 2003). These values are similar to the Si/Fe ratio for the 1.00 mM Si concentration used in the current study, 0.045.

Previous efforts to model silicic acid adsorption on hydrous metal oxides have considered the formation of both monomeric and polymeric Si surface complexes from solution monomers and dimmers (Hansen et al., 1994a, 1994b; Swedlund and Webster, 1999; Meng et al., 2000; Gustafsson, 2001). Unfortunately, the inclusion of solution phase dimers in the Swedlund and Webster (1999) study did not improve complexation model predictions. Solution phase silica dimers and trimers were included in our modeling efforts in an attempt to improve the predictive capability of the CD-MUSIC model; however, their inclusion yielded adsorbed Si concentrations 20 to 50% greater than the experimentally determined surface concentrations. Moreover, given that no clear spectroscopic evidence exists to preclude sorbed polymeric silica species at the goethite –solution interface, adsorbed Si surface dimers were also used to evaluate the adsorption behavior of Si without any dramatic improvement to the model predictive capabilities. The inclusion of larger Si surface polymers was not considered due to the inherent complexity of trying to describe their formation within the limited electrostatic framework of 3-plane surface complexation models.

Equilibrium constants for As(III) and Si adsorption on goethite by ligand exchange with singly coordinated hydroxyls are presented in Table 1 (Reactions 12–19). Log K values reported in Table 1 have been calculated based on the fully protonated weak acids as to allow for comparison between the values based on the similarity of the pKa₁ values. The log K calculated for As(III) adsorption were greater than those for Si adsorption, indicating As(III) forms a more thermodynamically stable surface complex (Table 1 Reactions 12–19). The relative magnitude of thermodynamic stability constants cannot be used to predict the results of a specific interaction between two adsorbates. However, they do provide an insight as to the relative stability of one adsorbate as compared with another. Based on the single ion modeling results the data indicates that Si may not adversely influence As(III) adsorption on goethite.

Dual (Binary) Ion Adsorption
Arsenite equilibrium adsorption for either initial solution concentration investigated was not influenced by the presence of the 0.10 mM Si below pH 10 regardless of the addition scenario (Fig. 3 ). However, the 1.00 mM Si reduced As(III) equilibrium adsorption by 7 to 20% for both As(III) concentrations and addition scenarios investigated (Fig. 3). Competition effects were more pronounced at pH values above 8 (Fig. 3). As with As(III) adsorption in the presence of the low concentration of Si (0.10 mM), Si equilibrium adsorption was not affected by the presence of either As(III) concentration (Fig. 4 ). Interestingly at the higher Si concentration, there was only a reduction in Si adsorption when As(III) was equilibrated with goethite before the addition of Si (Fig. 4). In the dual ion experiments the reduction in Si adsorption (when As was added first) was always greater that the reduction in As(III) adsorption (when Si was added first).

View larger version (26K): [in this window] [in a new window]   Fig. 3. Arsenite adsorption in the presence of silicic acid as a function of As and Si concentration and addition order. Suspension density 1 g L–1, background electrolyte 0.01 M NaNO3.

View larger version (27K): [in this window] [in a new window]   Fig. 4. Silicic acid adsorption in the presence of As(III) as a function of Si and As concentration and addition order. Suspension density 1 g L–1, background electrolyte 0.01 M NaNO3.

The reduction in As(III) adsorption in the presence of the highest Si concentration is consistent with competition for sorption sites below pH 8, and a combination of electrostatic repulsion and surface site competition at pH values above 8. The enhanced reduction in the equilibrium quantity of As(III) adsorbed in the presence of 1.00 mM Si at pH values > 8 (Fig. 3) suggests increased influence of Si on the goethite surface potential. The increase in the negative potential of the surface creates an unfavorable electrostatic environment for oxyanion adsorption (Ali, 1996; Geelhoed et al., 1998; Waltham and Eick, 2002). CD-MUSIC model predictions of the single ion adsorption equilibria indicate that As(III) forms a more stable surface complex with goethite (greater log K), however this does not totally prevent Si adsorption and some competitive interaction [reduction in the quantity of As(III) adsorbed at equilibrium] is expected. Similar results for the release of As due to competitive adsorption with bicarbonate from ferrihydrite have been reported in the literature when solution oxyanion concentrations exceeded surface site density (Anawar et al., 2004; Appelo et al., 2002). Furthermore, similar reductions in As(III) adsorption in the presence of Si and an increased inhibition with increasing pH, were reported by Waltham and Eick (2002).

