Chemical Geology 368 (2014) 31–38

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Chemical Geology

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Kinetics of competitive adsorption/desorption of arsenate and phosphateat the ferrihydrite–water interface

Ghanashyam Neupane a,b,⁎, Rona J. Donahoe c, Yuji Arai d

a University of Idaho-Idaho Falls, 1776 Science Center Drive, Suite 306, Idaho Falls, ID 83402, United Statesb Center for Advanced Energy Studies, 995 University Boulevard, Idaho Falls, ID 83401, United Statesc Department of Geological Sciences, University of Alabama, Box 870338, Tuscaloosa, AL 35487, United Statesd Department of Natural Resources and Environmental Sciences, 1102 South Goodwin Avenue, University of Illinois at Urbana-Champaign, Urbana, IL 61801, United States

⁎ Corresponding author at: University of Idaho-Idaho FSuite 306, Idaho Falls, ID 83402, United States. Tel.: +1 27929.

E-mail address: gneupane@uidaho.edu (G. Neupane).

0009-2541/$ – see front matter © 2014 Elsevier B.V. All rihttp://dx.doi.org/10.1016/j.chemgeo.2013.12.020

a b s t r a c t

a r t i c l e i n f o

Article history:Received 21 July 2013Received in revised form 24 December 2013Accepted 29 December 2013Available online 9 January 2014

Editor: J. Fein

Keywords:ArsenatePhosphateFerrihydriteAdsorptionDesorptionAs K-edge EXAFS

Metal hydroxides (e.g. ferrihydrite) present in geomedia play significant roles in regulating the environmentalmobilities of arsenate (As(V)) and inorganic phosphate (Pi) because of their high adsorption affinities for theseoxyanions. In this study, results are presented of experiments aimed at determining individual and competitiveadsorption/desorption kinetics of As(V) and Pi on ferrihydrite at pH 4 and 8. Selected samples were also subject-ed to As K-edge EXAFS study for understanding the changes with time in As(V) complexation on ferrihydrite inthe presence/absence of Pi. Both oxyanions showed similar behavior in single ion adsorption experiments. How-ever, when both oxyanions were loaded together, more As(V) was adsorbed than Pi. Furthermore, more pre-equilibrated Pi was desorbed by sequentially added As(V) than vice versa. Interactions of As(V) and Pi with fer-rihydrite slowed down after the initial rapid adsorption/desorption. The experimentally determined adsorption/desorption kinetic data for As(V) and Pi showedgood compliancewith pseudo-secondorder, Elovich, and power-function equations. Both oxyanions competed for adsorption on ferrihydrite, and each of them showed a limitedcapacity to desorb the other. EXAFS analysis of selected samples indicated the presence of mononuclear (2E) andbinuclear (2C) bidentate As(V) surface complexes. The Fe coordination numbers (CN) increased with increasingtime and decreased with addition of Pi into the system. A higher proportion of Fe CN associated with 2E As(V)surface complexes decreased after the addition of Pi, compared to Fe CN associated with 2C As(V) surface com-plexes. The competitive desorption study indicates that the excessive input of Pi due to the overuse of fertilizerscould mobilize As(V) from contaminated geomedia. Furthermore, insights into Pi-induced desorption of As(V)could also provide an opportunity for developing chemical treatment methods to intercept the mobilizedAs(V) by co-precipitation in apatite-like phases.

© 2014 Elsevier B.V. All rights reserved.

1. Introduction

Despite being in the same periodic group, As is a toxic elementwhileP is an essential element to the biosphere. Elevated concentrations of Asin the environment can result from natural and anthropogenic activities(Smedley andKinniburgh, 2002). For example, high levels of As in soil atseveral industrial sites in the southeastern United States have beentraced to the heavy application of As-based herbicides in the past(Yang and Donahoe, 2007). In the environment, inorganic As can existin several valance states (0, +3, +5), depending on the prevailing pHand redox conditions (Cullen and Reimer, 1989). In natural waters, Asmostly occurs as inorganic arsenite (As(III)) and arsenate (As(V)), al-though some organic As compounds are also reported in water severelyimpacted by industrial pollution (Smedley and Kinniburgh, 2002).Most

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of the P contamination of surfacewaters is linked to the input of excess Pfromagricultural land runoff (Withers et al., 1999). The P contaminationof surface waters is reported to cause severe environmental problemssuch as eutrophication (Correll, 1998). Phosphorous only occurs in pen-tavalent form in the environment, as ortho-phosphates (Pi), pyro-phosphates, longer-chain polyphosphates, and several organic phos-phates (Correll, 1998).

The mobility of these oxyanions in the environment, particularly ingeomedia, depends on the redox potential and pH conditions, the min-eralogy and organic matter content of the substrate material, and thecompeting oxyanions present (Smedley and Kinniburgh, 2002). BothAs(V) and Pi show a high degree of adsorption on Fe-, Mn-, and Al-hydroxides in geomedia (Violante and Pigna, 2002). Adsorption ofAs(V) and Pi on ferric hydroxides involves ligand exchange for OH2

and OH‐ in the coordination sphere of surface-exposed structural Featoms.

Individual ion adsorption experiments showed a rapid initial uptakeof both As(V) and Pi on ferric hydroxide followed by slow sorption(Fuller et al., 1993; Strauss et al., 1997). During competitive adsorption

32 G. Neupane et al. / Chemical Geology 368 (2014) 31–38

experiments on soil, Zhang and Selim (2007) observed a decrease in therate and extent of As(V) sorption as the Pi concentration increased.Phosphate-induced desorption of pre-adsorbed As(V) on goethite wasfound by O'Reilly et al. (2001) to be relatively rapid, releasing N35% ofthe total adsorbed As. Although As(V) and Pi showed similar adsorptionbehavior during individual ion adsorption experiments on gibbsite andgoethite, their adsorption behavior differed in competitive adsorptionexperiments (Hingston et al., 1971). Based on competitive adsorptionof oxyanions on these adsorbents, Hingston et al. (1971) suggestedthe presence of oxyanion-specific and competitive adsorption sites onoxide surfaces. Several researchers, (e.g., Gao and Mucci, 2001; Liuet al., 2001; Violante and Pigna, 2002; Carabante et al., 2010) have re-ported greater adsorption of As(V) than Pi on several iron oxides andiron-rich clay minerals in competitive adsorption experiments. Howev-er, the likely mechanism of preferential adsorption of As(V) over Pi onferrihydrite is poorly understood. Similarly, we still lack an understand-ing of changes over time in As(V) surface complexation in the presenceof Pi.

In this study, we conducted experiments examining As(V) and Pi in-teraction over time with ferrihydrite-type ferric hydroxide at pH 4 andpH 8. Ferrihydrite is one of the common ferric hydroxide minerals inlow-temperature oxic environmentswith a poorly-crystalline structure,large surface area, and high affinity for sequestering As(V) and Pi. Thismineral primarily exists as a surface coating material for soil and sedi-ment particles, and As has been reported to be bound with Fe-oxides/hydroxides in contaminated soils (Yang and Donahoe, 2007). Similarly,several ferric hydroxide-based methods have been used for removingAs from drinking water and for in situ fixation of As in contaminatedsoils. Owing to the similar chemical behavior of these oxyanions in theenvironment, this study was aimed at elucidating the individual ionand competitive adsorption kinetics of As(V) and Pi on ferric hydroxideat pH 4 and pH 8. This research was also aimed at studying the co-oxyanion induced desorption kinetics of As(V) and Pi on the same sub-strate. Time-dependent adsorption/desorption experiments usingAs(V) and Pi with individual, simultaneous, and sequential ion loadingswere performed for quantifying the adsorption/desorption kinetics andstudying the associated changes in As(V) complexation on ferrihydrite.

