Effect of major anions on arsenate desorption from ferrihydrite-bearing natural samples

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Applied Geochemistry 23 (2008) 1451–1466

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AppliedGeochemistry

Effect of major anions on arsenate desorption fromferrihydrite-bearing natural samples

Franco Frau *, Riccardo Biddau, Luca Fanfani

Department of Earth Sciences, University of Cagliari, Via Trentino 51, I-09127 Cagliari, Italy

Received 26 June 2007; accepted 12 January 2008Editorial handling by D. Polya

Available online 6 February 2008

Abstract

The influence of background electrolytes (Na2HPO4 � 2H2O, NaHCO3, Na2SO4, NaNO3 and NaCl) on arsenate(As(V)) desorption from 3 environmental samples (a tailings sample, a stream-bed sediment and a top soil) containing fer-rihydrite as the main As-bearing phase has been studied by means of kinetic batch experiments and geochemical simula-tions. The experimental results indicate that As(V) release increases greatly in the presence of dissolved phosphate andcarbonate species. Similarly to PO3þ

4 , a strong surface interaction of inner-sphere type between ferrihydrite and aqueouscarbonate species is suggested. Nitrate and Cl� have negligible effects on the As(V) desorption reaction, whereas SO2�

4

exhibits intermediate behavior depending on its dissolved concentration that probably influences the type of surface com-plex (i.e. outer-sphere or inner-sphere). The process of As(V) release follows the first-order rate equation of Lagergrenmodified for desorption; most values of the desorption rate constant kdes are in the range of 0.0012–0.0030 min�1. Mod-eling of the desorption experiments with PHREEQC, with ferrihydrite as the main As-bearing phase, indicates that theinfluence of pH is notably less important than the displacement action of carbonate species in determining the amountof As(V) released to solution. Simulation of As(V) desorption totally fails when the carbonate surface complexes areexcluded from the model. In the NaHCO3 experiments with the tailings sample the best match between observed and cal-culated data is obtained also including dissolution of scorodite and arsenopyrite in the model. Moreover, modeling hasstressed the poor quality of the adsorption constants for sulfate species that leads to strong overestimation of As(V)desorption at pH 4 and underestimation at pH 7.5. Although the findings of this study are consistent with the resultsof recent studies from other authors, they cannot be generalized or directly applied to natural systems. However, environ-mental implications concerning As mobility, as well as possible application in various fields (e.g. irrigation agriculture, soildecontamination, water treatment and mine site remediation), might be derived from these findings.� 2008 Elsevier Ltd. All rights reserved.

1. Introduction

Arsenic is known as one of the most toxic ele-ments, and is therefore of great interest in environ-

0883-2927/$ - see front matter � 2008 Elsevier Ltd. All rights reserveddoi:10.1016/j.apgeochem.2008.01.006

* Corresponding author. Fax: +39 070 282236.E-mail address: [email protected] (F. Frau).

mental studies. Arsenate (As(V)) and arsenite(As(III)) are the two main forms of As commonlyfound in soils, sediments and waters. Arsenate gen-erally predominates under oxidizing conditions,whereas As(III) occurs when conditions becomesufficiently reducing (Smedley and Kinniburgh,2002).

.

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1452 F. Frau et al. / Applied Geochemistry 23 (2008) 1451–1466

Sorption is the main factor controlling mobilityof As in natural systems. Many studies have beendevoted to As adsorption on a wide variety ofphases (Arai et al., 2001; Chiu and Hering, 2000;Manning and Goldberg, 1997; Pierce and Moore,1982). In particular, adsorption and desorptionreactions between As and hydrous ferric oxides(e.g. ferrihydrite) are important As-controllingmechanisms because hydrous ferric oxides are wide-spread in the environment and strongly adsorb Asat acidic and near-neutral pH conditions (Dixitand Hering, 2003; Pierce and Moore, 1982; Ravenet al., 1998). Hydrous ferric oxides are also impor-tant for the removal of As from a variety of wastewaters; in fact, precipitation of hydrous ferric oxidesis presently the most utilized disposal practice forAs in the metallurgical and mining industries (Riv-eros et al., 2001).

The reaction of As with hydrous ferric oxideshas been studied by EXAFS and FTIR (Sun andDoner, 1996; Waychunas et al., 1993). Accordingto these studies, As(V) is strongly adsorbed ontothe surfaces of hydrous ferric oxides predominantlyby forming inner-sphere bidentate complexes.Moreover, an adsorption maximum at pH 4–5, fol-lowed by a gradual decrease in adsorption withincreasing pH, has been emphasized (Pierce andMoore, 1982).

The effectiveness of adsorption/desorption reac-tions is determined by a number of variables,including pH, temperature, ionic strength, presenceof competing ions and structural changes in solidphases at the atomic level (Smith, 1999). Focusingon competing ions, it is well known that PO3�

4

and As(V) have similar geochemical behaviors,and both anions compete for adsorption sites onthe surfaces of hydrous ferric oxides (Anteloet al., 2005). The competitive role of SO2�

4 , NO�3and Cl� has been less well investigated, but theseanions generally appear much less effective in pro-moting As mobility (Goh and Lim, 2005; Rietraet al., 2000) or are considered ineffective as poten-tial competitors. Bicarbonate is often a majoranion in As-affected natural waters (Smedley andKinniburgh, 2002) and for this reason some studieshave recently focused on it as a potential competi-tor or extractor of As, finding either a negligible tomoderate effect (Meng et al., 2002; Wilkie and Her-ing, 1996) or a moderate to notable effect (Anawaret al., 2004; Appelo et al., 2002; Goh and Lim,2005), as well as equivocal results (Arai et al.,2004).

The aim of this study is to investigate the effectof dissolved PO3�

4 , CO2�3 , SO2�

4 , NO�3 and Cl� onAs(V) desorption processes in environmental mate-rials from an old mine area, containing ferrihy-drite (HFO) as the main As-bearing phase, bymeans of kinetic batch experiments and geochem-ical simulations. In other terms, the main goal/challenge of this study is to verify if the adsorp-tion/desorption processes involving As(V) andpotential competitors at the HFO surface may bemacroscopically recognized and modeled by usingnatural complex heterogeneous samples insteadof a single pure phase (i.e. ferrihydrite) synthesizedin the laboratory. As known, the use of naturalcomplex heterogeneous samples is a source ofnotable difficulties both in experiment set-up anddata interpretation. However, these drawbackscan be partly overcome or at least minimizedwhen well-characterized samples are used. More-over, verification with natural samples of experi-mental results obtained in controlled simplifiedartificial systems with synthetic phases representsa challenging task needed for an in-depth under-standing of the natural systems.