At the highest initial Si concentration, As(III) competition is expected based on a decrease in the available surface sites for inner-sphere complexation, as in the case when As(III) is added before Si. However, the addition scenario where Si adsorption is unaffected by the presence of As(III) was surprising. The effect of the addition scenario on Si adsorption suggests the formation of two different surface species depending on the order of Si addition. For the addition scenario when Si was added before As(III), we propose the formation of Si polymers at the goethite surface creating silica, oxygen, silica (Si-O-Si) linkages. When As(III) is added to solution we hypothesize that As(III) is able to displace the adsorbed Si, but the Si-O-Si linkages prevent its release back into the bulk solution (Fig. 5 ). This would explain the apparent lack of Si desorption from the goethite surface on the addition of As(III). As previously stated Si will polymerize on ferrihydrite surfaces to form Si-O-Si linkages at Si/Fe ratios similar to those used in the current study (Vempati et al., 1990; Doelsch et al., 2001, 2003). The observed reduction in Si adsorption when As(III) was added before Si, may result from As(III) preventing/poisoning the formation of Si polymers indicating that the polymer formation is a surface controlled process. However, it is important to note that while the order of Si addition had a significant influence on Si adsorption there was little influence on the quantity of As(III) adsorbed under either addition scenario.

View larger version (38K): [in this window] [in a new window]   Fig. 5. Schematic representation of depicting polymerized silica on the goethite surface being displaced by As(III), but not released into the bulk solution.

Dual Ion Surface Complexation Modeling
The CD-MUSIC model accurately predicted oxyanion adsorption and zeta potential for either As(III) concentration and the 0.10 mM Si concentration (Fig. 6 ). Results from modeling the low concentration of Si with either As(III) concentration suggest that the binding mechanism and surface speciation, as determined by previous spectroscopic studies, were able to describe the adsorption and zeta potential data (As 0.05 mM data not shown) (Goldberg and Johnston, 2001; Manning et al., 1998; Sun and Doner, 1996). At the 1.00 mM Si concentration the CD-MUSIC model predictions described the adsorption behavior of Si when Si was added after As(III) (Fig. 7 ). However, the CD MUSIC model did not accurately predict Si adsorption when Si was added before As(III), indicating a possible change in the Si surface complex.

View larger version (22K): [in this window] [in a new window]   Fig. 6. (A) Charge Distribution Multisite Surface Complexation (CD MUSIC) model predictions for oxyanion adsorption for the competitive adsorption of Si and As(III) on goethite. (B) CD MUSIC model predictions for zeta potential for the competitive adsorption of Si and As(III) on goethite. Model predictions are based on the equilibrium constants calculated from single ion adsorption of As(III) and Si. Symbols represent experimental data, lines represent CD-MUSIC model predictions. Suspension density = 1 g L–1, background electrolyte 0.01 M NaNO3.

View larger version (21K): [in this window] [in a new window]   Fig. 7. (a) Charge Distribution Multisite Surface Complexation (CD MUSIC) model predictions for oxyanion adsorption for the competitive adsorption of Si and As(III) on goethite. (b) CD MUSIC model predictions for zeta potential for the competitive adsorption of Si and As(III) on goethite. Model predictions are based on the equilibrium constants calculated from single ion adsorption of As(III) and Si. Symbols represent experimental data, lines represent CD-MUSIC model predictions. Suspension density = 1 g L–1, background electrolyte 0.01 M NaNO3.

