Kinetics of sorption and abiotic oxidation of arsenic(III)

15 Pages·2006·0.35 MB·English

GeochimicaetCosmochimicaActa70(2006)533–547
www.elsevier.com/locate/gca
Kinetics of sorption and abiotic oxidation of arsenic(III)
by aquifer materials
Aria Amirbahmana,*, Douglas B. Kentb, Gary P. Curtisb, James A. Davisb
aDepartmentofCivilandEnvironmentalEngineering,UniversityofMaine,Orono,ME04469,USA
bUSGeologicalSurvey,345MiddlefieldRd.,MenloPark,CA94025,USA
Received16May2005;acceptedinrevisedform4October2005
Abstract
The fate of arsenic in groundwater depends largely on its interaction with mineral surfaces. We investigated the kinetics of As(III)
oxidationbyaquifermaterialscollectedfromtheUSGSresearchsiteatCapeCod,MA,USA,byconductinglaboratoryexperiments.
Five different solid samples with similar specific surface areas (0.6–0.9m2g(cid:1)1) and reductively extractable iron contents (18–
26lmolm(cid:1)2), but with varying total manganese contents (0.5–3.5lmolm(cid:1)2) were used. Both dissolved and adsorbed As(III) and
As(V)concentrationsweremeasuredwithtimeupto250h.TheAs(III)removalratefromsolutionincreasedwithincreasingsolidman-
ganesecontent,suggestingthatmanganeseoxideisresponsiblefortheoxidationofAs(III).Underallconditions,dissolvedAs(V)con-
centrationswereverylow.Aquantitativemodelwasdevelopedtosimulatetheextentandkineticsofarsenictransformationbyaquifer
materials.Themodelincluded:(1)reversiblerate-limitedadsorptionofAs(III)ontobothoxidativeandnon-oxidative(adsorptive)sites,
(2)irreversiblerate-limitedoxidationofAs(III),and(3)equilibriumadsorptionofAs(V)ontoadsorptivesites.Rateconstantsforthese
processes,aswellasthetotaloxidativesitedensitieswereusedasthefittingparameters.Thetotaladsorptivesitedensitieswereestimated
basedonthemeasuredspecificsurfaceareaofeachmaterial.Thebestfitwasprovidedbyconsideringonefastandoneslowsiteforeach
adsorptiveandoxidativesite.Thefittingparameterswereobtainedusingthekineticdataforthemostreactiveaquifermaterialatdif-
ferentinitialAs(III)concentrations.UsingthesameparameterstosimulateAs(III)andAs(V)surfacereactions,themodelpredictions
werecomparedtoobservationsforaquifermaterialswithdifferentmanganesecontents.Themodelsimulatedtheexperimentaldatavery
wellforallmaterialsatallinitialAs(III)concentrations.TheAs(V)productionratewasrelatedtotheconcentrationsofthefreeoxida-
tive surface sites and dissolved As(III), as r ¼k0 ½(cid:1)MnðIVÞOH(cid:2)½H AsO (cid:2) with apparent second-order rate constants of
AsðVÞ ox 3 3
kf0 ¼6:28(cid:3)10(cid:1)1andks0 ¼1:25(cid:3)10(cid:1)2M(cid:1)1s(cid:1)1forthefastandtheslowoxidativesites,respectively.TheAs(III)removalratedecreased
ox ox
approximately by half for a pH increase from 4 to 7. The pH dependence was explained using the acid–base behavior of the surface
oxidative sites by considering a surface pK =6.2 (I=0). In the presence of excess surface adsorptive and oxidative sites, phosphate
a
diminishedtherateofAs(III)removalandAs(V)productiononlyslightlyduetoitsinteractionwiththeoxidativesites.Theobserved
As(III)oxidationratehereisconsistentwithpreviousobservationsofAs(III)oxidationovershorttransportdistancesduringfield-scale
transportexperiments.Themodeldevelopedheremaybeincorporatedintogroundwatertransportmodelstopredictarsenicspeciation
andtransport in chemically heterogeneous systems.
(cid:1)2005Elsevier Inc. Allrights reserved.
1. Introduction contaminantlevel(MCL)of10lgL(cid:1)1forarsenicindrink-
ing water (WHO, 1993). Depending on its source, arsenic
Arsenic contamination of drinking water is an issue of concentrations in natural waters may range up to several
great concern. Due to its acute toxicity to humans, the hundred milligrams per liter. Arsenic contamination of
World Health Organization (WHO) has set a maximum groundwaterhasbeenofgreatinterestinseveralcountries,
most notably in Bangladesh and West Bengal, where two-
thirdsofthepopulation areatriskofserioushealtheffects
* Correspondingauthor.Fax:[phone].
E-mailaddress:[email protected](A.Amirbahman). duetohighconcentrationsofarsenicintheshallowalluvial
0016-7037/$-seefrontmatter (cid:1)2005ElsevierInc.Allrightsreserved.
doi:10.1016/j.gca.2005.10.036
534 A.Amirbahmanetal.70(2006)533–547
and deltaic aquifers that are the main sources of drinking volvesaseriesofsteps,includingdiffusionofAs(III)tooxi-
water (Manzurul Hassan et al., 2003). dative sites, ligand exchange between As(III) and
Mobilization of sediment-bound arsenic in natural and manganese oxide surface functional groups, and inner-
contaminatedgroundwaterhasbeenobservedunderavari- sphere electron transfer. Alternatively, outer-sphere elec-
ety of chemical conditions (Smedley and Kinniburgh, trontransfermayoccurbetweenoxidativesitesandAs(III)
2002). Increasing arsenic mobility in groundwater in re- through their respective coordination sheaths. The pre-
sponse to elevated pH values (>10) has been reported dominantly Mn(III) oxide manganite (c-MnOOH) has
(Marineretal.,1996).Inoxicenvironments,releaseofsed- been shown to oxidize As(III) and to adsorb both As(III)
iment-bound arsenic has been observed in response to ele- and As(V) species (Chui and Hering, 2000). Adsorption
vated concentrations of phosphate (Kent and Fox, 2004) of As(III) onto manganese oxides consisting predominant-
and oxidation of sulfide minerals at low pH values ly of Mn(IV) has not been observed. Adsorption of As(V)
(McCreadie et al., 2000; Schreiber et al., 2000). In anoxic onto manganese(IV) oxides, however, has been reported
environments, sediment-bound arsenic can be mobilized (e.g., Tani et al., 2004), and spectroscopic studies have
by reductive dissolution or transformation of Fe(III) shown that the structure of the As(V) surface complex on
hydroxide coatings and reduction of arsenate (As(V)) to manganese oxides is similar to that on Fe(III) hydroxides
arsenite (As(III)) (e.g., Belzile and Tessier, 1990; Ahmann (Manning et al., 2002; Deschamps et al., 2003; Foster
et al., 1997; Harrington et al., 1998; Cummings et al., et al., 2003).