Although natural solids were used, the kineticbatch experiments performed in this study did notdirectly intend to simulate the natural conditions.In fact, uncommonly high concentrations (0.01and 0.1 M) of some anionic species (especiallyPO3�

4 and CO2�3 Þ were used according to the fre-

quent practice in many laboratory experiments thataim to verify the occurrence of a particular process/mechanism, thereby minimizing the possible inter-ference of a secondary parallel process/mechanism.This is particularly true in this study where the con-centrations of the competing anions in the batchsolutions had to exceed those deriving from possi-ble dissolution of soluble salts contained in thedried solid samples. Therefore, the results obtainedin this study cannot be directly applied to naturalsystems where lower concentrations of the consid-ered anionic species generally occur, but they pro-vide a contribution to the debated questionconcerning the role of potential competitors infavoring the release of As from sediments towaters.

Moreover, a simple adsorption model (i.e. Gen-eralized Two-Layer Model by Dzombak and Morel,1990) was used to simulate the experimental results,since more refined models were considered less suit-able for batch experiments with natural complexheterogeneous solid samples.

F. Frau et al. / Applied Geochemistry 23 (2008) 1451–1466 1453

2. Experimental

2.1. Materials

The samples used in the batch experiments arefrom the inactive mine area of Baccu Locci(south-eastern Sardinia, Italy), exploited for Asand Pb for about a century (1873–1965). The maincarrier of As in the environment is the Baccu Loccistream, whose waters (pH 7–8, Eh = 0.4 � 0.5 V,Ca–Mg–SO4 to Ca–Mg–HCO3 composition, Asup to 1 mg/L) continuously interact with the tailings(containing up to 24,000 mg/kg of As) which weredischarged from the flotation plant directly intothe stream-bed and redistributed downstream forseveral kilometers as far as the alluvial coastal plain.

Various sampling campaigns for waters and sol-ids have been carried out at Baccu Locci over thelast years, and numerous samples of waste rocks,flotation tailings, stream-bed sediments and soilshave been previously studied and characterized byusing various analytical techniques (XRD,WDXRF, SEM/EDX, TEM/EDX, XPS/AES andchemical extractions) (Frau and Ardau, 2003,2004; Frau et al., 2005).

Three samples were chosen for this study: a tail-ings sample (T), a stream-bed sediment (S) and atop soil (TS). Samples T and S were collected alongthe middle Baccu Locci stream course, downstreamof the flotation plant; sample TS is a top soil (0–20 cm) collected in the alluvial coastal plain of Quir-ra located about 10 km downstream of the minedarea (Frau and Ardau, 2003). The main mineralog-ical composition of the samples, as determined byXRD (Frau and Ardau, 2004), consists of quartz,K-feldspar, clinochlore, white mica and biotite.Arsenical two-line ferrihydrite (HFO) with mostFe/As molar ratios in the range of 1.6–3.2 (Frauet al., 2005) is by far the main phase containingAs. Ferrihydrite generally occurs as a coating on sil-icate grains wherein As, mainly as As(V), is presentas an adsorbed or co-precipitated species. Other As-bearing phases, definitely less abundant than HFOin the samples considered, are scorodite and arseno-pyrite, as determined by SEM/EDX and TEM/EDX (Frau and Ardau, 2004; Frau et al., 2005).Scorodite and arsenopyrite will be included in thegeochemical simulations along with HFO, whereasa potential As-bearing mineral such as calcite (Alex-andratos et al., 2007) was not included becauseSEM/EDX and TEM/EDX analyses never showedAs bound to calcite in the samples considered (Frau

and Ardau, 2004; Frau et al., 2005). Moreover,chemical extractions showed that the amount ofAs extracted in the carbonatic fraction (referableto calcite) was generally below 1% of total extractedAs, whereas the amount of As extracted in the mod-erately reducible fraction (referable to HFO) was75% of total extracted As on average (Frau andArdau, 2004). The top soil sample TS also containsgoethite that, however, will be considered togetherwith HFO as a whole in this study. Total As andFe concentrations and their solid speciation, as esti-mated in previous studies by chemical extractions(Dadea, 2003; Frau and Ardau, 2004), are reportedin Table 1. The sequential selective extraction proce-dure after Londesborough et al. (1998) was appliedto the T and S samples (Frau and Ardau, 2004). Itconsisted of 6 steps related to 6 fractions (compo-nents extracted): (1) exchangeable/soluble(exchangeable ions and elements in soluble phases);(2) carbonatic (elements bound to carbonate miner-als); (3) easily reducible (elements bound to oxides/hydroxides of Mn); (4) moderately reducible (ele-ments bound to amorphous or poorly crystallineoxides/hydroxides of Fe); (5) sulfidic/organic (ele-ments bound to sulfide minerals and organic mat-ter); (6) residual (elements bound to moreresistant, non-silicate minerals). The main results,as shown in Table 1, indicated that most As wasbound to amorphous or poorly crystalline oxides/hydroxides of Fe (HFO), secondly to sparingly sol-uble arsenates (scorodite) and sulfides (arsenopy-rite). Negligible or no organic matter was presentin the solids. The remaining fractions accountedfor less than 3% of As extracted. The sequentialselective extraction procedure after Li et al. (1995)was applied to the TS sample (Dadea, 2003). It con-sisted of only 5 steps because the two fractions refer-able to the oxides/hydroxides of Mn and Fe weremerged into one fraction. Results showed As to betotally bound to this fraction (see Table 1) thatcan be completely referable to the oxides/hydrox-ides of Fe since the Mn content in the soil sampleis very low. The fact that the sum of the As concen-trations bound to HFO, scorodite and arsenopyrite(as determined by chemical extraction) is not equalto the total As content (as determined by WDXRF)in the T and S samples is mainly due to loss of sam-ple during the extraction procedure. Moreover,since the silicate minerals were not attacked by thereagents used in the chemical extraction (Frau andArdau, 2004), this might have prevented As-bearingminerals embedded in silicate grains from being

Table 1Some physical–chemical features of the solid samples used in the batch desorption experiments, with estimated speciation for As and Fe,and percentage of As bound to ferrihydrite desorbed after 24 h from the samples in the presence of different background electrolytes andpure water

Sample T (Tailings) S (Stream-bed sediment) TS (Top soil)