2. Materials and methods

2.1. Ferric hydroxide synthesis

In this study, a ferrous sulfate heptahydrate (FeSO4·7H2O) solutionwas selected to prepare freshly-precipitated ferric hydroxide (ferrihy-drite) because it was previously used as a treatment solution for insitu chemical fixation of As and other trace elements in contaminatedsoil and coal fly ash samples (Yang et al., 2007; Bhattacharyya et al.,2009; Neupane et al., 2010). About 200 g of analytical grade FeSO4·7H2Owas placed in a 2 L reaction vessel containing 1 L of≥18 MΩ doubly de-ionized (DDI) water and stirred continuously with a magnetic stirrer. A30% H2O2 solution was added in 4 steps (about 10 mL per addition) tothe ferrous sulfate solution to promote rapid oxidation. Simultaneously,a 5 MNaOH solutionwas also gradually added to raise the pH of the slur-ry to 8. The ferric hydroxide precipitate was collected by sedimentationand rinsed in two cycles: the first rinsing cycle was performed withoutpH adjustment; during the second rinsing cycle, pH was adjusted to 11using additional 5 M NaOH solution to remove any remaining sulfate.Each rinsing cycle consisted of several serial additions of DDI water tothe reaction vessel followed by agitation, sedimentation, and decantationof the supernatant. The sulfate concentration in the decanted rinse waterwas measured by ion chromatography (Dionex LC 20). Once the sulfateconcentration in the decanted rinse water fell below 0.2 mg/L duringthe second rinsing cycle, the synthesized ferric hydroxide was separatedby vacuum filtration and freeze-dried. The X-ray powder diffraction(XRD) spectrum of the synthesized ferric hydroxide corresponded to atwo-line ferrihydrite. The single-point BET surface area (TriStar 3000

BET Surface Analyzer) and the Point of Zero Net Proton Charge (PZNPC)(acid–base titration) of the ferrihydrite were determined to be273 m2/g, and 8.1, respectively.

2.2. Adsorption/desorption kinetic experiments and modeling

The kinetics at pH 4 and 8 of single ion adsorption, competitive ad-sorption, and co-ion induced desorption of As(V) and Pi were studiedexperimentally at room temperature and pressure. Arsenate and Pistock solutions were prepared from analytical grade Na2HAsO4·7H2Osalt (Alfa Aesar, Ward Hill, MA) and 85% H3PO4 acid (Fisher Scientific,Hampton, NH), respectively, and their concentrations were determinedby inductively coupled plasma-optical emission spectroscopy (ICP-OES)using a Perkin Elmer Optima 3000DV instrument. Each set of experi-ments consisted of 1 L of pH-adjusted 2 g/L ferrihydrite suspension in0.1 M NaCl electrolyte placed in a 1.5 L reaction vessel. The suspensionwas spiked to 1.5 mM of As(V), Pi, or both, depending on the loadingscheme used. Three loading schemes were used at each pH:

a) As(V) before Pi,b) Pi before As(V), andc) simultaneous loading of As(V) and Pi.

In schemes (a) and (b), Pi and As(V) were sequentially loaded after48 h of the initial loadings of As(V) and Pi, respectively. From the twosequential loading schemes, individual ion adsorption kinetic datawere obtained during the first 48 h for As(V) and Pi.With sequential ad-dition of the co-oxyanion (Pi for scheme (a) and As(V) for scheme (b)),the competitive adsorption kinetics for Pi and As(V)were obtained. Fur-thermore, sequential loading of the co-oxyanion also provided compet-itive desorption kinetics for the pre-equilibrated oxyanion (As(V) forscheme (a) and Pi for scheme (b). In addition to the sequential loadingschemes, competitive adsorption kinetic data were also obtained by si-multaneous loading of both oxyanions (scheme (c)). Changes in pHwere monitored continuously during the experiments and adjusted tothe target valuewith 0.1 MNaOH/HCl solutions. Throughout the exper-iment, the suspension was agitated by a magnetic stirrer. From each re-action vessel, 20 mL of suspension was extracted at each samplingevent. Sampling frequencywas higher immediately after loadings (sep-arately, in sequence, or simultaneously). After each extraction, the sam-ple was centrifuged at 13,000 g for 10 min (2 min for the early kineticsamples) and filtered throughWhatman puradisc 0.2 μm nylon syringefilters, acidified to 2% HNO3 with OPTIMA® ultrapure nitric acid andstored in a refrigerator until chemical analysis by ICP-OES.

The adsorption of As(V) (or Pi) was determined using mass-balancecalculationswith reference to the initial concentration and the superna-tant concentration at each sampling point. The Elovich (Low, 1960),power function (Aharoni and Sparks, 1991), and pseudo-second order(Ho and McKay, 1999) equations (Table S1) were used for modelingthe single ion as well as competitive adsorption/desorption kinetics ofAs(V) and Pi on ferrihydrite. Our attempts to use other kinetic models(e.g. first order, second order, etc.) failed because of their inability torepresent the experimental data.

2.3. As K-edge EXAFS spectroscopy

2.3.1. EXAFS samples and data collectionSome kinetic experiment samples were selected for study using As

K-edge extended X-ray absorption fine structure (EXAFS) spectroscopy.Ferrihydrite pastes with sorbed As(V) (with or without Pi) were packedin Teflon sample holders, sealed with a Kapton® tape, and stored(b1 week) in a refrigerator/ice cooler. The EXAFS spectrawere collectedon bending magnet beamline 9 BM B at the Advanced Photon Source,Argonne National Laboratory, Argonne, IL. The storage ring was operat-ed at 7.0 GeV with a maximum ring current of 100 mA. Two siliconcrystals (presenting the Si(111) and Si(220) panes) were used as a

33G. Neupane et al. / Chemical Geology 368 (2014) 31–38

monochromator which was calibrated to 11,867 eV using a 0.1 Na2-HAsO4·7H2O solution prior to As EXAFS data collection. The sample hold-er was oriented at 45° to the unfocussed incident beam. Arsenic EXAFSspectra were collected from pre-edge to extended regions in 0.3 to10 eV steps with a 2 sec measurement time and 0.5 sec settling time ateach step. For each sample, at least three spectrawere collected influores-cencemode at room temperature by a 13-elementGe solid-state detector.