Environmental Implications
Although competitive adsorption between silica and As(III) had only a slight effect on the amount of As(III) adsorbed on goethite, regardless of the order of oxyanion addition (Fig. 3), the small quantity of As(III) competitively mobilized by silica represents a dramatic increase in solution concentration. In the absence of silica, the equilibrium concentration of As(III) in solution was below or close to the current EPA maximum contaminant load (MCL) of 10 µg L^–1 over the entire pH range examined. However, in the presence of silica at concentrations representative of natural surface waters (0.10 mM or 2.8 mg L^–1) the concentration of As(III) in solution increased by 10 to 15 times, resulting in solution As levels substantially in excess of the current MCL (Fig. 8A ). The competitive release of As(III) was even more pronounced at a silica concentration characteristic of groundwater (1.0 mM or 28 mg L^–1), increasing solution As(III) concentrations by more than 100 times over the entire pH range (Fig. 8B). These results highlight the importance of understanding the mechanisms responsible for elevated As levels in source waters, such as competitive desorption of As by naturally occurring ligands like silica, and will require innovative and cost-effective treatment and remediation strategies to assure compliance with drinking water standards and minimize the health risk posed by As.

View larger version (27K): [in this window] [in a new window]   Fig. 8. Arsenite equilibrium solution concentration in the presence and absence of silicic acid. (a) As(III)T = 0.05 mM (b) As(III)T = 0.10 mM.

	REFERENCES

TOP ABSTRACT INTRODUCTION MATERIALS AND METHODS RESULTS AND DISCUSSION REFERENCES

 

