Carbonate Ions and ArsenicDissolution by GroundwaterM Y O U N G - J I N K I M , †

J E R O M E N R I A G U , * , † A N DS H E R I D A N H A A C K ‡

Department of Environmental Health Science, School ofPublic Health, University of Michigan, Ann Arbor,Michigan 48109, and Water Resources Division,U.S. Geological Survey, Lansing, Michigan 48911

Samples of Marshall Sandstone, a major source ofgroundwater with elevated arsenic levels in southeastMichigan, were exposed to bicarbonate ion under controlledchemical conditions. In particular, effects of pH andredox conditions on arsenic release were evaluated. Therelease of arsenic from the aquifer rock was strongly relatedto the bicarbonate concentration in the leaching solution.The results obtained suggest that the carbonation ofarsenic sulfide minerals, including orpiment (As2S3) andrealgar (As2S2), is an important process in leaching arsenicinto groundwater under anaerobic conditions. The arseno-carbonate complexes formed, believed to be As(CO3)2

-,As(CO3)(OH)2

-, and AsCO3+, are stable in groundwater. The

reaction of ferrous ion with the thioarsenite fromcarbonation process can result in the formation ofarsenopyrite which is a common mineral in arsenic-richaquifers.

IntroductionThe chemistry of arsenic in groundwater remains murkydespite the large number of publications on the topic (1-8).In many parts of the world, elevated levels of arsenic arefound in groundwaters that are mostly reducing, as evidencedby low redox potentials and dissolved oxygen concentrations(typically < 0.5 mg L-1), high concentrations of dissolvediron and manganese and a preponderance of As(III) in thesoluble phase. The groundwater environments are oftencharacterized by high alkalinity. The most common arsenic-bearing minerals in the host rocks, believed to be the sourceof the arsenic in groundwater, include arsenic-rich (arsenian)pyrite and various arsenic sulfides and sulfosalts (2, 6, 7,9-11). Although arsenic can be released from these thio-mineral phases by an oxidation process (12-15), there is noknown abiotic mechanism for leaching the arsenic from theseminerals in the host rocks under reducing conditions.

The purpose of this investigation is to ascertain the likelyfactors and mechanisms involved in leaching arsenic fromthe fresh core samples of Marshall Sandstone (MSS), theprincipal bedrock aquifer in southeastern Michigan. Ground-water in MSS is reducing, has elevated levels of arsenic, andshares most of the characteristics noted above (16). Similargeochemical conditions prevail wherever elevated levels ofnaturally occurring arsenic are found in groundwater ofmidwestern United States (1, 17, 18). The origin of the

anomalous arsenic in groundwater in this region of thecountry has been a matter for much conjecture andspeculation.

In this report, the role of bicarbonate in leaching arsenicinto groundwater was investigated by conducting batchexperiments using core samples of MSS and NaHCO3

solutions. The effects of pH and redox conditions on arsenicdissolution were evaluated. Nickson et al. (7, 8) have recentlyproposed that arsenic in groundwater derives from reductivedissolution of As-rich iron oxyhydroxides that exist asdispersed phase in the aquifer rock. The sorbed arsenic isreleased and reduced during the dissolution of the oxy-hydroxides. The relationship between dissolved As and HCO3

-

was assumed to be indirect and attributed to the catenatedeffect of microbial reduction of organic matter (CH2O):

The present study suggests that the leaching of As intogroundwater is driven by direct interaction between HCO3

-

and As minerals in the aquifer rocks.

Materials and MethodsEach apparatus and bottle utilized in the experiment waswashed using nitric acid. All reagents were of analytical grade,and all solutions were prepared with deionized water. Arsenicand iron in solution were determined by a graphite furnaceatomic absorption spectrophotometer (GFAAS, Perkin-Elmer4100 ZL), and an electrodeless discharge lamp was used forarsenic analysis. During the GFAAS analysis, a matrix modifierof 5 µg of Pd and 3 µg of Mg(NO3)2 was used for each 20 µLof sample. The detection limit of arsenic analysis was 3 µg/L.A flame atomic absorption spectrophotometer (VarianTechtron Model 1200) was used to determine total ironconcentration in acid-digested solutions of rock samples.Ion chromatography (IC) used was made by Alltech andconnected to a conductivity detector. The redox potential(Eh) was determined using a platinum electrode (Ag-AgCl)and reference hydrogen ion electrode (Orion Co.). The pHof each solution was measured with an Orion Model 250ApH meter. A dissolved oxygen (DO) meter (YSI 5000) wasused to measure DO in solution.

For the speciation of As(III) and As(V), ion exchangemethods of Ficklin (19) and Grabinski (20) were adapted andmodified (21). Water samples were acidified using HCl to pH4-7 and allowed to pass through a column packed with stronganion-exchange resin (AG1-X8, 100 to 200 mesh chlorideform), followed by 0.1 M HCl as eluent. It has been shownthat adjustment of pH to 4-7 does not affect the As(III)/As(V) ratio in samples (16). The As(III) species from the samplewas in the pass-through solution, and the As(V) species wasadsorbed into the resin and later eluted with 0.1 M HCl.