Grain-size (70% fraction) (lm) 250–500 500–1000 63–125Total As (mg/kg) 19,400 2300 900Total Fe (mg/kg) 52,400 27,600 39,870As bound to HFO (mg/kg) 13,600 (70) 1220 (53) 900 (100)As bound to Scorodite (mg/kg) 1600 (8.2) 620 (27) –As bound to Arsenopyrite (mg/kg) 740 (3.8) 250 (11) –Fe bound to HFO (mg/kg) 15,720 (30) 4400 (16) 3190 (8)Fe/As in HFO (molar ratio) 1.5 4.8 4.7HFO in batch solutions (mg/L) 125 35 25

AsHFO desorbed after 24 h

PO4 (0.01–0.1 M) (%) 8.6–7.0 11.4–8.2 14.0–14.3CO3 (0.01–0.1 M) (%) 15.3–23.0 6.4–9.4 11.5–20.0SO4 (0.01–0.1 M) (%) 0.36–0.69 4.1–4.9 8.0–7.3NO3 (0.01–0.1 M) (%) 0.15–0.12 3.7–2.5 6.2–5.8Cl (0.01–0.1 M) (%) 0.13–0.16 2.3–2.7 5.7–5.6H2O (%) 0.11 2.1 5.1

HFO stands for ferrihydrite (and also goethite for the TS sample) and AsHFO for As bound to ferrihydrite. HFO is expressed asFe2O3 � H2O. Percentages on total As and total Fe are given in brackets.


solubilized. However, this fraction of As is to beconsidered ‘‘insoluble” and, thus, not involved inthe batch experiments of this study.

Speciation of water-leached As by IC–ICPMS(Ion Chromatography–Inductively Coupled PlasmaMass Spectroscopy) showed that only 2–3% oftotal leached As was released as As(III) from thetailings sample T after 24 h reaction time, whereasthe stream-bed sediment sample S released As(V)solely (Frau and Ardau, 2003). No organic As spe-cies were detected. Only 2% of total water-leachedAs was released as As(III) from the soil sampleTS after 24 h reaction time, as measured by polar-ography (Dadea, 2003). On the basis of theseresults, desorption reactions investigated in thisstudy are considered to involve inorganic As(V)species only.

2.2. Methods

All salts used (Na2HPO4 � 2H2O, NaHCO3,Na2SO4, NaNO3 and NaCl) in the batch experi-ments were high grade analytical reagents (CarloErba�) with As content less than 0.00005%. Milli-Q� water (Millipore, 18 MX cm) was always usedto prepare the solutions. Glassware and plasticwarewere accurately and properly cleaned prior to use.Solid samples were air dried, homogenized by quar-tering and sieved to 2 mm.

Batch desorption experiments were conducted at0.01 and 0.1 M concentrations of the followinganions: PO3�

4 , CO2�3 , SO2�

4 , NO�3 and Cl�. Experi-ments with pure water solutions were used as‘‘blanks”. A solid/solution ratio of 1:200 (g:mL)was always used.

All experiments were performed at room temper-ature and the solutions were continuously stirred ata low rate with a teflon coated magnetic bar. Ideallyall experiments should be run with a pH-Stat, butsome preliminary tests performed on the tailingssample T in Cl� solution demonstrated that it wasimpossible to raise the natural solution pH (pH4.0) in spite of the adding of up to 50 mL of 1 MNaOH. This probably happens because the tailingssample is a ‘‘variable surface charge” solid and, assuch, tends to release H+ ions from the hydroxyl-ated surface, with consequent neutralization ofOH� added (Smith, 1999). For this reason solutionpH was left to drift freely to the value generated bythe interaction with the solids. A stable pH valuewas quickly reached a few minutes after the begin-ning of the experiments. This rapidity is an indica-tion that solution pH was mainly due to the buffercapacity of the electrolyte or/and the formation offunctional groups on the solid surface as a conse-quence of the dissociative adsorption of water mol-ecules, even if the dissolution of very solubleminerals (e.g. metal sulfates) may have contributed


too, whereas the dissolution of less soluble minerals,such as arsenopyrite and scorodite, might affectsolution pH and the amount of As released onlyfor more prolonged reaction times. The tailingssample T instantly generated acidic conditions (pH4.0) in the presence of pure water, Cl�, NO�3 andSO2�

4 solutions, whereas the stream-bed sedimentsample S and the top soil sample TS produced nearneutral conditions (pH 7.5). Slightly more alkalineconditions (pH 8.0) were generated by the 3 samplesin the presence of PO3�

4 and CO2�3 solutions. The

redox potential values measured were always oxidiz-ing (Eh = 0.44 � 0.54 V). Both pH and Eh werechecked throughout the experiment (pH-meter andelectrodes Orion�); no significant variations wereobserved (pH ± 0.2, Eh ± 0.05 V).

A reaction time of 24 h was used, and at appro-priate time intervals a 5 mL aliquot of solutionwas sampled from the reaction beaker, filteredthrough a membrane filter (0.45 lm pore size), acid-ified with suprapure HNO3 and stored at 4 �C untilanalysis. The duration of each experiment was lim-ited to 24 h in order to minimize as much as possiblethe dissolution of As-bearing minerals (e.g. scoro-dite, arsenopyrite) with respect to the investigateddesorption process, even if this may have sometimesprevented the achievement of equilibrium As con-centrations in the batch solutions.

Arsenic was determined by ARL 3520 ICP–AES(Inductively Coupled Plasma–Atomic EmissionSpectroscopy). Analytical precision was estimatedat ±5%, or better. Arsenic concentrations were cor-rected for the volume effect of aliquots removalaccording to the following equation:

Cncorr ¼ Cn�1

corr þCn

meas � Cn�1meas

� �� ðV in � ðn� 1ÞV remÞV in

;

ð1Þ

where n is the number of the sampling, Ccorr andCmeas are respectively, the corrected and the mea-sured concentrations, Vin is the initial solution vol-

Table 2Chemical composition of the pure water solutions after 24 h of interac

Sample pH Ca Mg Na K Alk Cl SO4 NO3

mg/L

T 4.0 5.38 2.02 1.86 2.66 – 1.24 23.9 <0.1S 7.5 4.02 0.28 0.22 3.24 18.3 2.82 0.20 <0.1TS 7.5 1.41 0.42 2.76 2.60 11.9 3.10 0.88 0.10

Alk is alkalinity expressed as HCO3.

ume in the batch reactor, Vrem is the solutionvolume removed (5 mL) at each sampling.