• Aggett, J., and L.S. Roberts. 1986. Insight into the mechanism of accumulation of arsenate and phosphate in hydro lake sediments by measuring the rate of dissolution with ethylenediaminetraactic acid. Environ. Sci. Technol. 20:183–186.[CrossRef]
• Ali, M.A. 1996. Competitive sorption of simple organic acids and sulfate on goethite. Environ. Sci. Technol. 30:1061–1071.[CrossRef]
• Alverez, R., and D.L. Sparks. 1985. Polymerization of silicate anions in solutions at low concentrations. Nature (London) 318:649–651.[CrossRef]
• Anawar, H.M., J. Akai, and H. Sakugawa. 2004. Mobilization of arsenic from subsurface sediments by effect of bicarbonate ions in groundwater. Chemosphere 54:753–762.[Medline]
• Anderson, M.A., and D.T. Malotky. 1979. The adsorption of protolyzable anions in hydrous oxides at the isoelectric pH. J. Colloid Interface Sci. 72:413–427.
• Appelo, C.A.J., M.J.J. Van der Weiden, C. Tournassat, and L. Charlet. 2002. Surface complexation of ferrous iron and carbonate on ferrihydrite and the mobilization of arsenic. Environ. Sci. Technol. 36:3096–3103.[Medline]
• Applin, K.R. 1987. The diffusion of dissolved silica in dilute aqueous solution. Geochim. Cosmochim. Acta 51:2147–2151.
• Arai, Y., E.J. Elzinga, and D.L. Sparks. 2001. X-ray absorption spectroscopic investigation of arsenite and arsenate adsorption at the aluminum oxide-water interface. J. Colloid Sci. 235:80–88.
• Bang, S., M.D. Johnson, G.P. Korfiatis, and X.G. Meng. 2005. Chemical reactions between arsenic and zero-valent iron in water. Water Res. 39:763–770.[Medline]
• Barrón, V., and J. Torrent. 1996. Surface hydroxyl configuration of various crystal faces of hematite and goethite. J. Colloid Interface Sci. 177:407–410.[CrossRef]
• Benner, S.G., C.M. Hansel, B.W. Wielinga, T.M. Barber, and S. Fendorf. 2002. Reductive dissolution and biomineralization of iron hydroxide under dynamic flow conditions. Environ. Sci. Technol. 36:1705–1711.[Medline]
• Chapman, D.L. 1913. A contribution to the theory of electrocapollarity. Phil. Mag. and J. Sci. 25:475–481.
• Clifford, D.A., and G.L. Ghurye. 2002. Metal-oxide adsorption, ion exchange, and coagulation-microfiltration for arsenic removal from water. p. 217–246. In W.T. Frankenberger, (ed.) Environmental chemistry of arsenic. Marcel Dekker, Inc., New York.
• Costa, M.C., A.M.B. do Rego, and L.M. Abrantes. 2002. Characterization of a natural and an electro-oxidized arsenopyrite: A study on electrochemical and X-ray photoelectron spectroscopy. Int. J. Miner. Process. 65:83–108.[CrossRef]
• Davis, C.C., W.R. Knocke, and M. Edwards. 2001. Implications of aqueous silica sorption to iron hydroxide: Mobilization of iron colloids and interference with sorption of arsenate and humic substances. Environ. Sci. Technol. 35:3158–3162.[Medline]
• Devitre, R., N. Belzile, and A. Tessier. 1991. Speciation and adsorption of arsenic on diagenetic iron oxyhydroxides. Limnol. Oceanogr. 36:1480–1485.
• Dixit, S., and J.G. Hering. 2003. Comparison of arsenic(V) and arsenic(III) sorption onto iron oxide minerals: Implications for arsenic mobility. Environ. Sci. Technol. 37:4182–4189.[Medline]
• Doelsch, E., A. Masion, J. Rose, W.E.E. Stone, J.Y. Bottero, and P.M. Bertsch. 2003. Chemistry and structure of colloids obtained by hydrolysis of Fe(III) in the presence of SiO4 ligands. Colloids Surfaces A. 217:121–128.[CrossRef]
• Doelsch, E., W.E.E. Stone, S. Petit, A. Masion, J. Rose, J.Y. Bottero, and D. Nahon. 2001. Speciation and crystal chemistry of Fe(III) chloride hydrolyzed in the presence of SiO4 ligands. 2. Characterization of Si-Fe aggregates by FTIR and Si-29 solid-state NMR. Langmuir 17:1399–1405.[CrossRef]
• Ford, R.G. 2002. Rates of hydrous ferric oxide crystallization and the influence on coprecipitated arsenate. Environ. Sci. Technol. 36:2459–2463.[Medline]