2.3.2. EXAFS data reductionArsenic K-edge EXAFS data reduction was performed with Athena

0.8.056 software (Ravel and Newville, 2005). At least two spectra ofeach sample with similar baseline and backgroundwere merged beforebackground removal with the AUTOBK algorithm (Newville et al.,1993). The spectra were normalized with respect to E0 determinedfrom the second derivative, converted from E-space to k-space, andexported to the Artemis 0.8.012 program (Ravel and Newville, 2005)for model fitting using the IFEFFIT code (Newville, 2001). The prelimi-nary structural parameters for the model compound were obtainedwith ATOMS (Ravel, 2001) using scorodite (FeAsO4·2H2O) crystallinedata (Kitahama et al., 1975). The theoretical amplitude and phase shiftsfor the selected paths (a single scattering path of As–O, two separatesingle scattering paths of As–Fe, and amulti scattering (within arsenatetetrahedra) path of As–O(MS) were calculated using the FEEF6 code(Zabinsky et al., 1995). The radial structural functions were obtainedby Fourier-transforming the k-space data over a Δk of 3–11 Å−1.

Fig. 1. EXAFS data showing a) normalized and background subtracted k3-weighted As K-edge speSolid lines represent experimental data for As(V)-equilibrated ferrihydrite samples at pH 4 and8tal data.

Then, the fitting of the EXAFS data was performed in R-space with afitting range of 1–3.5 Å. During fitting, the values of coordination num-ber (CN) and inter-atomic bond distance (R) were allowed to float. TheE0 value was obtained by fitting the first As–O shell distance, and thenfixed for the fitting of all other shells. The amplitude reduction factor(S02) was fixed at 0.85. The Debye–Waller factor (σ2) for As–O single-scattering was allowed to float, while a fixed value of 0.005 was usedfor the As–Fe shells and 0.007 was used for As–O(MS). The accuraciesof the model-derived parameters were ±0.02 Å for RAs–O, ±0.03 Å forRAs–Fe, ±14% for CNAs–O, and ±41% for CNAs–Fe.

3. Results and discussion

3.1. As K-edge EXAFS analysis

Fig. 1 shows the k3-weighted As K3-edge EXAFS spectra and the cor-responding Fourier-transformed spectra for the kinetic experimentsamples collected at 5 min, 4 h (pH 4 only), and 24 h after the additionof As(V) at pH 4 and 8. Each As(V) EXAFS spectrum shown in Fig. 1a ischaracterized by strong sinusoidal oscillations originating from theback-scattering of the surrounding oxygen atoms. Primarily, the nearsymmetrical k-space spectral feature between 4 and 6 Å−1 was qualita-tively used as an indication of an adsorption-dominatedmechanism forAs(V) interaction on ferrihydrite (Chen et al., 2009). For the equilibriumaqueous concentrations of Fe and initial concentration of As(V),

ctra, and b) Fourier-transformed k3-weighted As K3-edge spectra (phase shift uncorrected)., whereas open circles represent the theoretical non-linear least-square fit to the experimen-

34 G. Neupane et al. / Chemical Geology 368 (2014) 31–38

PHREEQC (Parkhurst and Appelo, 1999) simulations indicated thatall experimental systems were undersaturated with respect to crys-talline scorodite (log Ksp = −25.83) and amorphous ferric arsenate(log Ksp = −23.0) (Langmuir et al., 2006).

The As–O oscillations resulted in a strong peak in the Fourier-transformed spectra (Fig. 1b). This peak represents the first coordina-tion shell of the four O atoms at a distance of 1.70–1.71 Å from the cen-tral As atom (Foster et al., 1998). Two As–Fe shells were identified forAs(V) complexation on ferrihydrite. The As–Fe1 shell is located at a dis-tance of 2.82–2.90 Å, and has been attributed to the bidentatemononu-clear edge-sharing complex (2E) (Waychunas et al., 2005). Similarly,the As–Fe2 shell, defined with a distance of 3.29–3.38 Å, has been at-tributed to the bidentate binuclear corner sharing complex (2C)(Waychunas et al., 2005). Multiple scattering of As–O(MS) contributeda broad peak at a distance of 3.18–3.32 Å to the overall Fourier-transformed spectrum. The inclusion of As–O(MS) slightly increasedthe quality of the fit; however, it was not enough to discard the 2Eshell, as suggested by Sherman and Randall (2003).

Although the Fe CN values obtained with EXAFS were relatively lessaccurate (±41%), the intensities of 2E and 2C varied in samples withtime (Table 1 and Fig. S1). Particularly, the Fe CN values for 2E increasedwith increasing contact time, but decreased after the addition of Pi to thesystem. After the addition of Pi to the As(V)-equilibrated system, nearly

Table 1Structural parameters obtained from least-squares analysis of As K-edge EXAFS spectra ofAs(V)-adsorbed ferrihydrite at pH 4 and 8 at different times and loading schemes.

Sample Shells σ2 (Å2)a ΔE0 (eV) R (Å)b CNc R-factor

pH 4: As(V) 5 min As–O 0.0010 8.55 1.71 4.06 0.022As–Fe1 0.005d 2.84 0.80As–Fe2 0.008d 3.32 1.87As–O(MS) 0.009d 3.26 12.30

pH 4: As(V) 4 h As–O 0.0036 9.97 1.71 4.25 0.010As–Fe1 0.005d 2.86 1.11As–Fe2 0.008d 3.38 1.82As–O(MS) 0.009d 3.32 10.02

pH 4: As(V) 24 h As–O 0.0008 9.57 1.71 3.84 0.006As–Fe1 0.005d 2.83 1.19As–Fe2 0.008d 3.38 1.75As–O(MS) 0.009d 3.30 10.26

pH 4: As(V)48 h Pi 5 min

As–O 0.0012 6.74 1.70 3.63 0.026As–Fe1 0.005d 2.90 0.76As–Fe2 0.008d 3.33 1.78As–O(MS) 0.009d 3.21 10.68

pH 4: As(V)48 h Pi 4 h

As–O 0.0023 5.81 1.70 4.24 0.021As–Fe1 0.005d 2.92 0.52As–Fe2 0.008d 3.28 1.75As–O(MS) 0.009d 3.20 12.85

pH 4: As(V)48 h Pi 24 h

As–O 0.0022 7.35 1.70 4.22 0.018As–Fe1 0.005d 2.86 0.48As–Fe2 0.008d 3.36 1.73As–O(MS) 0.009d 3.22 9.76

pH 8: As(V) 5 min As–O 0.0003 8.21 1.71 3.51 0.019As–Fe1 0.005d 2.88 0.67As–Fe2 0.008d 3.39 1.32As–O(MS) 0.009d 3.30 8.75

pH 8: As(V) 24 h As–O 0.0018 7.31 1.70 4.24 0.020As–Fe1 0.005d 2.84 0.72As–Fe2 0.008d 3.29 1.81As–O(MS) 0.009d 3.24 10.84

pH 8: As(V)48 h Pi 5 min

As–O 0.0011 8.46 1.71 3.96 0.016As–Fe1 0.005d 2.82 0.45As–Fe2 0.008d 3.32 1.82As–O(MS) 0.009d 3.26 11.88

pH 8: As(V)48 h Pi 24 h

As–O 0.0014 8.43 1.70 3.38 0.027As–Fe1 0.005d 2.89 0.40As–Fe2 0.008d 3.31 1.76As–O(MS) 0.009d 3.22 11.26

a Debye–Waller factor.b Interatomic distance.c Coordination number.d Fixed.