1999; McArthur et al., 2001; Ho¨hn et al., 2001; Harvey In the presence of natural composite surfaces, both oxi-
et al., 2002; Horneman et al., 2004; Kent and Fox, 2004; dation and adsorption reactions should be considered for
Swartz et al., 2004; van Geen et al., 2004). the removal of arsenic from solution. The Fe(III)- and Al-
Speciationofarseniccontrolsitsmobilityandtoxicityin bearing minerals in natural soil and sediment surfaces ad-
naturalwaters.Inmostnaturalwaters,arsenicspeciationis sorb arsenic (Pierce and Moore, 1980; Fuller et al., 1993;
dominated by the oxyanions As(III) and As(V) (Smedley Manning and Goldberg, 1996; Wilkie and Hering, 1996;
and Kinniburgh, 2002). In general, As(V) adsorption to Fendorf et al., 1997; Raven et al., 1998; Dixit and Hering,
mineral oxides is favored at low pH, whereas maximum 2003; Arai et al., 2005). Arsenic, however, has a higher
As(III) adsorption occurs at circumneutral pH and affinity to Fe(III) hydroxides than to Al hydroxides and
decreases with an increase or a decrease in pH (Raven clay minerals (Manning and Goldberg, 1997;Arai et al.,
et al., 1998; Arai et al., 2001; Dixit and Hering, 2003). 2001). Adsorption of arsenic species onto Fe(III) hydrox-
As(III), however, is known to be more toxic than As(V) ides is pH-dependent. At acidic pH, As(V) generally sorbs
to humans (National Research Council, 1999). Therefore, more favorably than As(III), whereas at basic pH, the
understanding the redox speciation of inorganic arsenic trend is reversed (Manning et al., 1998; Raven et al.,
involvingAs(III)andAs(V)speciesisofgreatimportance. 1998; Dixit and Hering, 2003). The pH behavior of
Both biotic and abiotic processes can lead to redox As(III) adsorption onto Fe(III) hydroxides, however,
transformation of arsenic species. Several microbial respi- shows a smaller pH dependency than that of As(V).
ratory and non-respiratory enzymatic systems for oxida- The exact pH adsorption behavior may be attributed to
tion of As(III) have been reported (Oremland and Stolz, the acid–base characteristics of aqueous As(V) and
2003).Innaturalsystems,themostimportantabioticpath- As(III), and acid–base characteristics and adsorption
way for the oxidation of As(III) is by minerals containing affinity of surface functional groups on the solids (Davis
Mn(III) and Mn(IV) (Oscarson et al., 1981a,b; Brannon and Kent, 1990; Dzombak and Morel, 1990). Competitive
and Patrick, 1987). Experimental studies using synthetic or cooperative effects induced by the presence of other
and biogenic manganese oxides have shown that these sol- adsorbing anions or cations can significantly influence
ids can effectively oxidize As(III), albeit at different rates the adsorption and mobility of As(III) and As(V) (Wilkie
that span an order of magnitude (Oscarson et al., 1983; and Hering, 1996; Darland and Inskeep, 1997; Gra¨fe
Moore and Walker, 1990; Scott and Morgan, 1995; Chui et al., 2004; Kent and Fox, 2004).
and Hering, 2000; Tournassat et al., 2002; Tani et al., Over short timescales, as might be encountered in labo-
2004). Tournassat et al. (2002) attributed differences in ratory experiments and some surface- and groundwater
As(III) oxidation rate to the degree of crystallinity and applications, arsenic interaction with mineral surfaces
structure of manganese minerals. An experimental study may be controlled kinetically (Johnson and Pilson, 1975;
on sediment samples showed that removal of manganese Seyler and Martin, 1989; Kuhn and Sigg, 1993). In batch
oxides resulted in a marked decrease in the As(III) oxida- laboratory experiments, arsenic adsorption kinetics onto
tion rate (Oscarson et al., 1981a). Fe(III) hydroxides have been shown to be relatively fast,
PreviousstudiesofAs(III)oxidationkineticsbysynthet- with a faster rate observed for As(V) at lower pH values,
ic manganese oxides show that the kinetics of aqueous and for As(III) at higher pH values (Raven et al., 1998;
As(III) removal do not display a simple first-order behav- Waltham and Eick, 2002). The extent and kinetics of
ior (Moore et al., 1990; Scott and Morgan, 1995; Chui As(V)andAs(III)adsorptionontoFe(III)hydroxideswere
and Hering, 2000; Manning et al., 2002; Tournassat shown to be dependent on the solute:sorbent ratio (Pierce
et al., 2002). As(III) oxidation by manganese oxides in- and Moore, 1980; Raven et al., 1998).