• Ford, R.G., P.M. Bertsch, and K.J. Farley. 1997. Changes in transition and heavy metal partitioning during hydrous iron oxide aging. Environ. Sci. Technol. 31:2028–2033.[CrossRef]
• Frankenberger, W.T. 2002. Environmetal chemistry of arsenic. Marcel Dekker, New York.
• Fredrickson, J.K., J.M. Zachara, D.W. Kennedy, H.L. Dong, T.C. Onstott, N.W. Hinman, and S.M. Li. 1998. Biogenic iron mineralization accompanying the dissimilatory reduction of hydrous ferric oxide by a groundwater bacterium. Geochim. Cosmochim. Acta 62:3239–3257.
• Garman, S., M.J. Eick, and T.P. Luxton. 2004. Kinetics of chromate adsorption on goethite in the presence of sorbed silicic acid. J. Environ. Qual. 33:1703–1708.[Abstract/Free Full Text]
• Geelhoed, J.S., T. Hiemstra, and W.H. van Riemsdijk. 1997. Phosphate and sulfate adsorption of goethite: Single anion and competive adsorption. Geochim. Cosmochim. Acta 61:2389–2396.[CrossRef]
• Geelhoed, J.S., T. Hiemstra, and W.H. van Riemsdijk. 1998. Competive interaction between phosphate and citrate on goethite. Environ. Sci. Technol. 32:1219–1223.
• Goldberg, S., and C. Johnston. 2001. Mechanisms of arsenic adsorption on amorphous oxides evaluated using macroscopic measurements, vibrational spectroscopy, and surface complexation modeling. J. Colloid Interface Sci. 234:204–216.[CrossRef][ISI][Medline]
• Gouy, G. 1910. Sur la constitution de la charge electrique a la surface d'un eletolute. J. Phys. Radium 9:457–468.
• Grafe, M., M.J. Eick, and P.R. Grossl. 2001. Adsorption of arsenate (V) and arsenite (III) on goethite in the presence and absence of dissolved organic carbon. Soil Sci. Soc. Am. J. 65:1680–1687.[Abstract/Free Full Text]
• Gustafsson, J.P. 2001. Modeling competitive anion adsorption on oxide minerals and an allophane soiling soil. J. Soil Sci. 52:639–653.[CrossRef]
• Hansen, H.C.B., B. Rabenlange, K. Raulundrasmussen, and O.K. Borggaard. 1994a. Monosilicate adsorption by ferrihydrite and goethite at pH 3–6. Soil Sci. 158:40–46.
• Hansen, H.C.B., T.P. Wetche, K. Raulund-Rasmussen, and O.K. Borggaard. 1994b. Stability constants for silicate adsorbed to ferrihydrite. Clay Miner. 29:341–350.[Abstract]
• Herbelin, A., and J. Westall. 1999. FITEQL a computer program for determination of chemical equilibrium constants from experimental data. Version 4.0 ed. Department of Chemistry Oregon State University, Corvallis, OR.
• Hiemstra, T., and W.H. van Riemsdijk. 1996. A surface structural approach to ion adsorption: The charge distribution (CD) model. J. Colloid Interface Sci. 179:488–508.[CrossRef]
• Hiemstra, T., and W.H. van Riemsdijk. 1999. Surface structural ion adsorption modeling of competitive binding of oxyanions by metal (hydr)oxides. J. Colloid Interface Sci. 210:182–193.[CrossRef][ISI][Medline]
• Hiemstra, T., and W.H. van Riemsdijk. 2000. Fluoride adsorption on goethite in relation to different types of surface sites. J. Colloid Interface Sci. 225:94–104.[CrossRef][ISI][Medline]
• Hiemstra, T., W.H. van Riemsdijk, and G.H. Bolt. 1989a. Multisite proton adsorption modeling at the solid/solution interface of (hydr)oxides: A new approach I. Model description and evaluation of intrinsic cation constants. J. Colloid Interface Sci. 133:91–104.[CrossRef]
• Hiemstra, T., J.M.C. De Wit, and W.H. van Riemsdijk. 1989b. Multisite proton adsorption modeling at the solid/solution interface of (hydr)oxides: A new approach II. Application to various important (hydr)oxides. J. Colloid Interface Sci. 133:105–115.[CrossRef]
• Hiemstra, T., P. Venema, and W.H. van Riemsdijk. 1996. Intrinsic proton affinity of reactive surface groups of metal (hydr)oxides: The bond valence principal. J. Colloid Interface Sci. 184:680–692.[CrossRef][ISI][Medline]
• Iller, R.K. 1979. The chemistry of silica. Jon Wiley & Sons, New York.