60% and 44% of the Fe CN associated with 2E was decreased at pH 4 and8, respectively, 24 h after the addition of Pi. However, changes in the FeCN associated with 2C after the addition of Pi were relatively small andinconclusive. Despite observing a greater desorption of 2E As(V) fromferrihydrite, there is no previous literature that supports the formationof a 2E Pi complex on ferrihydrite. Rose et al. (1997) only demonstratedthe formation of corner-sharing complexes during the hydrolysis ofFeCl3 in the presence of Pi. Therefore, it is likely that Pi directly substi-tutes for 2C complexed As(V), whereas the displacement mechanismfor 2E As(V) on ferrihydrite by Pi needs further evaluation.

3.2. Adsorption kinetics of As(V) and Pi

Figs. 2 and 3 show the results for individual ion as well as competi-tive adsorption/desorption of As(V) and Pi on ferrihydrite as a functionof time. In our experiments, both oxyanions showed rapid initial ad-sorption until 4 h (Fig. 2). At pH 4, adsorption of As(V) and Pi reached0.55 mmol/g (73%) and 0.61 mmol/g (82%), respectively, in 5 min.Such rapid adsorption of As(V) and Pi on iron oxides has been previous-ly reported by Fuller et al. (1993) and Luengo et al. (2007). However,after the initial rapid adsorption, the kinetics became more sluggish,reaching 0.73 mmol/g (97%) adsorption of both oxyanions in 48 h. Sim-ilar trends, but with lower overall adsorption, occurred at pH 8. In thefirst 5 min, 0.31 mmol/g (41%) of the As(V) and 0.33 mmol/g (44%) ofthe Pi were transferred from solution to the solid phase. In 48 h of reac-tion, these figures reached to 0.43 mmol/g (57%) and 0.44 mmol/g(58%) for As(V) and Pi, respectively.

Competitive adsorption kinetics for Pi and As(V) were studied usingboth sequential and simultaneous loadings. Experiments for competi-tive adsorption kinetics with sequential loadings started after additionof As(V) or Pi to a system where the co-oxyanion had been previ-ously equilibrated with ferrihydrite for 48 h (Fig. 2). At pH 4 and8, 0.30 mmol/g (40%) and 0.27 mmol/g (36%) of the Pi, and0.26 mmol/g (35%) and 0.20 mmol/g (27%) of the As(V), were compet-itively adsorbed 5 min after sequential loadings, respectively. Withsimultaneous loading (Fig. 3), 0.41 mmol/g (55%) of the As(V) and0.33 mmol/g (44%) of the Pi were adsorbed within 5 min at pH 4. Sim-ilarly, 0.24 mmol/g (31%) of the As(V) and 0.20 mmol/g (27%) of the Piwere adsorbed within 5 min at pH 8. After 4 h, however, adsorption ofboth As(V) and Pi slowed with time.

Fig. 2. Individual ion and competitive desorption kinetic data for As(V) and Pi. Initial spikesusing As(V) or Pi stock solutions produced a 1.5 mM oxyanion concentration at pH 4 or 8in a pH-equilibrated 0.1 M NaCl solution with 2 g/L ferrihydrite, which was allowed toreact for 48 h to determine of individual ion adsorption kinetic data. After 48 h, the sus-pensionwas spiked to 1.5 mMof the co-oxyanion to obtain competitive adsorption kineticdata for the sequentially added co-oxyanion and competitive desorption kinetic data forthe initially loaded oxyanion.

Fig. 3. Competitive adsorption kinetic data for As(V) and Pi in a 0.1 M NaCl solution withsimultaneous loadings.

Fig. 4. Comparison of competitive adsorption/desorption kinetic data for each oxyanion(As(V) or Pi) upon adsorption of the sequentially loaded co-oxyanion (Pi or As(V)). De-sorption of the pre-equilibrated oxyanion (open symbols) is significantly lower than thecompetitive adsorption of the sequentially added co-oxyanion (filled symbols).

35G. Neupane et al. / Chemical Geology 368 (2014) 31–38

3.3. Competitive desorption kinetics of As(V) and Pi

Competitive desorption of an oxyanion (e.g., As(V)) occurred afterthe sequential addition of the co-oxyanion (e.g., Pi) to the equilibratedoxyanion/ferrihydrite system (Fig. 2). Both As(V) and Pi were partiallyeffective in desorbing the respective co-oxyanion from ferrihydrite.The desorbed amounts of As(V) after 5 min of Pi spiking to 1.5 mM atpH 4 and 8 were 0.06 mmol/g and 0.01 mmol/g, accounting for about8% and 2% of the adsorbed of As(V), respectively. The desorbed amountsof Pi after 5 min of As(V) spiking to 1.5 mmol at pH 4 and 8 were rela-tively higher, 0.10 mmol/g and 0.05 mmol/g, accounting for about 14%and 11% of the adsorbed amounts of Pi, respectively. Competitive de-sorption of Pi and As(V) occurred persistently with increasing contacttime, but at diminishing rates. Within 4 h, the co-oxyanion induced de-sorption of As(V) and Pi increased to 0.10 mmol/g and 0.17 mmol/g atpH 4, and to 0.04 mmol/g and 0.10 mmol/g at pH 8, respectively. Littledesorption of either oxyanion occurred after 4 h at pH 4 or 8. The smallamount of Pi-induced desorption of As(V) at pH 8 was also reflected byrelatively small decreases in the CNs of As-Fe shells (Fig. S1).

3.4. Kinetic modeling

The pseudo-second order, Elovich, and power-function kineticmodels(Table S1) were used to describe the individual/competitive adsorptionand desorption kinetic data. The kinetic parameters for these models arepresented in Table S2, alongwith corresponding coefficient of determina-tion (R2) values. Linear plots of pseudo-second order, Elovich, and powerfunction kinetic equations for single ion adsorption, competitive adsorp-tion after sequential and simultaneous loadings, and competitive desorp-tion, are shown in Figs. S2 through S4. All three models were able topredict the general adsorption/desorption kinetics of As(V) and Pi.

Pseudo-second order kinetic equations (Table S2) produced excel-lent fits (R2 N 0.99) for adsorption/desorption data measured for bothoxyanions during different loading schemes. Other researchers havealso shown good correspondence between pseudo-second order equa-tions and empirical kinetic data for Pi adsorption on multi-componentAl(III)–Fe(III) hydroxide (Harvey and Rhue, 2008) and on α-Al2O3

(Del Nero et al., 2010). The Elovich equation adequately describedAs(V) and Pi adsorption/desorption kinetics, but with somewhatlower coefficient of determination values (R2 N 0.88) (Table S2). TheElovich equation, which is based on the assumption that the adsorptionenergy increaseswith surface coverage (Low, 1960), has been previous-ly used to describe the kinetics of adsorption/desorption of As(V)(Raven et al., 1998) and Pi (Chien and Clayton, 1980) on soil and its con-stituents. Chien and Clayton (1980) suggested that an increase in α,with or without changes in β, would indicate a relatively rapid reaction.However, any mechanistic inference drawn from the nature of α and βwas deemed questionable for other adsorption conditions (Sparks,

1989). The power-function equation described the experimental datawith R2 N 0.76 (Table S2). The relatively lower R2 values were particu-larly associated with desorption, and therefore might indicate that thepower function equation is not equally suitable for describing the de-sorption kinetics of these oxyanions.