Some experiments were repeated as a check onreproducibility; an average error of 15% wasobserved at the end of the experiments (after 24 h).

The computer program PHREEQC (version2.12.5.669 with revised As thermodynamic data inthe database WATEQ4F) (Parkhurst and Appelo,1999) was used to model the desorption experiments.

3. Results and discussion

3.1. Desorption experiments

3.1.1. Effect of pH and solution composition on As

release

The chemical composition of pure water solutionafter 24 h of interaction with the 3 solids is reportedin Table 2. The T sample produces a Ca–Mg–SO4

composition, pH 4 and TDS (Total DissolvedSolids) = 45 mg/L; the S sample gives rise to a Ca–HCO3 composition, pH 7.5 and TDS = 30 mg/L;the TS sample produces an HCO3-dominated solu-tion with a more balanced cationic composition(Na–K–Ca), pH 7.5 and TDS = 23 mg/L.

The highest concentrations of metals are releasedfrom the T sample (800 lg/L Fe, 600 lg/L Pb,70 lg/L Zn); they can be mainly referred to dissolu-tion of soluble metal sulfates, such as melanterite,anglesite and bianchite, that have been observedboth in the field and in microscopic scale analyses(Frau and Ardau, 2004). The fast dissolution ofthese metal sulfates deriving from sulfide oxidationmay have contributed to the low pH value (pH 4)generated by the T sample that, however, is thoughtto derive mainly from the dissociative adsorption ofwater molecules onto the HFO surface (see the Sec-tion 2.2). A contribution to SO2�

4 in solution is cer-tainly given by the dissolution of soluble metalsulfates, although the Ca–Mg–SO4 composition(Table 2) clearly indicates that minerals such as gyp-sum and epsomite are the main sources of dissolved

tion with the 3 solid samples T, S and TS

F Si Mn Fe Co Ni Cu Zn Pb As

lg/L

0.07 3.65 8.15 802 1.40 3.60 12.8 68.8 601 770.02 0.50 13.5 133 2.00 1.40 0.30 62.6 28.3 1300.18 – 35.1 395 4.06 4.50 2.87 15.0 27.8 230


SO2�4 . The absence of alkalinity is a consequence of

the low pH, since at pH 4 the aqueous carbonatespeciation is dominated by H2CO3, and is linkedto the low content of calcite in the T sample thatis insufficient to buffer the acidity generated.

The concentrations of the anions released (Table2) are so small with respect to those of the back-ground electrolytes used in the batch experimentsthat their potential interference in the desorptionprocess involving As can be disregarded.

The kinetics experiments of As desorption fromthe 3 environmental samples (T, S and TS), in the

1

10

100

1000

10000

100000

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

PO4 - pH 8.0 CO3 - pH 8.0 SO4 - pH 4.0

NO3 - pH 4.0 Cl - pH 4.0 H2O - pH 4. 0

Sample T0.01 M

10

100

1000

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

PO4 CO3 SO4

NO3 Cl H2O

Sample S0.01 M

pH 7.5-8.0

10

100

1000

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

PO4 CO3 SO4

NO3 Cl H2O

Sample TS0.01 M

pH 7.5-8.0

Fig. 1. Desorption of As from 3 environmental samples (T = tailings; Sand 0.1 M concentrations of different anionic species (PO3�

4 , CO2�3 , SO

presence of 0.01 and 0.1 M concentrations of thedifferent anions, produced the results shown inFig. 1. Table 1 reports the percentages of As boundto ferrihydrite (AsHFO) desorbed after 24 h.

Although data interpretation is complicated bythe partial superposition of the two main factors,i.e. pH and solution composition, the experimentalresults shown in Fig. 1 indicate that the highestAs(V) mobility is obtained by interaction of solidsamples with PO3�

4 and CO2�3 solutions. In fact,

modeling will show that the strong difference indesorption efficiency (up to 3 orders of magnitude)

1

10

100

1000

10000

100000

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

PO4 - pH 8.0 CO3 - pH 8.0 SO4 - pH 4.0

NO3 - pH 4.0 Cl - pH 4.0 H2O - pH 4.0

Sample T0.1 M

1

10

100

1000

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

PO4 CO3 SO4

NO3 Cl H2 O

Sample S0.1 M

pH 7.5-8.0

10

100

1000

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

PO4 CO3 SO4

NO3 Cl H2 O

Sample TS0.1 M

pH 7.5-8.0

= stream-bed sediment; TS = top soil) in the presence of 0.01 M2�4 , NO�3 and Cl�) and pure water.


between PO3�4 =CO2�

3 solutions and other solutionscan be only partly attributed to the pH variable.

With regard to the T sample experiments, theywere conducted at different pH (i.e. pH 4 withCl�, NO�3 , SO2�

4 and pure water solutions, pH 8with PO3�

4 and CO2�3 solutions) and, therefore, the

results cannot be directly compared. However, itwas considered useful to run these experiments atdifferent pH conditions because a comparisonbetween the well known competitive effect of PO3�

4

and the debated competitive effect of CO2�3 was pos-

sible at pH 8, as well as a comparison between apotential competitor such as SO2�

4 and probableineffective competitors such as NO�3 and Cl� at pH4.

The As concentrations released after 24 h fromthe T sample are influenced by the solution compo-sition and pH according to the following sequences:

pH 8; CO3 > PO4; ð2ÞpH 4; SO4 > NO3; Cl; H2O: ð3Þ

Desorption of As from the S and TS samplesoccurred at about constant pH (7.5–8.0) and canbe thus considered as mainly influenced by the solu-tion composition:

0:01 M; PO4 >CO3 > SO4 >NO3; Cl; H2O; ð4Þ0:1 M; CO3 P PO4 > SO4 >NO3; Cl; H2O: ð5Þ

In order to better understand the potential com-petitive role of each background electrolyte, Fig. 2contrasts the different behavior of the same samplesas a function of the anion. The As concentrationsreleased in the 0.01 M and 0.1 M CO2�

3 and PO3�4

solutions decrease according to the following sam-ple order: T� TS P S. In contrast, the NO�3 , Cl�

and pure water solutions desorb decreasing Asamounts from the 3 samples as follows: TS > S > T.An intermediate behavior is shown for the SO2�

4

solutions: in the more diluted solutions (0.01 M)the As concentrations released after 24 h decreasein the following sample order: TS(360 lg/L) > S(250 lg/L) > T(240 lg/L), while in the moreconcentrated solutions (0.1 M) the order of Asrelease is: T(470 lg/L) > TS(330 lg/L) > S(300 lg/L).