• Inskeep, W.P., T.R. McDermott, and S.E. Fendorf. 2002. As(V)/(III) cycling in soils and natural waters: Chemical and microbiological processes. p. 183–216. In W.T. Frankenberger (ed.) Environmental chemistry of arsenic. Marcel Dekker, Inc., New York.
• Jain, A., and R.H. Loeppert. 2000. Effect of competing anions on the adsorption of arsenate and arsenite by ferrihydrite. J. Environ. Qual. 29:1422–1430.[ISI]
• Korte, N.E., and Q. Fernando. 1991. A review of arsenic(III) in groundwater. Crit. Rev. Environ. Control 21:1–39.
• Kosmulski, M. 2001. Chemical properties of material surfaces. Marcel Dekker, Inc., New York.
• Loeppert, R.L., and W.P. Inskeep. 1996. Iron. p. 639–664. In D.L. Sparks (ed.) Methods of soil analysis. Part 3. SSSA Book Ser. No. 5. SSSA, Madison, WI.
• Manning, B.A., and S. Goldberg. 1996. Modeling competitive adsorption of arsenate with phosphate and molybdate on oxide minerals. Soil Sci. Soc. Am. J. 60:121–131.[Abstract/Free Full Text]
• Manning, B.A., S.E. Fendorf, and S. Goldberg. 1998. Surface structures and stability of As(III) on goethite: Evidence for inner sphere complexes. Environ. Sci. Technol. 32:2383–2388.[CrossRef]
• McBride, M.B. 1994. Environmental chemistry of soils. Oxford Press, New York.
• Meng, X.G., S. Bang, and G.P. Korfiatis. 2000. Effects of silicate, sulfate, and carbonate on arsenic removal by ferric chloride. Water Res. 34:1255–1261.[CrossRef]
• National Institute of Standards and Technology. 2002. NIST X-ray Photoelectron Spectroscopy Database: Version 2.0 [Online]. Available by National Institute of Science and Technology http://www. nist.gov/srd/nist20.htm (posted 2003; verified 19 Oct. 2005).
• Nesbitt, H.W., L.J. Muir, and A.R. Pratt. 1995. Oxidation of Arsenopyrite by air and air-saturated, distilled water, and implications for mechanism of oxidation. Geochim. Cosmochim. Acta 59:1773–1786.
• Nickson, R.T., J.M. McArthur, P. Ravenscroft, W.C. Burgess, and K.M. Ahmed. 2000. Mechanism of Arsenic release to groundwater. Bangladesh and West Bengal. App. Geochem. 15:403–413.
• Nordstrom, D.K. 2002. Public health—Worldwide occurrences of arsenic in ground water. Science (Washington, DC) 296:2143–2145.[Abstract/Free Full Text]
• NRC. Arsenic in drinking water. 1999. Washington, DC, National Academy Press.
• NRC. Arsenic in drinking water 2001 Update. 2001. Washington, DC, National Academy Press.
• Oscarson, D.W., P.M. Huang, and W.K. Liaw. 1980. The oxidation of arsenite by aquatic sediments. J. Environ. Qual. 9:700–703.
• Oscarson, D.W., P.M. Huang, C. Defosse, and A. Herbillon. 1981. Oxidative power of Mn(IV) and Fe(III) oxides with respect to As(III) in terrestrial and aquatic environments. Nature (London) 291:50–51.[CrossRef]
• Pauling, L. 1929. The principals determining the structure of complex ionic crystals. J. Am. Chem. Soc. 51:1010–1026.[CrossRef]
• Raven, K.P., A. Jain, and R.H. Loeppert. 1998. Arsenite and arsenate adsorption on ferrihydrite: Kinetics, equilibrium, and adsorption envelopes. Environ. Sci. Technol. 32:344–349.[CrossRef]
• Reitra, R.P.J.J., T. Hiemstra, and W.H. van Riemsdijk. 1999a. The relationship between molecular structure and ion adsorption on variable charge minerals. Geochim. Cosmochim. Acta 63:3009–3015.[CrossRef]
• Rietra, R.P.J.J., T. Hiemstra, and W. van Riemsdijk. 1999b. Sulfate Adsorption on Goethite. J. Colloid Interface Sci. 218:511–521.[Medline]
• Rietra, R.P.J.J., T. Hiemstra, and W.H. van Riemsdijk. 2000. Electrolyte anion affinity and its effect on oxyanion adsorption on goethite. J. Colloid Interface Sci. 229:199–206.[Medline]
• Roden, E.E., and M.M. Urrutia. 2002. Influence of biogenic Fe(II) on bacterial crystalline Fe(III) oxide reduction. Geomicrobiol. J. 19:209–251.
• Sadiq, M. 1997. Arsenic chemistry in soils: An overview of thermodynamic predictions and field observations. Water Air Soil Pollut. 93:117–136.[CrossRef]