During the kinetic experiments, individual ion aswell as competitiveadsorption/desorption of As(V) and Pi to/from ferrihydrite (Figs. 2–4)showed initially rapid, followed by slow, adsorption/desorption withtime. At pH 4 and 8, significant adsorption (N0.55 mmol/g) of As(V)and Pi occurred by 5 min. This rapid initial uptake of oxyanions by ferri-hydrite is attributed to adsorption on surface/near-surface sites, where-as the later slow uptake (absorption) could result either from thesaturation of specific adsorption sites, or the migration of oxyanionsinto the ferrihydrite structure (Willett et al., 1988).

The excellent correspondence of the oxyanion adsorption kineticdata with pseudo-second order equations may indicate a sorptionmechanism in which the rate-limiting step is associated with chemi-sorption that involves valence forces through the sharing or exchangeof electrons between the sorbent and sorbate (Ho and McKay, 1999).This is consistent with the observation that both As(V) and Pi adsorbon ferric-hydroxide by forming inner-sphere surface complexes(Tejedor-Tejedor and Anderson, 1990; Waychunas et al., 1993). Fur-thermore, As K-edge EXAFS analysis of selected samples used in thisstudy showed the presence of two surface complexes that have differentcomplexation trends over time. At pH 4, the 2C coordination number(CN) did not change from the earliest sample (5 min) to the sampletaken after 24 h, or for the sample taken after 72 h (i.e., 24 h after intro-duction of Pi) (Table 1 and Fig. S1). On the other hand, the CN associatedwith 2E increased consistently over 24 h, and then decreased after the in-troduction of Pi. Similarly, at pH 8, the CN associated with 2C was higherfor the 24 h sample than for the 5 min sample, and remained similarafter the addition of Pi. However, the increase in CN associated with 2Efrom the 5 min sample to the 24 h sample was very small. At both pHconditions, the CN associated with 2E decreased with the introductionof Pi. The consistently uniform CN associated with 2C at pH 4 after5 min could indicate that the As(V)–ferrihydrite interactions formingthis complex reach an equilibrium within a short timeframe (b5 min),and therefore the trend might have been missed. However, the changesin CN associated with 2E, particularly at pH 4, could suggest that therate-limiting adsorption of the oxyanion (e.g., As(V)) is related to satura-tion of specific surface sites. It has been previously reported that forma-tion of the 2C surface complex is energetically favored over formation ofthe 2E surface complex (Sherman and Randall, 2003). During single ionadsorption experiments, the initial rapid adsorption could have repre-sented formation of the 2C complex; however, as the specific surfacesites for this complex became less available, the oxyanions might havestarted to form the energetically less favorable 2E complex. However,

Fig. 5.Molar adsorption ratio (MAR) values calculated as adsorbed As(V)/adsorbed Pi (oras desorbed As(V)/desorbed Pi) for different loading schemes.

36 G. Neupane et al. / Chemical Geology 368 (2014) 31–38

thismechanism is not conclusive at pH 8. It also should be noted that theexistence of the 2E complex has not been confirmed for Pi by previousstudies (e.g., Rose et al., 1997).

Another likely mechanism for the observed slower adsorption overtime is the migration of these oxyanions into the ferrihydrite particles.Previously,Willett et al. (1988) provided evidence for inwardmigrationof Pi from the surface of ferrihydrite particles with time. Similarly,Strauss et al. (1997) also reported that Pi ions were rapidly adsorbedon the charged external surface of goethite and then slowly diffusedinto the particles. Assuming Pi as an analog of As(V), we could have asimilar mechanism at work for the latter. Likewise, the co-oxyanion in-duced desorption of As(V) and Pi was observed to be rapid initially, butslowedwith time (Fig. 4). Therefore, the early rapid desorption could beattributed to displacement of the equilibrated oxyanion (e.g., As(V))from surface sites by the competing oxyanion (e.g., Pi), whereas thelater slower desorption may result from outward-displacement of inte-rior ions by the competing oxyanion as it migrates into the aggregateparticles (Willett et al., 1988; Strauss et al., 1997). Finally, the goodagreement between the experimental data and the kinetic model equa-tionsmight indicate that early rapid adsorption/desorption is controlledby surface interactions,whereas the inwardmigration of oxyanion is therate-limiting factor during the remainder of the adsorption/desorptionexperiments (Low, 1960; Willett et al., 1988; Aharoni and Sparks, 1991).

3.5. Comparison of individual ion and competitive adsorption kinetics ofAs(V) and Pi

At both pH 4 and 8, the adsorption of each oxyanion was greaterwhen it was loaded separately, compared to simultaneous loading ofboth oxyanions (Figs. 2 and 3). The observed decrease in adsorption ofPi was relatively greater than the decrease in adsorption of As(V) inthe competitive adsorption experiments. At pH 4, nearly complete(0.73 mmol/g or 97%) adsorption of each oxyanion occurred during in-dividual ion adsorption experiments, while only 76% (0.57 mmol/g) ofAs(V) and 57% (0.43 mmol/g) of Pi were adsorbedwhen loaded togeth-er. Similarly, 57% (0.43 mmol/g) of the As(V) and 58% (0.44 mmol/g) ofthe Pi were adsorbed during individual ion adsorption experiments atpH 8, but decreased to 50% (40% for Pi) at the same pH during compet-itive adsorption experiments. The loading scheme used also resulted ina minor change in the competitive adsorption of As(V): slightly greateradsorption of As(V) occurred with simultaneous loading (0.57 mmol/gor 76% at pH 4; 0.37 mmol/g or 50% at pH 8) thanwith sequential load-ing (0.53 mmol/g or 70% at pH 4 and 0.33 mmol/g or 44% at pH 8).However, a similar adsorption of Pi [57% (0.43 mmol/g) vs. 61%(0.46 mmol/g) at pH 4 and 40% (0.30 mmol/g) vs. 42% (0.33 mmol/g)at pH 8] occurred with simultaneous vs. sequential loading schemes.The pseudo-second order reaction rates (k and h), Elovich constant(α), and power function rate constant (a) for competitive adsorptionof As(V) and Pi are lower than the corresponding rates/constantsobtained for individual ion adsorption (Table S2). Therefore, thepresence of the co-oxyanion not only decreased the partial surfacecoverage of the initially loaded oxyanion, but also changed theadsorption kinetics.

3.6. Adsorption kinetics and pH

The influence of pH on As(V) and Pi adsorption/desorption kineticscan be evaluated by the observation of linear Elovich and power func-tion plots (Figs. S3 and S4). The model fitting of the experimental dataat pH 4 and 8 produced lines which were parallel or sub-parallel toeach other, except for the data for competitive adsorption with sequen-tial loading. This indicates that adsorption/desorption reaction rates arelargely independent of pH when the oxyanions are loaded either sepa-rately or simultaneously. Zhang and Stanforth (2005) also reported par-allel linear Elovich fits to the experimental data collected at pH 4, 7, and

10, and suggested pH-independent kinetics for As(V) adsorption ongoethite.