The experimental data reported in Fig. 2 (exceptfor the lower right-hand plot) show that the releaseof As is roughly proportional to the amount ofAsHFO in the solids (see Table 1) for the experimentsconducted at the same pH (CO2�

3 and PO3�4 solutions

at pH 8). In contrast, the influence of different pH

conditions on the amount of As(V) desorbed fromthe 3 solids is evident in the experiments withNO�3 , Cl� and pure water solutions. In fact,although the T sample contains much more AsHFO

than the S and TS samples, it releases less Asbecause the desorption of As(V) from HFO is muchless effective at pH 4 than at pH 7.5. The experi-ments with SO2�

4 solutions show that the pH depen-dence is still visible at 0.01 M SO2�

4 concentration,but it is overcome by the probable competitive effectof SO2�

4 as its concentration is increased to 0.1 M.These aspects will be discussed in detail in Section3.1.3.

Finally, it should be noted that the S sample(AsHFO = 1220 mg/kg) generally releases less Asthan the TS sample (AsHFO = 900 mg/kg). This isprobable due to the coarser grain-size (see Table1) and, consequently, the smaller reactive surfacearea of the former.

3.1.2. Effect of reaction time

Fig. 1 shows that desorption of As increases withtime and the desorption equilibrium is not alwaysattained at the end of the experiments (24 h). Inorder to assess the effect of reaction time, the exper-imental data were analyzed according to the first-order rate equation of Lagergren (Altundoganet al., 2000) modified for desorption:

kdes ¼2:303

t� log

Ce

Ce � Ct

; ð6Þ

where kdes is the desorption rate constant, Ce and Ct

are the concentrations of As measured in the batchsolution at pseudo-equilibrium condition (after24 h) and at any sampling time t, respectively.Moreover, the reaction time t1/2 needed for a des-orbed As concentration of Ce/2 was calculated asfollows:

t1=2 ¼0:6932

kdes

: ð7Þ

Plots of log (Ce � Ct) vs. t for all experimental dataare shown in Fig. 3. Linear relationships with highcorrelation coefficients (R2) are generally found,clearly indicating the first-order nature of Asdesorption. Applicability of second- and third-orderrate equations was unsuccessfully tested.

Table 3 reports the values of kdes, as calculatedfrom the slopes of the lines in the plots of Fig. 3,and t1/2 as well. The same results are obtainableusing the residual amounts of As(V) at the surfaceof HFO instead of Ce and Ct in Eq. (6). As an

10

100

1000

10000

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

T S TS

PO4 0.1 MpH 8.0

1

10

100

1000

10000

100000

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

T S TS

CO3 0.1 MpH 8.0

0

100

200

300

400

500

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

T - pH 4.0 S - pH 7.5 TS - pH 7.5

SO4 0.01 M

0

100

200

300

400

500

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

T - pH 4.0 S - pH 7.5 TS - pH 7.5

SO4 0.1 M

0

100

200

300

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

T - pH 4.0 S - pH 7.5 TS - pH 7.5

NO3 0.1 M

0

100

200

300

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

T - pH 4.0 S - pH 7.5 TS - pH 7.5

Cl 0.1 M

0

100

200

300

0 4 8 12 16 20 24 28

Time (hours)

As

(µg/

L)

T - pH 4.0 S - pH 7.5 TS - pH 7.5

Milli-Q H2O

0

20

40

60

80

100

0 2 4 6 8 10 12 14 16 18 20 22 24 26

Time (hours)

% A

s ad

sorb

ed

PO4 CO3SO4 Cl

Synthetic 2-line ferrihydrite 0.1 M pH = 8.3

Fig. 2. Comparison of As desorption from 3 environmental samples in the presence of the same batch solution. Results at 0.01 M anionconcentrations are not plotted, with the exception of SO2�

4 solutions, because results are very similar to the corresponding 0.1 Mexperiments shown. The lower right-hand plot shows the percentage of As(V) adsorbed by synthetic 2-line ferrihydrite at pH 8.3 in thepresence of 0.1 M concentrations of different anionic species (PO3�

4 , CO2�3 , SO2�

4 , and Cl�). The initial As(V) concentration in solution was2 mg/L and the concentration of ferrihydrite (expressed as Fe2O3 � H2O) was 370 mg/L.


Fig. 3. Modified Lagergren plots for As(V) desorption obtained from the results of the kinetic batch experiments. Ce and Ct are the concentrations of As measured in the batchsolution at pseudo-equilibrium condition (after 24 h) and at any sampling time t, respectively.

F.

Fra

uet

al./A

pp

liedG

eoch

emistry

23

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00

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61459

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le3

Val

ues

of

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ant

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et 1

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ion

mea

sure

dat

the

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the

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imen

t)

PO

4C

O3

SO

4N

O3

Cl

H2O

0.01

M0.

1M

0.01

M0.

1M

0.01

M0.

1M

0.01

M0.

1M

0.01

M0.

1M

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

kd

es

(min�

1)

t 1/2

(h)

T0.

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80.

0012

100.

0012

100.

0039

30.

0030

40.

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60.

0014

80.

0018

60.

0016

70.

0023

5S

0.00

186

0.00

186

0.00

1210

0.00

0912

0.00

216

0.00

216

0.00

235

0.00

186

0.00

167

0.00

255

0.00

235

TS

0.00

186

0.00

186

0.00

413

0.00

167

0.00

284

0.00

393

0.00

304

0.00

284

0.00

343

0.00

235

0.00

462


example, the desorption rate constant for As(V) inthe S sample experiments with SO2�

4 solutions is0.0021 min�1 that yields a reaction time t1/2 of 5.5 h.

3.1.3. Considerations on competitive effect

Because HFO is the main As-bearing phase in the3 samples, the results of desorption experiments canbe interpreted considering the combination of twoeffects: the pH effect and the competitive effect ofdissolved anions (i.e. solution composition). In thecase of desorption experiments, ‘‘competitive effect”means the capacity of an ion to displace and substi-tute another ion previously adsorbed at the surfaceof an adsorbent.