• Schwertmann, U., and R.M. Cornell. 1991. Iron oxides in the laboratory: Preparation and characterization VCH, Weinheim, FR of Germany.
• Schwertmann, U., and R.M. Cornell. 1996. The Iron oxides: Structures, properties, reactions, occurrences and uses. VCH, Weinheim, FR of Germany.
• Sigg, L., and W. Stumm. 1981. The interaction of anions and weak acids with the hydrous goethite (-FeOOH) surface. Colloids Surfaces 2:101–117.
• Smedley, P.L., and D.G. Kinniburgh. 2002. A review of the source, behavior and distribution of arsenic in natural waters. Appl. Geochem. 17:517–568.
• Smedley, P.L., H.B. Nicolli, D.M.J. Macdonald, A.J. Barros, and J.O. Tullio. 2002. Hydrogeochemistry of arsenic and other inorganic constituents in groundwaters from La Pampa, Argentina. Appl. Geochem. 17:259–284.[CrossRef]
• Smith, A.H., P.A. Lopipero, M.N. Bates, and C.M. Steinmaus. 2002a. Public health- Arsenic epidemiology and drinking water standards. Science (Washington, DC) 296:2145–2146.[Abstract/Free Full Text]
• Smith, R.M., and A.E. Martel. 1992.Critical stability constants. Plenum Press, New York.
• Smith, E., R. Naidu, and A.M. Alston. 1999. Chemistry of arsenic in soils: I. Sorption of arsenate and arsenite by four Australian soils. J. Environ. Qual. 28:1719–1726.
• Smith, E., R. Naidu, and A.M. Alston. 2002b. Chemistry of inorganic arsenic in soils: II. Effect of phosphorus, sodium, and calcium on arsenic sorption. J. Environ. Qual. 31:557–563.[Abstract/Free Full Text]
• Stollenwerk, K.G. 2003. Arsenic in groundwater. Kluwer Academic Publishers, Boston.
• Sun, X., and H.E. Doner. 1996. An investigation of arsenate and arsenite bonding structures on goethite by FTIR. Soil Sci. 161:865–872.
• Sun, X.H., and H.E. Doner. 1998. Adsorption and oxidation of arsenite on goethite. Soil Sci. 163:278–287.
• Swedlund, P.J., and J.G. Webster. 1999. Adsorption and polymerization of silicic acid on ferrihydrite, and its effect on arsenic adsorption. Water Res. 33:3413–3422.[CrossRef]
• Tadanier, C.J., and M.J. Eick. 2002. Formulating the charge-distribution multisite surface complexation model using FITEQL. Soil Sci. Soc. Am. J. 66:1505–1517.[Abstract/Free Full Text]
• Tadanier, C.J., T.P. Luxton, and M.J. Eick. 2002. Competitive adsorption of oxyanions. Geochim. Cosmochim. Acta 66:A759.
• United States Environmental Protection Agency. 2001. National Drinking Water Regulations; Arsenic and Clarifications to Compliance and New Source Contaminant Monitoring; Proposed Rule. 40 CFR parts 9, 141, and 142, 20579–20584. Fed. Regist 66.
• Vempati, R.K., and R.H. Loeppert. 1989. Influence of Structural and Adsorbed Si on the Transformation of Synthetic Ferrihydrite. Clays Clay Miner. 37:273–279.[Abstract]
• Vempati, R.K., R.H. Loeppert, D.C. Dufner, and D.L. Cocke. 1990. X-ray photoelectron spectroscopy as a tool to differentiate silicon-bonding state in amorphous iron oxides. Soil Sci. Soc. Am. J. 54:695–698.[Abstract/Free Full Text]
• Waltham, C., and M.J. Eick. 2002. Kinetics of arsenic adsorption on goethite in the presence of sorbed silicic acid. Soil Sci. Soc. Am. J. 66:818–825.[Abstract/Free Full Text]
• Waychunas, G.A., B.A. Rea, C.C. Fuller, and J.A. Davis. 1993. Surface-chemistry of ferrihydrite.1. Exafs studies of the geometry of coprecipitated and adsorbed arsenate. Geochim. Cosmochim. Acta 57:2251–2269.
• Welch, A.H., D.B. Westjohn, D.R. Helsel, and R.B. Wanty. 2000. Arsenic in ground water of the United States: Occurrence and geochemistry. Ground Water 38:589–604.[CrossRef][ISI]
• Winship, K.A. 1984. Toxicity of inorganic arsenic salts. Adverse Drug Reac. Toxicol. Rev. 3:129–160.
• Zhang, W., P. Singh, E. Paling, and S. Delides. 2004. Arsenic removal from contaminated water by natural iron ores. Miner. Eng. 17:517–524.[CrossRef]

This Article Abstract Figures Only </