3.7. Molar adsorption ratios

The adsorption preference of ferrihydrite for As(V) versus Pi duringindividual ion as well as competitive adsorption settings was evaluatedby calculating the molar adsorption ratio (MAR) of adsorbed As(V) toadsorbed Pi at each sampling time. AMARvalue N1 indicates greater ad-sorption of As(V) than Pi at that sampling point whereas a MAR b1shows greater adsorption of Pi than As(V). The MAR values for individ-ual ion adsorptionwere very close to 1 (N0.94) after 20 min and 40 minat pH 4 and pH 8, respectively (Fig. 5). Competitive adsorption with si-multaneous loading resulted in MAR values of 1.18 to 1.35 at both pHvalues throughout the experiments (Fig. 5). For competitive adsorptionwith sequential loading, the calculated MAR values at pH 4 were N1after 20 min; however, MAR values at pH 8 increased from 0.74 tovalues close to 1 between 0 and 24 h (Fig. 5). Calculated MAR valuesfor the ratio of desorbed As(V) to desorbed Pi at pH 4 and 8 (Fig. 5)were significantly lower than 1, indicating greater competitive desorp-tion of Pi by added As(V) than vice versa.

The MAR values for sequential loadings were also calculated usingthe ratio of the remaining adsorbed amount of thepreviously equilibrat-ed oxyanion to the adsorbed amount of the sequentially loaded co-oxyanion (Fig. S5). With this calculation, the MAR values for sequentialaddition of As(V) to the Pi-equilibrated system gradually increased fromb0.5 to ~1 at both pH 4 and pH 8. In contrast, MAR values decreasedfrom 2.27 to 1.37 at pH 4, and from 1.55 to 1.15 at pH 8, after the se-quential addition of Pi to the As(V)-equilibrated system. Apparently,greater MAR values at pH 4 than at pH 8 for As(V)-before-Pi systemsare observed. However, almost identical MAR values are obtained atpH 4 and 8 when As(V) was loaded in sequence after Pi.

3.8. Competitive adsorption/desorption

The competitive adsorption of As(V) or Pi after sequential loading re-sults in desorption of the previously equilibrated oxyanion (Fig. 4).Withrespect to the adsorbed amount of the previously equilibrated oxyanionat 48 h of reaction time, sequentially added As(V) caused greater de-sorption of Pi than vice versa (Fig. 4). The sequentially added co-oxyanion resulted in 11–14% desorption of the previously equilibratedAs(V), and 26–28% of the previously equilibrated Pi. Previously, Liuet al. (2001), O'Reilly et al. (2001), and Carabante et al. (2010) reporteda similar finding. The greater As(V)-induced desorption of Pi from ferri-hydrite is consistentwith the observed greater adsorption preference ofAs(V) after simultaneous loading of both oxyanions, aswell as the great-er amount of non-exchangeable As(V) which remained adsorbed afterthe sequential loading of Pi.

37G. Neupane et al. / Chemical Geology 368 (2014) 31–38

3.9. Potential mechanisms for As(V) adsorption preference on ferrihydrite

The competitive adsorption/desorption data (Figs. 2, 3,and 4) andevolution of MAR values (Figs. 5 and S5) illustrate the apparent prefer-ential adsorption of As(V) by ferrihydrite in competitive adsorption set-tings. The greater adsorption preference of As(V) over Pi on Fe-oxidesalso has been reported previously by several researchers (e.g., Gao andMucci, 2001; Liu et al., 2001; Violante and Pigna, 2002). Other adsorp-tion experiments conducted by Neupane (2012) as a function of so-lution pH (adsorption envelopes) and adsorbent surface coverage(adsorption isotherms) also indicated that adsorption preference offerrihydrite for As(V) over Pi is dependent upon both of these factors.In general, ferrihydrite adsorption preference for As(V) increasedwith increasing surface coverage (in competitive experiments), anddecreased gradually with increasing pH (Neupane, 2012).

The greater adsorption of As(V) over Pi on ferrihydrite is an interest-ing observation because both ortho-arsenic (H3AsO4) and ortho-phosphoric (H3PO4) acids have similar aqueous species and dissociationconstants, and As(V) and Pi both have tetrahedral structures with simi-lar thermochemical radii of 2.48 Å and 2.38 Å, respectively (da Silva andWilliams, 2001). Furthermore, for both oxyanions themono-protonatedaqueous species at neutral pH have similar (−0.895 for HAsO4

2 − and−0.952 for HPO4

2 −) partial negative charges on O in As(V) and Pi,with a partial charge of +1.125 and +1.263 on the central atom ofAs and P, respectively (Kish and Viola, 1999). Even if the small differ-ences in thermochemical radii and partial negative charges on O inAs(V) and Pi influence the adsorption preference between thesetwo oxyanions, Pi should be favored because of its slightly smallerthermochemical radius and greater partial negative charge on O.However, the competitive adsorption experiment shows the oppo-site, with an adsorption preference of As(V) over Pi on ferrihydrite.

The greater adsorption preference of ferrihydrite for As(V) over Pi inthe competitive setting could be attributable to other factors, such as agreater affinity for the formation of Pi-Na+ ion pairs than As(V)–Na+

ion pairs (Gao and Mucci, 2001), the presence of greater specific sorp-tion sites on ferrihydrite for As(V) than for Pi, or because As(V) makesstronger sorption bonds than Pi, thus leading to a greater adsorption.However, it is unlikely that variation in ion-pair formation with theelectrolyte is the sole reason for the preferential adsorption of As(V)on ferrihydrite because Violante and Pigna (2002) not only reportedgreater adsorption preference for As(V) on Fe-oxides (or Fe-richclays), but also reported a greater adsorption preference for Pi on Al-oxides (or Al-rich clays) than for As(V). Although Violante and Pigna(2002) used a KCl solution as the electrolyte, their results for competi-tive adsorption of As(V) and Pi on Fe-oxides and Fe-rich clays are consis-tent with the results of this study.

The presence of specific sorption sites for As(V) on ferrihydrite sur-faces that exclude Pi could also explain the greater adsorption prefer-ence for As(V). If these sites exist, then greater adsorption of As(V)than Pi would also be expectedwith higher surface coverage for individ-ual ion adsorption experiments. However, during individual ion adsorp-tion experiments, adsorption of As(V) and Pi by ferrihydrite at highsurface coverages was very similar (Neupane, 2012). Furthermore, thecombined total adsorption of As(V) and Pi in competitive adsorption ex-periments as a function of surface coverage (competitive adsorption iso-therms) was comparable to the amount of adsorption observed in theindividual ion setting for either As(V) or Pi (Neupane, 2012). Theseobservations indicate the ferrihydrite surface lacks significant As(V)-specific sorption sites.

Another possible reason for preferred uptake of As(V) during com-petitive adsorption experiments could be caused by its ability to makestronger sorption bonds with ferrihydrite than Pi. This is very similarto the assumption that ferrihydrite has As(V)-specific sites which Pican occupy in individual ion adsorption experiments, but Pi is not thepreferred species if As(V) is also present in the system. A greateramount of Pi than As(V) in solution would be required to compete for

the As(V)-preferred sites in competitive adsorption experiments. Al-though this hypothesis was not tested in this study, Carabante et al.(2010) previously reported that five times more aqueous Pi thanAs(V) was required to desorb an equivalent amount of the previouslyadsorbed As(V) on ferrihydrite. Therefore, of the three alternativemechanismsdiscussed, it seemsmost likely that As(V) sorbswith stron-ger bonds on ferrihydrite than Pi, resulting in its preferential adsorptionin a competitive adsorption setting.