The CO2�3 and PO3�

4 solutions exhibit the stron-gest influence on As(V) mobility, and the competi-tive effect in these experiments seems to beprevalent on the pH effect, although the distinctionof the two effects will be possible only with model-ing. The ability of PO3�

4 to compete with As(V)for HFO surface sites was expected since PO3�

4 , likeAs(V), is adsorbed as an inner-sphere complex via aligand-exchange mechanism (Arai and Sparks,2001; Gao and Mucci, 2001). The experiments haveshown that CO2�

3 has a competitive effect similar toPO3�

4 . This suggests, as for PO34, a strong surface

interaction, most probably of inner-sphere nature,between CO2�

3 and HFO surface. As reported inthe introduction, this is a much debated questionthat, however, is supported by some ATR–FTIRspectroscopic studies indicating the formation ofinner-sphere monodentate bonds of CO2�

3 on thesurface of Fe or Al oxides/hydroxides (Villalobosand Leckie, 2000, 2001; Wijnja and Schulthess,1999, 2001). Moreover, according to Appelo et al.(2002) the displacement effect of HCO�3 may offeran explanation for high As concentrations in Ban-gladesh ground water.

A less strong competitor such as SO2�4 is shown

to be partly affected by solution pH. This behaviorcan be interpreted as an indication that the degreeof competition of SO2�

4 increases as a function ofits concentration in solution. Kinjo and Pratt(1971) proposed that when SO2�

4 concentration is60.05 M, SO2�

4 is adsorbed via outer-sphere com-plexation, but inner-sphere adsorption is the preva-lent mechanism at higher SO2�

4 concentration. Morerecently, many researchers have investigated theSO2�

4 adsorption mechanism and the results haveshown that SO2�

4 can be adsorbed as either an outer-sphere complex (Zhang and Sparks, 1990) or aninner-sphere complex (Peak et al., 1999). Therefore,

Table 4Surface parameters of ferrihydrite (HFO) and equilibriumconstants used in the desorption modeling

HFO surface parameters

Stoichiometry: Fe2O3 � H2O; 89 g HFO per mole FeSurface area: 600 m2/gSurface site density: 0.2 mole sites per mole Fe = 2.25 sites/nm2

Adsorption reactions LogK

HFO wOHþHþ ¼ HFO wOHþ2 7.29HFO_wOH = HFO_wO� + H+ �8.93HFO wOHþAsO3�

4 þ 3Hþ ¼ HFO wH2AsO4 þH2O 29.31HFO wOHþAsO3�

4 þ 2Hþ ¼ HFO wHAsO�4 þH2O 23.51HFO wOHþAsO3�

4 ¼ HFO wOHAsO3�4 10.58

HFO wOHþ PO3�4 þ 3Hþ ¼ HFO wH2PO4 þH2O 31.29

HFO wOHþ PO3�4 þ 2Hþ ¼ HFO wHPO�4 þH2O 25.39

HFO wOHþ PO3�4 þHþ ¼ HFO wPO2�

4 þH2O 17.72HFO wOHþ CO2�

3 þHþ ¼ HFO wCO�3 þH2O 12.56HFO wOHþ CO2�

3 þ 2Hþ ¼ HFO wHCO3 þH2O 20.62HFO wOHþ SO2�

4 þHþ ¼ HFO wSO�4 þH2O 7.78HFO wOHþ SO2�

4 ¼ HFO wOHSO2�4 0.79

Only weak binding sites HFO_w are active for anions.


at 0.1 M SO2�4 concentration the mechanism of sur-

face binding of As(V) and SO2�4 is probably similar,

i.e. as an inner-sphere complex, and in that case thecompetitive effect of SO2�

4 prevails over the pHeffect. In fact, the highest As concentration wasreleased from the tailings sample T, that containsmuch more AsHFO than the S and TS samples, eventhough the desorption experiment for the T sampleoccurred under acidic pH conditions that notori-ously favor As(V) adsorption, not desorption. Incontrast, since the competitive capacity of SO2�

4 isstrongly reduced at 0.01 M concentration, it beingadsorbed as an outer-sphere complex, the pH effectprevails and As desorption from the T sample waslower than from the S and TS samples.

The NO�3 and Cl� solutions have negligibleeffects on As(V) desorption, proving to be typicalexamples of outer-sphere complexes that do notinterfere with the inner-sphere binding peculiar toAs in HFO.

The results obtained in this study are supportedby a set of kinetic batch experiments of As(V)adsorption by synthetic 2-line ferrihydrite (paperin preparation). Two-line ferrihydrite was synthe-sized following the procedure of Pierce and Moore(1982). The experiments were conducted at pH val-ues from 4 to 10 and at different anionic concentra-tions. As an example, the lower right-hand plot inFig. 2 reports the percentage of As(V) adsorbed(Asads) versus time at pH 8.3 in the presence of0.1 M concentration of potential competing anions(PO3�

4 , CO2�3 , SO2�

4 and Cl�). The plot shows thatPO3�

4 and CO2�3 strongly inhibit As(V) adsorption

(18% and 28% of Asads, respectively), as a conse-quence of their competitive effect, SO2�

4 behaves asa moderate competitor (Asads = 57%), whereas Cl�

has negligible influence on As(V) adsorption(Asads = 74%).

3.2. Desorption modeling

Modeling of the desorption experiments was per-formed using PHREEQC with the databaseWATEQ4F that comprises the equilibrium con-stants for cation and anion adsorption on HFOcompiled by Dzombak and Morel (1990) andrecently implemented with the carbonate speciesafter Appelo et al. (2002). Table 4 reports theHFO surface parameters and the surface complexa-tion reactions according to the diffuse double layermodel proposed by Dzombak and Morel (1990) alsoknown as GTLM (Generalized Two-Layer Model).

The surface complexation constants for carbonatespecies reported in Appelo et al. (2002)(LogK = 12.78 ± 0.48 and Log K = 20.37 ± 0.20for the singly charged complex and the unchargedcomplex, respectively) are slightly different fromthose included in the WATEQ4F database (seeTable 4) and also contain an indication of parame-ter uncertainty. It is obvious that the uncertainty ofgeochemical simulations performed in this studygenerally derive from the approximations impliedin the estimation of As and Fe speciation in the sol-ids and in the definition of the HFO surface param-eters (i.e. surface area and surface site density) morethan from the uncertainty interval of equilibriumconstants.

The initial surface composition of each samplehas been calculated by equilibration with pure waterwith the concentration of As desorbed after 24 h ofinteraction with the solid. Since that interactiontime may have been insufficient to achieve equilib-rium, the experimental data have been extrapolatedto 96 h with a mathematical function (naperian log-arithm), and the calculated As concentration usedas input value to the simulation.