As(III)oxidationbyaquifermaterial 535
The above-mentioned studies have investigated arsenic imately 2km downgradient at a site previously designated
interactions with synthetic and, in a few cases, naturally the ‘‘oxic’’ site (Kent et al., 1995). These materials consist
occurring manganese oxides. However, studies that model ofapproximately90%quartzwiththeremainderasplagio-
arsenic interaction kinetics with natural composite materi- clase and K-feldspars, magnetite, hematite, goethite, glau-
als are lacking. In this work, we have conducted batch conite, and lithic fragments (Barber, 1990; Wood et al.,
experiments to examine the adsorption and oxidation of 1990; Bau et al., 2004). The surfaces of quartz and other
As(III) by aquifer materials from five different locations mineral grains are, however, coated with nm-size Fe(III)-
at the USGS research site on Cape Cod, MA, USA (Le- and Al-hydroxides and/or silicates (Coston et al., 1995;
Blanc, 1984; LeBlanc et al., 1991; Kent et al., 1995; Kent Banfield and Hamers, 1997). Aquifer materials were col-
and Fox, 2004). The aquifer is shallow and unconfined, lected using a wire-line core barrel (Zapico et al., 1987)
and its mineralogy is dominated by quartz, feldspars, and andfrozen. Theywereair-dried,sievedtoremove material
other silicate minerals. These materials contain relatively greater than 2mm in diameter, and stored dry until use.
similar specific surface areas and reductively extractable All chemicals used in this study were reagent grade. A
ironcontents.However,theydiffergreatlyinthereductive- 0.05M concentration stock solution of As(III) was pre-
lyextractablemanganesecontents.Theobjectiveofthispa- paredinananaerobicglovebagbydissolvingNaAsO salt
2
peristodevelopamacroscopickineticmodelbasedonthe (Baker) (brand names are for identification purposes only
laws of surface complexation and mass-action that ac- and should not be taken to construe endorsement by the
counts for uptake of arsenic by adsorption and oxidation US Geological Survey) in an oxygen-free 0.01M solution
inthepresenceofaquifermaterialsfromtheCapeCodsite. of HCl. This solution was stored in a Teflon bottle in the
Thismodelinvolvesamulti-stepreactionschemeinvolving dark in an anaerobic glove bag. A 0.05M concentration
adsorption of As(III) onto oxidative and adsorptive sites of As(V) stock solution was prepared by dissolving Na
2-
on the aquifer material, electron transfer at the oxidative HAsO Æ7H O salt (Baker) in a 0.01M solution of HCl.
4 2
sites followed by As(V) release, and the reaction of As(V) This solution was stored in a Teflon bottle in the dark. A
with the adsorptive sites. The model developed here is cal- 2M stock solution of NaCl (Aldrich) was used to adjust
ibrated using experimental data with aquifer materials un- the background ionic strength of the samples. The buffer
der varying chemical conditions and may be used to solutions of sodium acetate (Baker) and 2,2-bis(hydroxy-
simulatethemajorfeaturesofreactivetransportofarsenic methyl)-2,20,200-nitrilotriethanol (Bis-Tris; Aldrich) were
in chemically heterogeneous groundwater systems. freshly prepared before the start of each experiment.
Reductiveextractionofironandmanganesefromtheaqui-
2. Experimental section fer materials was performed with hydroxylamine hydro-
chloride (NH OHÆHCl; Aldrich). Oxalic acid, HCl,
2
2.1. Materials HNO , and NaOH were obtained from Aldrich.
3
Aquifer materials from five different locations in the 2.2. Experimental setup
USGS groundwater experiment site in Cape Cod, MA
(USA)wereused(Table1).Allsolidsampleswereobtained The As(III) oxidation experiments were conducted by
fromaregionoftheaquiferthatwasoxic(250–350lMdis- adding 17.50±0.05g of dry aquifer material to a 25ml
solved oxygen), had low concentrations of dissolved salts buffer solution in 50mL polycarbonate centrifuge tubes.
(specificconductance<100lScm(cid:1)1),andacidicpHvalues The buffers were 5mM sodium acetate (pH 5.2 and 6.0)
(4.8–5.8)atlocations described indetailelsewhere.Sample and5mMBis-Tris(pH6.9).Theionicstrengthwasadjust-
F415-19 was obtained adjacent to a site previously desig- ed to 10mM by addition of NaCl. The tubes were rotated
nated the ‘‘suboxic’’ site (Kent et al., 1994). Sample end-over-end at approximately 13rpm in the dark for at
R23AW was a composite of materials obtained from 2 least 16h prior to the addition of As(III). A relatively
cores(R23AWC02andC03)adjacenttoasitewherearsen- low mixing rate was chosen to avoid excessive abrasion
ic tracer tests had been conducted (Stadler et al., 2001; of the solid materials. The experiments were initiated by
KentandFox,2004).SamplesF168wereobtainedapprox- adding aliquots of As(III) from the stock solution to the
Table1
Characteristicsoftheaquifermaterialsa
Aquifermaterial Altitude Specificsurface Reductivelyextracted Reductivelyextracted
(mabovesealevel) area(m2g(cid:1)1) Fecontent(lmolm(cid:1)2) Mncontent(lmolm(cid:1)2)
F168-15 6.9±0.8 0.608±0.074 20.06±1.05 3.36±0.20
F168-20 5.3±0.8 0.600±0.005 22.74±1.82 2.51±0.23
F168-25 3.8±0.8 0.845±0.131 18.18±3.02 1.41±0.23
F415-19 13±0.3 0.418±0.042 26.07±0.84 1.75±0.14
R23AW 13.15±0.45 0.531±0.115 23.76±1.01 0.59±0.03
a Errorisbasedononestandarddeviationfromthemeanwithn=3.
536 A.Amirbahmanetal.70(2006)533–547
suspensions. Initial As(III) concentrations ranging from analytically to the known concentration) were maintained
approximately 25 to 500lM were used. The tubes were attheoptimalvaluesof5and110%,respectively,byinsur-
thenrotatedinthedarkforthedurationoftheexperiment. ing arsenic concentrations exceeded the limit of quantita-
Experiments were conducted both in the presence and ab- tion of 0.13lM, at which concentration the relative
senceofoxygen.Thelattersetofexperimentswasconduct- precision and accuracy greatly increased.