3.10. Environmental implications

Interactions of Pi and As(V)with soil/sediment constituents (e.g., Fe-hydroxides) in near surface environments could have several environ-mental implications. Although Pi has no such stringent environmentalregulations as for As, the excess input of Pi to the environment could di-rectly deteriorate the quality of surface waters by eutrophication(Correll, 1998). Moreover, despite less desorption of adsorbed As(V)on ferrihydrite by Pi than vice versa, the increased mobility of As(V)due to excessive use of Pi-fertilizer could have a severe environmentalimpact because of its toxicity. However, Pi interaction with As-contaminated geomedia could also offer an opportunity to interceptthe mobilized As(V) by co-precipitation in Pi-phases. Recently,Neupane and Donahoe (2013) showed in situ fixation of As (presentas As(V)) in soil with precipitation of apatite-like phases by simulta-neous delivery of both Pi and Ca. Therefore, understanding the As(V)and Pi interactions with soil/sediment constituents could help evaluatethe risk and develop a new remediation method to safeguard the near-surface environment.

4. Conclusions

Individual ion as well as competitive adsorption/desorption experi-ments utilizing As(V) and Pi indicated early rapid adsorption of bothoxyanions at the ferrihydrite–water interface, followed bymuch sloweradsorption kinetics. The initial rapid adsorption/desorption is attributedto formation/displacement of surface complexes, whereas the latersluggish adsorption/desorption is either related to saturation of specificadsorption sites or tomigration of oxyanions to/from the particles' inte-rior. Both oxyanions showed limited ability to desorb the previouslyequilibrated co-oxyanion from the ferrihydrite. The competitive desorp-tion of the previously equilibrated oxyanion (e.g., As(V)) was less thanthe competitive adsorption of the sequentially loaded co-oxyanion(e.g., Pi). Although the addition of As(V) induced greater desorptionof Pi than vice versa, Pi also showed a limited ability to desorb thepreviously equilibrated As(V) from ferrihydrite. It is therefore possi-ble that input of Pi-rich fertilizers could mobilize As from contami-nated soils into natural waters. Furthermore, it is likely that the Pi-mobilized As(V) could be fixed in geomedia by co-precipitating itin Pi-phases with the simultaneous delivery of Ca.

The experimental adsorption/desorption kinetics showed goodagreement with pseudo-second order, Elovich, and power-functionequations. The As K-edge EXAFS analysis indicated the presence oftwo types of inner-sphere adsorption complexes: 2E and 2C. The Fe co-ordination numbers (CN) increasedwith increasing time and decreasedwith addition of Pi into the system. Particularly, a higher proportion ofthe Fe CNs associated with 2E As(V) surface complexes decreased afterthe addition of Pi than CN associatedwith 2C. Despite the similar ferrihy-drite adsorption affinities observed during individual ion adsorption ex-periments, As(V) was the preferred species in competitive adsorptionsettings, indicating a stronger bonding of As(V) than Pi on the ferrihy-drite surface.

Acknowledgements

We would like to thank Dr. S. Bhattacharyya and E.Y. Graham fortheir assistance with ICP-OES analysis of our samples. The use of the

38 G. Neupane et al. / Chemical Geology 368 (2014) 31–38

Advanced Photon Source, an Office of Science User Facility operated forthe U.S. Department of Energy (DOE) Office of Science by ArgonneNational Laboratory, was supported by the U.S. DOE under ContractNo. DE-AC02-06CH11357. We greatly appreciate beamline scientistT. Bolin's help during XAS data collection.

Appendix A. Supplementary data

Supplementary data to this article can be found online at http://dx.doi.org/10.1016/j.chemgeo.2013.12.020.

References

Aharoni, C., Sparks, D.L., 1991. Kinetics of soil chemical reactions — a theoretical treat-ment. In: Sparks, D.L., Suarez, D.L. (Eds.), Rate of soil chemical processes. SSSA Spec.Publ., 27, pp. 1–18.

Bhattacharyya, S., Donahoe, R.J., Patel, D., 2009. Experimental study of chemical treatmentof coal fly ash to reduce the mobility of priority trace elements. Fuel 88, 1173–1184.

Carabante, I., Grahn, M., Holmgren, A., Hedlund, J., 2010. In situ ATR-FTIR studies on thecompetitive adsorption of arsenate and phosphate on ferrihydrite. J. Colloid InterfaceSci. 351, 523–531.

Chen, N., Jiang, D.T., Cutler, J., Kotzer, T., Jia, Y.F., Demopoulos, G.P., Rowson, J.W., 2009.Structural characterization of poorly-crystalline scorodite, iron(III)-arsenate co-precipitates and uranium mill neutralized raffinate solids using X-ray absorptionfine structure spectroscopy. Geochim. Cosmochim. Acta 73, 3260–3276.

Chien, S.H., Clayton, W.R., 1980. Application of Elovich equation to the kinetics of phos-phate release and sorption in soils. Soil Sci. Soc. Am. J. 44, 265–268.

Correll, D.L., 1998. Role of phosphorus in the eutrophication of receiving waters: a review.J. Environ. Qual. 27, 261–266.

Cullen, W.R., Reimer, K.J., 1989. Arsenic speciation in the environment. Chem. Rev. 89,713–764.

da Silva, J.J.R.F., Williams, R.J.P., 2001. The Biological Chemistry of the Elements: The Inor-ganic Chemistry of Life, 2nd ed. Oxford University Press, New York USA (600 pp.).

Del Nero, M., Galindo, C., Barillon, R., Halter, E., Madé, B., 2010. Surface reactivity of α-Al2O3 and mechanisms of phosphate sorption: in situ ATR-FTIR spectroscopy and ζpotential studies. Soil Sci. Soc. Am. J. 342, 437–444.

Foster, A.L., Brown Jr., G.E., Tingle, T.N., Parks, G.A., 1998. Quantitative arsenic speciation inmine tailings using X-ray absorption spectroscopy. Am. Mineral. 83, 553–568.

Fuller, C.C., Davis, J.A., Waychunas, G.A., 1993. Surface chemistry of ferrihydrite: part 2. Ki-netics of arsenate adsorption and coprecipitation. Geochim. Cosmochim. Acta 57,2271–2282.

Gao, Y., Mucci, A., 2001. Acid base reactions, phosphate and arsenate complexation, andtheir competitive adsorption at the surface of goethite in 0.7 M NaCl solution.Geochim. Cosmochim. Acta 65, 2361–2378.

Harvey, O.R., Rhue, R.D., 2008. Kinetics and energetics of phosphate sorption in a multi-component Al(III)–Fe(III) hydr(oxide) sorbent system. Soil Sci. Soc. Am. J. 322,384–393.

Hingston, F.J., Posner, A.M., Quirk, J.P., 1971. Competitive adsorption of negatively chargedligands on oxide surfaces. Discuss. Faraday Soc. 52, 334–342.