Calibration of simulation was made on the PO3�4

experiments with the T sample since the surfacecomplexation constants for PO3�

4 and the resultsof T sample experiments are considered more reli-able than the others. A site density lower than thedefault value (1.44 sites/nm2 vs. 2.25 sites/nm2)was also used, when necessary, in order to better


fit the calculated data with the observed ones. Thesimulation was also performed excluding the surfacecarbonate species from calculations in order toassess their importance in the desorption processcompared to pH. The results of modeling for thesamples T, S and TS are reported in Table 5.

Strong undersaturation with respect to mineralsrelated to the background electrolytes used (e.g.trona, natron, nahcolite, thenardite, mirabilite)was calculated in all batch solutions.

With regard to the T sample, overestimation ofAs(V) desorption is obtained in the PO3�

4 experi-ments considering an HFO site density of2.25 sites/nm2, whereas the model is satisfactoryusing 1.44 sites/nm2. This can be explained with apartial incorporation of As in the HFO structure.The incorporated As is not affected by the compet-itive effect of PO3�

4 and, thus, is not released intosolution. Since the amount of incorporated As isunknown, the model uses a lower HFO site density(1.44 sites/nm2) to simulate the incorporation effect.

The amount of As released in the CO2�3 experi-

ments is higher than in the PO3�4 experiments

(Fig. 1) and this is thermodynamically incongruentif only the desorption/adsorption process is consid-ered. As a consequence, a clear underestimation ofAs(V) desorption is calculated for both HFO sitedensities, and a good match is only obtained when,using the speciation data reported in Table 1, com-plete dissolution of scorodite in 0.01 M CO2�

3 solu-tion, or scorodite and arsenopyrite in 0.1 M CO2�

3

solution, with ferrihydrite equilibrium constraint issimulated. The incongruent dissolution of scoroditeand the oxidation of arsenopyrite are processesalready observed and described in the Baccu Loccitailings (Frau and Ardau, 2003, 2004). The overallreactions under alkaline oxidizing conditions canbe schematized as follows:

FeAsO4 � 2H2OðsÞ þ 2OH�ðaqÞ

! FeðOHÞ3ðsÞ þHAsO2�4ðaqÞ þH2OðlÞ; ð8Þ

FeAsSðsÞ þ 3:5O2ðaqÞ þ 4OH�ðaqÞ

! FeðOHÞ3ðsÞ þHAsO2�4ðaqÞ þ SO2�

4ðaqÞ; ð9Þ

where (s), (aq) and (l) stand for solid, aqueous andliquid, respectively. In both reactions As is releasedinto solution, whereas Fe is immobilized in a solidphase. The consumption of OH� is partly bufferedby the couple HCO�3 =CO2�

3 , so that solution pHdoes not change significantly. It is interesting tonote that the simulation totally fails when the car-

bonate surface complexes are not included in themodel; in this case, the surface coverage of As(V) in-creases from 12% to 20% in 0.01 M CO2�

3 solution,and from 4% to 20% in 0.1 M CO2�

3 solution.A strong overestimation of As(V) desorption is

obtained in the SO2�4 experiments for both HFO site

densities. When the sulfate surface complexes arenot included in the model, the simulation providesa good result only in 0.01 M SO2�

4 solution, whereasit is unsatisfactory in 0.1 M SO2�

4 solution. Thisconfirms that at low concentration SO2�

4 behavesas an outer-sphere complex that does not affectAs(V).

With regard to the S sample, the model generallyprovides good results for the PO3�

4 and CO2�3 exper-

iments (especially for an HFO site density of1.44 sites/nm2). It was not necessary to considerreactions other than desorption/adsorption toexplain the concentrations of As(V) released to solu-tion, although partial dissolution of scorodite and/or arsenopyrite may have occurred in the CO2�

3

experiments. The simulation fails when the carbon-ate surface complexes are not considered, leading tounderestimation of As(V) desorption and increaseof As(V) surface coverage from 6% to 12% in0.01 M CO2�

3 solution and from 1% to 12% in 0.1M CO2�

3 solution. The worst fit is for the SO2�4

experiments but, contrary to the T sample model-ing, a clear underestimation of As(V) desorption isobtained.

With regard to the TS sample, it was not neces-sary to consider a lower HFO site density since avery good fit is obtained for the PO3�

4 experimentsusing the default value of 2.25 sites/nm2. This couldbe compatible with the fact that the As bound toHFO in the soil is mainly or totally present asadsorbed species. The model underestimates theamount of As(V) released in the CO2�

3 experiments(especially in 0.1 M solution) but, in this case, a lessprecise assessment of As and Fe speciation (seeTable 1) has not allowed simulation of As-releasingreactions other than desorption. The model clearlyfails in the absence of carbonate surface complexa-tion. The results for the SO2�

4 experiments are likethose obtained in the S sample modeling.

In summary, the best fitting for the experimentsof the T and S samples is obtained using an HFOsite density of 1.44 sites/nm2, whereas the defaultvalue of 2.25 sites/nm2 works better with the TSsample experiments. This finding can be explainedas follows: (i) HFO contained in the tailings/streamsediments formed by precipitation from the As-rich

Table 5Results of PHREEQC desorption modeling for the kinetic batch experiments

Surface coverage (%) As(V) released (lg/L)

As(V) Competing anion Calculated Observed

2.25/nm2 1.44/nm2 2.25/nm2 1.44/nm2 2.25/nm2 1.44/nm2 24 h 96 ha

T sample

Pure water pH 4 48.4 48.1 – – – – 77 100PO4 pH 8

0.01 M 0.08 0.06 63.2 67.2 10150 6492 5830 71500.1 M 0.01 0.01 88.8 93.6 10164 6498 4752 5550

CO3 pH 80.01 M – noSurf 20.6 19.8 – – 5833 3823 10413 118100.01 M 12.2 11.7 39.1 40 7598 49180.01 M – ScorDis 13.8 13.7 41.3 41.8 14776 121490.1 M – noSurf 20.2 20.6 – – 5923 3721 15593 177200.1 M 4.12 4.56 82.9 83.9 9301 58850.1 M – ScorDis 5.07 6.09 81.6 81.8 16601 131770.1 M – Scor + ArsDis 5.36 6.51 81.4 81.5 20242 16822

SO4 pH 40.01 M – noSurf 46.6 44.6 – – 371 477 242 3100.01 M 28.4 22.7 28.7 31.7 4210 34350.1 M – noSurf 42.7 37 – – 1195 1508 472 5800.1 M 17.5 13.1 35.1 35 6490 4727