ed inside an anaerobic glove bag (Coy Laboratory Approximately 5g of each aquifer material sample was
Products) in a N atmosphere with 4% H . At specific used for the N -BET specific surface area measurements
2 2 2
times, the tubes were centrifuged and the supernatant was (Micromeritics Tristar). Specific surface area determined
removed for the analysis of dissolved arsenic concentra- by N adsorption on this instrument was checked against
2
tion. Kinetic experiments were continued for times longer a quartz sample whose specific surface area of
than150h until steady-stateconditions,especiallywithre- 0.33m2g(cid:1)1hadbeendeterminedbyKrandN adsorption
2
spect to the production of As(V), were nearly reached. (Kohleretal.,1996).Theinstrumentproducedreliablespe-
Speciation of the dissolved arsenic was determined by cific surface area values as long as the measurement was
passing 5mL of the supernatant followed by 5mL of made on a sample large enough to provide at least 1m2
deionized water through a strong anion exchange resin of surface area. Chemical extractions were used to deter-
(SAX; Alltech) with an exchange capacity of 0.15meq/ mine the reductively extractable iron and manganese con-
100mg. The resin was conditioned prior to use with tents of the materials. The extractions were performed by
2mL of methanol followed by 10mL of deionized water adding the same samples used for the BET measurements
(Bednar et al., 2002). Under our experimental conditions, to a 25ml solution of 0.25M NH OHÆHCl and 0.25M
2
As(V) is present predominantly as H AsO (cid:1) and is re- HCl. The suspension was then shaken in the dark at
2 4
moved by the SAX resin, whereas As(III) is present pre- 50(cid:2)C for 96h. The supernatant was filtered and diluted
dominantly as H AsO and elutes through the resin at a 1:1 ratio in 0.15M HNO and analyzed for iron and
3 3 3
column. manganese contents. The NH OHÆHCl extraction tech-
2
Solid-bound arsenic species were desorbed by adding nique used here was found to be effective in assessing var-
25mL of 0.2M oxalic acid after the removal of the super- iability in adsorption properties of materials from the site
natant.The tubes werethenrotatedinthedarkfor15min (Fulleretal.,1996).Surfaceareameasurementsandextrac-
and centrifuged to remove the supernatant containing the tions for all materials were performed in triplicates.
desorbed arsenic species. Preliminary experiments showed
that arsenic species were effectively desorbed using a 2.4. Modeling
15min extraction, and did not change their oxidation
states, as also shown by De Vitre et al. (1991). Longer Theinteractionsofarsenicspecieswiththeaquifermate-
extractiontimesdidnotyieldsignificantlyhigherextracted rialsdescribedbyasetofequilibriumandkineticreactions
arsenic concentrations. Extraction with oxalic acid was were modeled using the computer program RATEQ (Cur-
found to be more effective and more expedient than other tis, 2005). RATEQ can simulate multiple equilibrium and
methods,includinghot4MHCland0.02MNaH EDTA. rate-controlledreactionsinbatchsystemsaswellasinmul-
2
Speciation of the adsorbed arsenic species was deter- tidimensional groundwater systems. For simultaneous
mined by adjusting the pH of the supernatant from the equilibrium and rate-controlled reactions, RATEQ solves
oxalic acid extract to approximately 3.8 by adding 10M a set of non-linear mole-balance and mass-action equa-
NaOH. A 1mL solution of the extract supernatant fol- tions. The code requires the user to specify the nature of
lowed by 7mL of deionized water were then passed the reactions (equilibrium or kinetic), the corresponding
through the preconditioned SAX resin. As(III) concentra- equilibrium or rate constants, total concentrations of the
tion was determined by measuring the resin eluate, and components, and the time intervals at which concentra-
the As(V) concentration was determined by the difference tions are calculated. The computer program UCODE was
between the total arsenic concentration in the supernatant usedtoperforminversemodelingtofitasetofequilibrium
and the As(III) concentration. andkineticequationstotheexperimentaldata(Poeterand
Hill,1998).UCODEsolvesthenon-linearregressionprob-
2.3. Analytical methods lemby minimizingaweightedleastsquaresobjectivefunc-
tion with respect to the parameter values. An estimated
Arsenicwasmeasuredusinginductivelycoupledplasma parameter is a quantity that appears in the input files of
optical emission spectrometry (ICP-OES; Thermo Jarrell- the application model. The parameters to be estimated in
AshIrisdual-view).Priortoanalysis,sampleswerefiltered thiscaseweretheequilibriumandrateconstants,andtotal
using 0.45lm polyvinylidene fluoride filters (Millex) and concentrationsofthesurfacesites.Parameterestimationis
acidified to pH 2 with 5M HNO . Samples with high performed by running the input files of the application
3
arsenicconcentrationswerediluteduptoa10:1ratio.Rel- model, in this case several RATEQ files, simultaneously.
ative precision (2 times the standard deviation divided by Each of the RATEQ files contained model parameters for
the average concentration determined analytically) and As(III)oxidationbyoneaquifermaterialatdifferentinitial
accuracy (ratio of the average concentration determined As(III) concentrations or pH. UCODE runs the
As(III)oxidationbyaquifermaterial 537
simulations,comparestheobservedandsimulatedconcen- 3.1. Experimental results of arsenic surface interactions
trations, and performs sensitivity calculations for each
modelparameter.Thisallowstheusertoobservethesignif- Figs.1–5showthereactionkineticsofarsenicspeciesat
icance of each parameter to the simulated values, and the pH 5.2 at initial As(III) concentrations ranging from
significance of the individual experimental observations approximately 30 to 500lM for the five different aquifer
to the estimation of the different parameters. Sensitivity materials. Production of the As(V) species takes place
analysis also aids in the identification of the most sensitive simultaneously as the depletion of aqueous As(III) and
parameters. Fitting parameters with high sensitivity coeffi- the appearance of solid-bound As(III). Due to its very
cientsareestimatedwithahigherprecisionthanthosewith low concentrations, aqueous As(V) is reported only for
low sensitivity coefficients. the most reactive material (F168-15) at the highest initial
As(III) concentration (Fig. 1D, inset).