Ho, Y.S., McKay, G., 1999. Pseudo-second order model for sorption processes. ProcessBiochem. 34, 451–465.

Kish, M.M., Viola, R.E., 1999. Oxyanion specificity of aspartate-β-semialdehyde dehydro-genase. Inorg. Chem. 38, 818–820.

Kitahama, K., Kiriyama, R., Baba, Y., 1975. Refinement of the crystal structure of scorodite.Acta Crystallogr. B31, 322–324.

Langmuir, D., Mahoney, J., Rowson, J., 2006. Solubility products of amorphous ferric arse-nate and crystalline scorodite (FeAsO4·2H2O) and their application to arsenic behav-ior in buried mine tailings. Geochim. Cosmochim. Acta 70, 2942–2956.

Liu, F., De Cristofaro, A., Violante, A., 2001. Effect of pH, phosphate and oxalate on the ad-sorption/desorption of arsenate on/from goethite. Soil Sci. 166, 197–208.

Low, M.J.D., 1960. Kinetics of chemisorption of gases on solids. Chem. Rev. 60, 267–312.Luengo, C., Brigante, M., Avena, M., 2007. Adsorption kinetics of phosphate and arsenate

on goethite. A comparative study. J. Colloid Interface Sci. 311, 354–360.

Neupane, G., 2012. Application of Phosphate and Surfactant-modified Zeolite for Remedi-ation of Trace Elements in Soil and Coal Fly Ash. University of Alabama (PhD disser-tation, 256 pp.).

Neupane, G., Donahoe, R.J., 2013. Calcium-phosphate treatment of contaminated soil forarsenic immobilization. Appl. Geochem. 28, 145–154.

Neupane, G., Donahoe, R.J., Qi, Y., 2010. Fate of arsenic, iron, and lanthanum in ferrous sul-fate and ferrous sulfate-lanthanum chloride treated arsenic-contaminated soils. G-073, in K.A. Fields and G.B. Wickramanayake (Chairs), Remediation of Chlorinatedand Recalcitrant Compounds—2010. Seventh International Conference on Remedia-tion of Chlorinated and Recalcitrant Compounds (Monterey, CA; May 2010). ISBN978-0-9819730-2-9, Battelle Memorial Institute, Columbus, OH, www.battelle.org/chlorcon.

Newville, M., 2001. IFEFFIT: interactive XAFS analysis and FEFF fitting. J. SynchrotronRadiat. 8, 322–324.

Newville, M., Livins, P., Yacoby, Y., Rehr, J., Stern, E., 1993. Near-edge X-ray-absorptionfine-structure of Pb — a comparison of theory and experiment. Phys. Rev. B 47,14126–14131.

O'Reilly, S.E., Strawn, D.G., Sparks, D.L., 2001. Residence time effects on arsenate adsorp-tion/desorption mechanisms on goethite. Soil Sci. Soc. Am. J. 65, 67–77.

Parkhurst, D.L., Appelo, C.A.J., 1999. User's guide to PHREEQC (Version 2)—A computerprogram for speciation, batch-reaction, one-dimensional transport, and inverse geo-chemical calculations. U.S. Geol. Surv. Water Resour. Invest. Rep. 99–4259 (326 pp.).

Ravel, B., 2001. ATOMS: crystallography for the X-ray absorption spectroscopist.J. Synchrotron Rad. 8, 314–316.

Ravel, B., Newville, M., 2005. Athena, Artemis, Hephaestus data analysis for X-ray absorp-tion spectroscopy using IFEFFIT. J. Synchrotron Rad. 12, 537–541.

Raven, K.P., Jain, A., Loeppert, R.H., 1998. Arsenite and arsenate adsorption on ferrihydrite:kinetics, equilibrium, and adsorption envelopes. Environ. Sci. Technol. 32, 344–349.

Rose, J., Flank, A.M., Masion, A., Bottero, J.Y., Elmerich, P., 1997. Nucleation and growthmechanisms of Fe oxyhydroxide in the presence of PO4 ions. 2. P K-edge EXAFSstudy. Langmuir 13, 1827–1834.

Sherman, D.M., Randall, S.R., 2003. Surface complexation of arsenic(V) to iron(III) (hydr)oxides: structural mechanism from ab initio molecular geometries and EXAFS spec-troscopy. Geochim. Cosmochim. Acta 67, 4223–4230.

Smedley, P.L., Kinniburgh, D.G., 2002. A review of the source, behaviour and distributionof arsenic in natural waters. Appl. Geochem. 17, 517–568.

Sparks, D.L., 1989. Kinetics of Soil Chemical Processes. Academic Press, New York(210 pp.).

Strauss, R., Brümmer, G.W., Barrow, N.J., 1997. Effects of crystallinity of goethite: II. Ratesof sorption and desorption of phosphate. Eur. J. Soil Sci. 48, 101–114.

Tejedor-Tejedor, M.I., Anderson, M.A., 1990. Protonation of phosphate on the surface ofgoethite as studied by CIR-FTIR and electrophoretic mobility. Langmuir 6, 602–611.

Violante, A., Pigna, M., 2002. Competitive sorption of arsenate and phosphate on differentclay minerals and soils. Soil Sci. Soc. Am. J. 66, 1788–1796.

Waychunas, G.A., Rea, B.A., Fuller, C.C., Davis, J.A., 1993. Surface chemistry of ferrihvdrite:Part 1. EXAFS studies of the geometry of coprecipitated and adsorbed arsenate.Geochim. Cosmochim. Acta 57, 2251–2269.

Waychunas, G., Trainor, T., Eng, P., Catalano, J., Brown, G., Davis, J., Rogers, J., Bargar, J.,2005. Surface complexation studied via combined grazing-incidence EXAFS and sur-face diffraction: arsenate on hematite (0001) and (10–12). Anal. Bioanal. Chem. 383,12–27.

Willett, I.R., Chartres, C.J., Nguyen, T.T., 1988. Migration of phosphate into aggregated par-ticles of ferrihydrite. J. Soil Sci. 39, 275–282.

Withers, P.J.A., Clay, S.D., Breeze, V.G., 1999. Phosphorus transfer in runoff following appli-cation of fertilizer, manure, and sewage sludge. J. Environ. Qual. 30, 180–188.

Yang, L., Donahoe, R.J., 2007. The form, distribution and mobility of arsenic in soils con-taminated by arsenic trioxide, at sites in southeast USA. Appl. Geochem. 22, 320–341.

Yang, L., Donahoe, R.J., Redwine, J.C., 2007. In situ chemical fixation of arsenic-contaminated soils: an experimental study. Sci. Total Environ. 387, 28–41.

Zabinsky, S.I., Rehr, J.J., Ankudinov, A., Albers, R.C., Eller, M.J., 1995.Multiple-scattering cal-culations of X-ray-absorption spectra. Phys. Rev. B 52, 2995–3009.

Zhang, H., Selim, H.M., 2007. Modeling competitive arsenate–phosphate retention andtransport in soils: a multi-component multi-reaction approach. Soil Sci. Soc. Am. J.71, 1267–1277.

Zhang, J., Stanforth, R., 2005. Slow adsorption reaction between arsenic species and goe-thite (r-FeOOH): diffusion or heterogeneous surface reaction control. Langmuir 21,2895–2901.