S sample

Pure water pH 7.5 15.9 15.9 – – – – 130 160PO4 pH 8

0.01 M <0.01 <0.01 63.0 67.2 936 596 695 8600.1 M <0.01 <0.01 88.8 93.7 936 597 501 603

CO3 pH 80.01 M – noSurf 12.6 12.2 – – 196 139 391 4370.01 M 6.52 6.02 45.2 47.4 553 3710.1 M – noSurf 12.5 12.4 – – 198 131 576 6340.1 M 1.12 1.20 87.8 88.9 871 551

SO4 pH 7.50.01 M 14.6 14.2 1.29 1.51 79 64 250 3110.1 M 13.4 12.8 3.78 4.50 148 115 301 371

TS sample

Pure water pH 7.5 16.5 – – – – – 230 230PO4 pH 8

0.01 M <0.01 – 63.2 – 707 – 632 7760.1 M <0.01 – 88.8 – 707 – 643 776

CO3 pH 80.01 M – noSurf 12.4 – – – 178 – 518 6880.01 M 6.07 – 46.4 – 448 –0.1 M – noSurf 12.3 – – – 180 – 898 10640.1 M 0.92 – 88.2 – 668 –

SO4 pH 7.50.01 M 14.7 – 1.48 – 77 – 361 4890.1 M 13.2 – 4.17 – 141 – 328 411

noSurf means that the corresponding surface complexes are excluded from the model; ScorDis means that scorodite dissolution is includedin the model; Scor + ArsDis means that scorodite and arsenopyrite dissolution is included in the model.

a Amount of As(V) released after 96 h extrapolated by fitting with a mathematical function (naperian logarithm) the values observed in24 h.


slurry discharged from the flotation plant (Frau andArdau, 2003, 2004) and, therefore, some As boundto HFO might be incorporated in the HFO struc-

ture owing to co-precipitation or diffusion from sur-face to bulk (absorption); (ii) HFO (also includingsome goethite) contained in the soil mainly formed


by pedogenetic processes, although a detrital frac-tion was also recognized (Dadea, 2003), and Aswas adsorbed from pore waters.

The failure of simulation when the carbonate sur-face complexes are not included in the model indi-cates that carbonate species have a clear influencein the desorption of As(V). The surface coverageof the carbonate species (40–50% and 80–90% forthe 0.01 M and 0.1 M experiments, respectively) isessential to account for the amounts of As(V)released in the experiments with CO2�

3 solutions.For instance, the pH factor only accounts for about35% and 23% of As(V) desorbed from the S sample,respectively in the 0.01 M and 0.1 M CO2�

3 solu-tions, whereas the competition effect of carbonatespecies accounts for the remaining 65% and 77%,respectively.

Also, the results indicate that the adsorption con-stants for SO2�

4 are unreliable, with strong overesti-mation of SO2�

4 adsorption at pH 4 (T samplemodeling) and underestimation at pH 7.5 (modelingof S and TS samples).

4. Environmental implications for As mobility and

possible applications

This study has shown that the adsorption/desorption mechanisms known from the literatureas involving As(V) and possible competitors canbe macroscopically recognized and modeled alsowhen natural complex heterogeneous samples areused instead of a single pure synthetic phase (i.e. fer-rihydrite), provided a detailed characterization ofthe natural solid samples is obtained.

This study confirms the authors’ preliminaryleaching/adsorption column tests on As-contami-nated samples from the Baccu Locci mine con-ducted with different solutions with 100 mg/L ofthe anions Cl�, SO2�

4 and HCO�3 (Frau et al.,2004). Moreover, it sheds further light on a particu-larly interesting aspect: the displacement action ofaqueous carbonate species on As adsorbed ontohydrous ferric oxides. This mechanism has beenproposed by Appelo et al. (2002) to explain the highAs concentrations in Bangladesh ground water and,after this study, is also thought to contribute to thesevere As contamination occurring in surface andground waters of the Baccu Locci – Quirra area inSardinia, Italy (Frau and Ardau, 2003). Upon mix-ing with an HCO3-dominated water from an artifi-cial lake, the Baccu Locci stream water chemistrylocally changes from a Ca–Mg–SO4 type to a Ca–

Mg–HCO3 type downstream, and this change iscombined with a sharp increase in dissolved As con-centration (from 200 lg/L to 350 lg/L) that, in spiteof the conservative behavior of most elements, is notexplained by mixing (Frau and Ardau, 2003).Arsenical ferrihydrite contained in the flotation tail-ings/stream sediments had precipitated from slightlyacidic SO4-dominated slurry waters discharged fromthe treatment plant during mine activity. After mineclosure in 1965, more dilute HCO3–(SO4) watershave progressively interacted with the ferrihydrite-bearing stream sediments/tailings, and the near-neu-tral pH conditions and the probable displacementaction of HCO�3 have favored As desorption, caus-ing diffuse and persistent contamination.

Another important aspect concerns the potentialcompetitive capacity of SO2�

4 as a function of its dis-solved concentration, although the need to recon-sider the adsorption constants for SO2�

4 clearlyemerges from modeling.

Possible applications in various fields, such as irri-gation agriculture, soil decontamination, water treat-ment and mine site remediation, might be derivedfrom these findings. As an example, the strong dis-placement capacity of aqueous carbonate speciesmight be used to naturally clean As-containing soils.On the other hand, irrigation of As-contaminatedsoils with common HCO3-dominated waters mightgreatly mobilize As, with consequent risk of Asuptake by plants and ground water contamination.

As a closing remark on this study, it is necessaryto make it clear that the results obtained cannot begeneralized or directly applied to natural systems.Further field and laboratory studies are necessaryto clarify and verify the role of potential competi-tors such as aqueous carbonate species in desorbingAs from natural sediments.

Acknowledgements

Many thanks to the students Deborah Congiuand Giorgia Usai that conducted most of the batchexperiments in the ambit of their Degree Thesiswork. A special thank to Dr. C.A.J. Appelo forhis kind help and useful hints on modeling. We alsothank two anonymous reviewers for their valuablecomments, and the AG Associate Editor Dr. DavidPolya for handling the manuscript. This study wassupported by the Italian Ministry of Education,University and Research (PRIN 2006 to L. Fanfani;ex-60% funds to F. Frau) and by Fondazione Bancodi Sardegna.


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Effect of major anions on arsenate desorption from ferrihydrite-bearing natural samples

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