3. Results and discussion Mass balance of arsenic species was achieved as the
summation of solid-phase As(III) and As(V) concentra-
Aquifer materials from Cape Cod contain a range of tions and total aqueous As(III) concentration. In all
reactive functional groups that originate from the mixture samples, mass balance was within 12% of the initial
of mineral phases and their surface coatings (Davis et al., As(III) concentration, suggesting that the 15min 0.2M
1998). These coatings are primarily comprised of Fe, Al, oxalic acid extraction is an effective method for the
and Si, with lesser amounts of Mn (Coston et al., 1995), desorption of arsenic species from aquifer materials.
likely in the form of Fe(III) hydroxides, Al hydroxides The 15min extraction period is also acceptable consid-
and silicates, and Mn oxides (Banfield and Hamers, ering the relatively slow transformation rates in this
1997). Because the sediments can adsorb oxyanions (Kent work. Lack of any measurable concentrations of aque-
et al., 1995; Stollenwerk, 1995) like As(III) and As(V), it ous As(V) also indicates that to correctly estimate the
was necessary to quantify redox speciation of arsenic in transformation kinetics of As(III) by natural materials,
both dissolved and solid phases. This was accomplished adsorbed As(V) concentration must be directly
using a 15min extraction with oxalic acid. measured.
A 30 B 60
solid As(V) 50
solid As(V)
20 40
30
10 20
aq. As(III)
aq. As(III)
M) 10
µ solid As(III) solid As(III)
n ( 0 0
o
ati 0 50 100 150 0 50 100 150
ntr
ce C100 D 500
n 3
o
c
80 400 2
aq. As(III) aq. As(III)
solid As(V) 1 aq. As(V)
60 300
0
0 50 100 150 200 250
40 200
solid As(V)
20 100
solid As(III)
solid As(III)
0 0
0 50 100 150 0 50 100 150 200 250
time (hr)
Fig.1. OxidationkineticsofAs(III)byF168-15material.pH5.2,andinitial[As(III)]=28.3lM(A),59.2lM(B),88.7lM(C),and483.0lM(D).All
linesaremodelfitstotheexperimentaldata.
538 A.Amirbahmanetal.70(2006)533–547
A 30 A 30
20 20
aq. As(III) solid As(V)
aq. As(III)
solid As(V)
10
10
solid As(III)
M) solid As(III)
µ
n ( 0 0
o
ati
ntr B 500 B 100 2
e
c
on M) 80
c 400 µ
aq. As(III) n (
o
300 ati 60 aq. As(III)
ntr
e
c 40
200 n solid As(III)
o
solid As(V) c
100 20 solid As(V)
solid As(III)
0
0
0 50 100 150 200
C 500 0
time (hr)
Fig. 2. Oxidation kinetics of As(III) by F168-20 material. pH 5.2, and 400
initial[As(III)]=24.7lM(A),and480.3lM(B).Alllinesaremodelfits
totheexperimentaldata. aq. As(III)
300
Arsenic displayed varying transformation rates depend- 200
solid As(III)
ingontheaquifermaterial.Thedisappearanceoftheaque-
ous As(III) species and the formation of the adsorbed 100
As(III) and As(V) species were the fastest with F168-15
solid As(V)
material (Fig. 1), and slowest with R23AW material 0
(Fig. 5). The extent of aqueous As(III) oxidation also fol- 0 50 100 150 200
lows the same order with respect to the aquifer materials. time (hr)
For F168-15 material, oxidation, as compared to adsorp-
tion, was clearly the dominant mechanism for aqueous Fig. 3. Oxidation kinetics of As(III) by F168-25 material. pH 5.2, and
initial [As(III)]=28.3lM (A), and 84.4lM (B), and 495.4lM (C). All
As(III)removal,wheremorethan90%oftheaddedAs(III)
linesaremodelfitstotheexperimentaldata.
at concentrations below 100lM oxidized to As(V) (Figs.
1A–C). Oxidation was also the dominant mechanism for
As(III) removal for F168-20 material at the concentration 3.2. Anaerobic and sterile controls
range studied here (Fig. 2). For F168-25 (Fig. 3) and
F415-19 (Fig. 4) materials, however, oxidation was the Arsenic(III)oxidationexperimentswereconductedboth
dominant mechanism for As(III) removal at initial As(III) in the presence and absence of dissolved oxygen at pH 5.2
concentrations below 30lM, whereas, oxidation and withF168-20material.Thesamematerialwasalsousedto
adsorption contributed nearly equally to the aqueous conduct As(III) oxidation experiments in the presence of
As(III) removal kinetics at initial As(III) concentrations 0.1Mformaldehyde,abactericide.Inbothcases,nodiffer-
near 90lM. At initial As(III) concentrations as high as encesintheaqueousAs(III)removalandAs(V)production
500lM, As(III) adsorption was clearly the dominant rates were observed (results not shown here). Addition of
mechanism for the removal of aqueous As(III) for these nitrate to the same aquifer material to promote the auto-
aquifermaterials(Figs.3Cand4C).ForR23AWmaterial, trophicoxidationofAs(III)bymicroorganisms(Oremland
As(III) adsorption was the dominant mechanism for and Stolz, 2003) also did not result in the enhanced oxida-
As(III) removal at all concentrations studied here (Fig. 5). tion of As(III) (S. Hoeft, USGS, personal communica-
As(III)oxidationbyaquifermaterial 539
A 30 A 30
20 20
aq. As(III) solid As(V) aq. As(III)
10 10
solid As(III)
solid As(III) solid As(V)
0
0
B100 0 50 100 150 200
B 100 0 50 100 150 200
80
µM) aq. As(III) M) 80
concentration ( 246000 solid As(V) µoncentration ( 4600 aq. As(III) solid As(III)
solid As(III) c 20
solid As(V)
0
0
C 500 0 50 100 150 200 C 500 0 50 100 150 200
400 aq. As(III) aq. As(III)
400
300
300
200
200
solid As(III)
100
solid As(III)
100
solid As(V)
0 solid As(V)
0 50 100 150 200 0
0 50 100 150 200
time (hr)
time (hr)
Fig. 4. Oxidation kinetics of As(III) by F415-19 material. pH 5.2, and
initial [As(III)]=28.3lM (A), and 88.7lM (B), and 499.8lM (C). All Fig. 5. Oxidation kinetics of As(III) by R23AW material. pH 5.2, and
linesaremodelfitstotheexperimentaldata. initial [As(III)]=28.5lM (A), and 85.1lM (B), and 500.3lM (C). All
linesaremodelfitstotheexperimentaldata.
tions). These observations suggest that neither dissolved
oxygennorbacteriawereresponsibleforAs(III)oxidation. ratio.However,itisgenerallyassumedthatFe(III)insoils
with a lower Fe(III) content than lake sediments does not
3.3. Aquifer material properties oxidize As(III) significantly (Oscarson et al., 1981b). The
specific surface areas, and the reductively extractable iron
Previous studies cited abovehaveshownthat the extent and manganese contents of the aquifer materials used in
and kinetics of As(III) oxidation may be attributed to the this study are reported in Table 1. These materials pos-
oxidized solid-phase manganese content of the material. sessed relatively similar specific surface areas
Adsorption of As(III), however, is largely controlled by (0.600±0.156m2g(cid:1)1), and reductively extractable iron
the oxidized iron content of the mineral (Deschamps contents(22.16±3.10lmolm(cid:1)2).However,thereductive-
et al., 2003). Oxidation of As(III) by Fe(III) hydroxide in lyextractablemanganesecontentsvariedbyafactorofsix,
lake sediment has been reported previously (De Vitre ranging from 0.59lmolm(cid:1)2 for R23AW material to
et al., 1991), possibly because of the high Fe(III):As(III) 3.36lmolm(cid:1)2 for F168-15 material (Table 1).
540 A.Amirbahmanetal.70(2006)533–547
3.4. Dissolved Mn(II) (cid:1)SOHþH AsO kf;a$d=kr;ad(cid:1)SOAsO H þH O ð1Þ
3 3 2 2 2
TheprocessofoxidationofAs(III)bymanganeseoxides where kf,ad and kr,ad are the forward and reverse rate con-
leadstotheproductionofMn(II)(Mooreetal.,1990;Scott stants for As(III) interaction with the adsorptive sites,
andMorgan,1995;Manningetal.,2002;Tournassatetal., respectively. It is assumed that As(III) oxidation does not
2002). However, dissolved manganese concentrations were take place at the surface of the adsorptive sites.
lessthan0.1lM(limitofquantification)inallexperiments. The processes involving the interaction between arsenic
Experimental studies with synthetic Mn(IV) oxides have speciesandthesurfacemanganeseoxidesareshowngener-
shown that dissolved manganese concentrations remained ically in Fig. 6. To obtain mass and charge balance, the
very low at pH values above approximately 5.8 (Scott reactions are shown in two dimensions. Assuming that
andMorgan,1995).Significantconcentrationsofdissolved Mn(IV) oxide is the primary oxidative site on the surface
manganese in oxic groundwater from the Cape Cod site of the aquifer material, the reaction scheme in Fig. 6 de-
have only been observed in samples with pH values below picts the formation of a As(III) precursor surface complex
approximately 5.0 (Savoie and LeBlanc, 1998). These re- with forward and reverse adsorption rate constants kf1,ox
sults suggest that Mn(II) binding to manganese oxide and and k , followed by electron transfer from the surface-
r1,ox
otherconstituentsofthesolidmaterialissufficientlyexten- bound As(III) to the Mn(IV) center with a rate constant
sive to maintain very low dissolved manganese concentra- of k , and the release of As(V) into solution with a rate
2,ox
tions over the pH range examined in this study. constant of k3,ox.
Ourmodeldoesnotconsidertheregenerationoftheoxi-
3.5. Conceptual model for arsenic–surface interactions dative surface sites. This assumption is reasonable for this
systemconsideringthattheproducedAs(V)concentrations
Similartothepreviousobservationsregardingthekinet- reachnearlysteady-statevalues(Figs.1–5),whereassignif-
ics of As(III) removal by oxidized manganese minerals icant regeneration of the surface oxidative sites would
(Scott and Morgan, 1995; Chui and Hering, 2000), the bring about a continual increase in the produced As(V)
kinetics of aqueous As(III) removal by the aquifer materi- concentration. One explanation for the lack of regenera-
als used here (Figs. 1–5) do not show a simple first-order tionofthesurfaceoxidativesitescouldbeduetotheasso-
behavior. Conceptually, this suggests that As(III) removal ciation of the produced Mn(II) with the surface Mn(IV)
and oxidation by aquifer material may be understood as a oxide, which would mask the surface sites from further
multi-stepsurfaceprocessinvolving(1)transportofAs(III) reaction.ScottandMorgan(1995)haveindeedshownthat
tothesurface,(2)As(III)adsorptionontothesurfacesites, theproducedMn(II)remainsassociatedwiththesurfaceof
(3) oxidation of As(III) to As(V) at the surface, (4) release birnessite to a significant extent, especially at pH values
ofproducedAs(V)fromthesurface,and(5)As(V)adsorp- above 4. Some As(V) may also remain on the manganese
tion to other surface sites following its production and re- oxide surface following its production, as suggested by
lease into solution. Adsorption of As(III) involves the spectroscopic studies (Manning et al., 2002; Deschamps
displacement of surface-bound OH(cid:1) and H O species via et al., 2003; Foster et al., 2003), inhibiting further As(III)
2
a ligand exchange mechanism, whereby the As(III) anion interaction with that oxidative site. Another explanation
forms an inner-sphere complex at the mineral surface
(Manning et al., 1998; Nesbitt et al., 1998; Arai et al.,
2001).
Mn O Mn OH Mn O Mn OH
The As(III) adsorption process may further be charac-
terized by the surface coordination of this species with O O k O O
f1,ox
thepurelyadsorptivesiteswherenoAs(III)oxidationtakes Mn O Mn OH + H3AsO3 Mn O Mn O AsIIIO2H2+ H2O
k
r1,ox
place,andwiththeoxidativesiteswhereAs(III)isoxidized O O O O
to As(V). The kinetic data suggest that within the time
Mn O Mn OH Mn O Mn OH
frame of this study, As(III) adsorption is a relatively slow
process with As(III) disappearance from solution having
a half-life in the order of several hours to several days Mn O Mn OH Mn O Mn OH
depending on the aquifer material (Figs. 1–5). Previous O OH O OH
k k
studies of oxyanion adsorption on Cape Cod materials 2,ox Mn O MnII O AsVO3H2 3,ox Mn O MnII+ + H2AsO4-
have found that approximately 3 days is required to reach
O OH O OH
equilibrium (Stollenwerk, 1995). Thus, on the timescale of
Mn O Mn OH Mn O Mn OH
these experiments, the rate of change of dissolved and ad-
sorbed As(III) must be taken into consideration. The pre- Fig. 6. Schematic showing the reversible adsorption of As(III) onto the
cise nature of surface complex formation of As(III) with oxidative surface Mn(IV) sites, followed by electron transfer and As(V)
release.k andk aretheforwardandbackwardrateconstantsfor
the aquifer material surface sites is unknown, and, there- f1,ox r1,ox
the adsorption and desorption of As(III), respectively, k is the rate
2,ox
fore, adsorption of As(III) on non-oxidative sites is de-
constantfortheelectrontransferreaction,andk istherateconstantfor
3,ox
scribed using a generic site, here abbreviated as (cid:1)SOH, detachmentofAs(V).
As(III)oxidationbyaquifermaterial 541
may bethe lackof reactivityof thenewly exposedMn(IV) approach is to develop the simplest possible model with
metal center at the solid–solution interface. the least number of fitting parameters that describes the
Surface spectroscopic and proton balance studies sug- main characteristics of solute–surface interaction in a
gest that As(III) oxidation by synthetic birnessite proceeds chemically heterogeneous system (Davis et al., 1998).
via a two-step process involving formation of Mn(III) as The reactions depicted in Eqs. (1) and (2), and in Fig. 6
the intermediate, followed by further reduction of Mn(III) were used to model the experimental data, with the corre-
to Mn(II) at pH 4 (Nesbitt et al., 1998; Tournassat et al., sponding equilibrium and rate constants as the fitting
2002).Wehavenotincludedthetwo-stepoxidationprocess parameters. To reduce the number of fitting parameters,
involving production of the intermediate Mn(III) in the certainsimplifyingassumptionsweremadeinthetreatment
proposed model, since the reaction stoichiometry with re- ofthekineticdata.Forchemicalreactionsinseries,suchas
spect to As(III) and surface Mn(IV) species remains the theonesdepictedinFig.6,whentheintermediatecomplex-
same for the two models. es(i.e.,adsorbedAs(III)andAs(V)complexesattheoxida-
It is assumed that in the multi-step reaction of As(III) tivesites)aresoreactivethattheydonotaccumulateatan
with MnO , the electron transfer and As(V) release pro- appreciable level compared to the reactants and the prod-
2
cesses are considerably faster than As(III) adsorption ucts, steady-state condition with respect to these species
(Tournassat etal., 2002); i.e., the As(III)precursor surface may be assumed (Espenson, 1995). The adsorbed As(III)
complex is very reactive. This assumption may be justified and As(V) concentrations at the surface of the oxidative
basedontheobservationthatAs(V)interactionwithnatu- sites cannot be directly measured in this system. However,
ral surfaces is controlled by its adsorption to Fe(III) based on the results of several previous studies of As(III)
hydroxides compared to Mn(IV) oxides (Deschamps oxidation by synthetic MnO minerals, we have assumed
2
et al., 2003). The presence of negligible concentrations of that the concentration of As(III) precursor surface com-
the produced aqueous As(V) species in this study, there- plex, ((cid:1)MnOAs(III)O H ; Fig. 6) is relatively small at any
2 2
fore, suggests that following its production, this species given time, and as such, reaches a steady-state value. Set-
interacts with the aquifer material extensively and rapidly. ting the formation rate of the As(III) precursor surface
Asdiscussedinthefollowingsection,undertheexperimen- complex equal to zero (i.e., d½(cid:1)MnOAsðIIIÞO2H2(cid:2)¼0), and solv-
tal conditions of this study, this interaction may be repre- ingforthesteady-stateconcentratiodtnof(cid:1)MnOAs(III)O H
2 2
sented as an equilibrium relationship with K as the
ad species from the set of reactions in Fig. 6, the formation
equilibrium constant,
rate of the aqueous As(V) species may be obtained as,
(cid:1)SOH+H AsO (cid:1) () (cid:1)SOAsO H(cid:1)+H O; K ð2Þ
2 4 3 2 ad d½H AsO (cid:1)(cid:2)
2 4 ¼k0 ½(cid:1)MnOH(cid:2)½H AsO (cid:2); ð3Þ
dt ox 3 3
3.6. Development of a macroscopic rate law for As(III)
where, the apparent rate constant k0 ¼ kf1;oxk3;ox if
oxidation k (cid:4)k , and k0 ¼ kf1;oxk2;ox if k (cid:4)okx k.r1T;oxhþek3;osxtea-
2,ox 3,ox ox kr1;oxþk2;ox 3,ox 2,ox
A rigorous treatment of the reactive functional groups dy-state assumption also implies that k +k (or
r1,ox 2,ox
of natural surfaces with mixtures of mineral phases and k )(cid:4)k , which suggests a slow rate of As(III)
3,ox f1,ox
their electrical double layer properties is a complicated adsorptionontotheoxidativesitescomparedtotheAs(III)
task. Part of this difficulty arises from the uncertainties in transformation and As(V) release rates. The steady-state
the estimation of the surface densities of reference mineral assumption simplifies the kinetic formulation considerably
phases and the associated surface coatings that are avail- by combining four rate constants (Fig. 6) into one, k0 .
ox
able for reaction with the solution phase species (Davis Obtaining a quantitative fit of experimental adsorption
etal., 2004). Here, we have used theGeneralized Compos- data over a wide range of chemical conditions on Cape
ite (GC) approach where solute–surface interactions are Cod aquifer materials with semi-empirical surface com-
described by mass balance and reaction thermodynamic plexation models required considering strong and weak
expressions containing ‘‘generic’’ surface functional bindingsites(Davisetal.,1998).Thislikelyreflectsthevar-
groups. The relevant stoichiometry and equilibrium con- iation in the adsorption site energies on complex mineral
stants for these expressions are estimated by fitting experi- surfaces. For the As(III) concentration range used here,
mentaldatafortheentiremineralassemblage(Davisetal., best fitsto thedatawere obtained byconsideringtwo sites
1998; Westall et al., 1998). Even though the GC approach with differing interaction energies for both the purely
is generally developed for equilibrium solute adsorption, adsorptivesites and theoxidativesites. Since thisstudy in-
we have used it to develop a set of mass action equations volves adsorption and oxidation kinetics of arsenic species
to model the kinetics of arsenic interaction and redox inthepresenceofasolidsurface,thesesitesaretermedfast
transformation at the surface of aquifer materials. The and slow, in analogy to the strong and weak sites used for
GCapproachalsoneglectsanysurfaceelectrostaticcorrec- adsorption equilibrium. We have further assumed a con-
tionfactors,andassuch,requiresfewerfittingparameters, stant total surface density for the adsorptive sites at 2.31
andiscomputationallyefficientwhenincorporatedinreac- sites nm(cid:1)2 (equivalent to 3.84lmol m(cid:1)2) for all aquifer
tive solute transport models. The general goal of this materials as proposed by Davis and Kent (1990). This
542 A.Amirbahmanetal.70(2006)533–547
Table2
Equilibriumandkineticexpressionsandconstantsusedtomodelthedataa
Reactionb korlogKb,c Ref.
HAsO () H++HAsO(cid:1) (cid:1)9.15 1
3 3 2 3
HAsO () H++HAsO(cid:1) (cid:1)2.3 1
3 4 2 4
HAsO () 2H++HAsO2(cid:1) (cid:1)9.46 1
3 4 4
HAsO () 3H++AsO3(cid:1) (cid:1)21.11 1
3 4 4
(cid:1)SfOH+HAsO M(cid:1)SfOAsOH+HO kf ¼2:43(cid:3)10(cid:1)2M(cid:1)1s(cid:1)1 2
3 3 2 2 f;ad
kf ¼5:25(cid:3)10(cid:1)7s(cid:1)1
r;ad
(cid:1)SsOH+HAsO M(cid:1)SsOAsOH+HO ks ¼6:85(cid:3)10(cid:1)3M(cid:1)1s(cid:1)1 2
3 3 2 2 f;ad
ks ¼4:72(cid:3)10(cid:1)5s(cid:1)1
r;ad
(cid:1)MnfOH+HAsO fiHAsO +otherproductsd kf0 ¼6:28(cid:3)10(cid:1)1M(cid:1)1s(cid:1)1 2
3 3 3 4 ox
(cid:1)MnsOH+HAsO fiHAsO +otherproductsd ks0 ¼1:25(cid:3)10(cid:1)2M(cid:1)1s(cid:1)1 2
3 3 3 4 ox
(cid:1)SfOH+HAsO(cid:1) () (cid:1)SfOAsOH(cid:1)+HO 6.37 2
2 4 3 2
(cid:1)SsOH+HAsO(cid:1) () (cid:1)SsOAsOH(cid:1)+HO 5.06 2
2 4 3 2
(cid:1)MnOH () (cid:1)MnO(cid:1)+H+ (cid:1)6.2 2
(forbothoxidativesites)
1,SmithandMartell(1976);2,thiswork.
a ConstantsforallsurfacereactionsestimatedusingdataforF168-15material.
b ‘‘f’’and‘‘s’’inthesuperscriptsrefertofastandslowsites,respectively,and‘‘f’’and‘‘r’’inthesubscriptsrefertoforwardandreverserateconstants,
respectively.
c ConstantsgivenforI=0.
d SeeFig.6andthetextfortheexplanationofthesereactions.
choice is in keeping with the development of a self-consis- taneously using the four different initial As(III) concen-
tentthermodynamicdatabasethatmaybeappliedtomod- trations by UCODE. The values of the calibrated
el solute adsorption onto natural composite mineral equilibrium and rate constants for the F168-15 material
surfaces. are reported in Table 2 and the fitted site densities for
Table 2 shows the set of reactions used to describe the the oxidative sites are reported in Table 3. For the non-
kinetics of arsenic interaction with the aquifer material. oxidative site densities, the best fit to the ktinetic data
Thesereactionsinvolvefiverateandequilibrium constants
for each fast and slow site that need to be optimized to
Table3
model the data. Kinetic data from the most reactive mate-
Fitteddensitiesfortheoxidativesites
rial (F168-15) were fitted with the constants in Table 2, as
Aquifermaterial (cid:1)MnfOH(lmolm(cid:1)2) (cid:1)MnsOH(lmolm(cid:1)2)
wellaswiththetotalsurfacedensityforbothfastandslow
F168-15 0.129 0.665
oxidativesites.TheF168-15materialpossessedthehighest
F168-20 0.053 0.303
reductively extractable manganese content (Table 1), and
F168-25 0.032 0.106
the fastest As(III) removal and As(V) production rates of F415-19 0.049 0.134
all materials (Fig. 1). The constants obtained were then R23AW 0.013 0.046
usedtosimulateAs(III)oxidationkineticsbyotheraquifer
materials. In this case, however, since each material had a
different reactivity with respect to As(III) oxidation, and a 1.0
y
different reductively extractable manganese content (Table sit
n
1), the oxidative site densities were used as the only fitting e de 0.8 total
parameters for individual materials; i.e., the rate and equi- sit
librium constants were the same for all materials, and as e slow
described above, the total density of adsorptive sites were urfac-2)m 0.6
estimated from the specific surface area measurements. e s mol
The approach used here, where the parameters were esti- ativµ( 0.4
mated based on only one material and used for materials d
xi
not used originally in the model development, is intended ed o 0.2
to validate the proposed model for different conditions. el fast
d
The reaction parameters were determined by fitting the o
m
model to the observed aqueous As(III), and adsorbed 0.0
0 1 2 3 4
As(III) and As(V) data for F168-15 material that was
reacted with initial As(III) concentrations that ranged extractable Mn content (µmol m-2)
from 30 to 500lM (Fig. 1). Initial parameter values were
Fig. 7. Reductively extractable manganese content vs. the modeled
estimated from observations at both 30 and 500lM, and oxidative surface site density. (d) Total density of the fitted sites, (s)
then all of the reaction parameters were estimated simul- fitteddensityoftheslowsites,and(n)fitteddensityofthefastsites.