A new well was drilled in Bad Axe, Huron County, MI inOctober 1997 by the U.S. Geological Survey (USGS), and coresamples were taken at different depths between 50 and 350ft. This interval is in the Marshall Sandstone, the primaryaquifer for potable water in southeast Michigan (22). Thelocation was chosen because it was close to existing wellswith high dissolved arsenic concentration (total As ) 167-278 µg/L, As(III) ) 73-92%). Immediately after the coresamples (4′′ diameter) were collected, they were placed insealed plastic bags filled with nitrogen. Subsequently, theMSS samples were crushed and placed in a freeze-drier (VirtisFreezemobile model 12 SL) to dry for 3-4 days. The freeze-

* Corresponding author phone: (734)936-0706; fax: (734)764-9424;e-mail: jnriagu@umich.edu.

† University of Michigan.‡ Water Resources Division, U.S. Geological Survey.

4FeOOH + CH2O + 7H2CO3 f

4Fe2+ + 8HCO3- + 6H2O (1)

Environ. Sci. Technol. 2000, 34, 3094-3100

3094 9 ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 34, NO. 15, 2000 10.1021/es990949p CCC: $19.00 2000 American Chemical SocietyPublished on Web 06/21/2000

dried samples were sieved through a 0.35 mm mesh andhomogenized by thorough mixing. The samples were storedin a freezer until further analysis.

To determine total arsenic concentration, the pulverizedrock samples were digested using concentrated nitric acid(16). One gram of dry core sample and 10 mL of concentratednitric acid were placed in a 100 mL glass beaker. Then thebeaker was covered with a watch glass and put on a heaterwhich was set at 100 °C for 1 h. After the beaker was cooled,approximately 30 mL of deionized water was added to thebeaker. The suspended mixture was filtered through a 0.45µm membrane. Finally the filtrate was transferred to a 100mL volumetric flask and brought up to the volume usingdeionized water. The same procedure was performed witha blank (nitric acid without core sample) and a standardreference material (2709, San Joaquin Soil) in each batch ofdigestion. The digestion procedure was effective in extractingAs from rock samples, showing 89-98% recoveries ofstandard material. The range of total arsenic concentrationsin rock samples was from 0.8 mg/kg to 72.1 mg/kg. Becausethe mass of each sample at different depth was not enoughto conduct the entire experiment with each sample separately,10 core samples with high arsenic concentrations (>9 mg/kg) were homogenized by mixing them together. The arsenicconcentration of the mixture was 26.4 mg/kg. Unlessotherwise specified, all further leaching experiments wereperformed with this mixed MSS sample.

Determination of Effects of Major Ions on Leaching ofArsenic. To determine which major ions most efficientlycause arsenic leaching, deionized water, groundwater, andsynthetic solutions containing the major ions found ingroundwater of study area, namely Ca2+, K+, Mg2+, Na+, Cl-,HCO3

-, and SO42-, were prepared. Table 1 shows average

concentrations of major and minor ions in groundwater ofsoutheast Michigan. Groundwater used in this leachingexperiment was collected from a local well in Ann Arborknown to contain negligible amounts (<3 µg/L) of arsenic.

Leaching tests were conducted using water (deionizedwater and groundwater) and 0.1 M solutions of KCl, Na2SO4,MgSO4, CaSO4, NaHCO3, KHCO3, and FeCl3. The level ofarsenic concentration in each starting solution was negligible(<3 µg/L). As shown in Table 2, the pH of test solutions ofdeionized water, groundwater, KCl, Na2SO4, and MgSO4 wereneutral, those of CaSO4, NaHCO3, and KHCO3 were basic,and that of FeCl3 was acidic.

One gram of MSS sample and 20 mL of each solutionwere placed in three polyethylene bottles, and each batchconsisting of 27 bottles were shaken using a platform shaker.One bottle of each solution was removed from the shakerafter each of the time intervals (4 h, 1 day, and 3 days). Allleaching tests were run under atmospheric conditions at roomtemperature. The sample mixtures were filtered through a0.45 µm membrane, and the arsenic concentrations in thefiltrates were determined by GFAAS.

Arsenic Leaching in Different NaHCO3 Concentrations.From the leaching test described above, it was found thatNaHCO3 leached arsenic from core samples most efficiently.Additional experiments were conducted to further explorethe role of HCO3

- in arsenic dissolution. Since the pH of 0.1M NaHCO3 was 8.5, it is conceivable that hydroxide ion wasprimarily involved in leaching arsenic rather than thebicarbonate ion. After the pH of 0.1 M NaCl and 0.1 M MgSO4

solutions was adjusted to 10.1 and 8.8, respectively, the sameleaching test described above was performed. Subsequently,varying concentrations of NaHCO3, ranging from 0.02 to 0.6M, were used to test the rate of arsenic leaching. The testswere performed for different time intervals (4 h, 1 day, and3 days) using the same experimental method described above.

Arsenic Leaching under Oxic/Anoxic Conditions. Toinvestigate how the rate of arsenic leaching changes underoxic or anoxic condition, three kinds of solution wereprepared as follows: experiment #1, 0.04 M NaHCO3 madein deionized water under air; experiment #2, 0.04 M NaHCO3

made in deionized water under nitrogen; and experiment#3, 0.04 M NaHCO3 made in groundwater under nitrogen.The last experimental system should approximate the air-free conditions in groundwater.

Oxic conditions were maintained by preparing the solutionin air-saturated deionized water under standard atmosphericconditions. MSS sample (2.4 g) and 20 mL of air-saturatedsolution (experiment #1) were placed in eight polyethylenebottles.

To provide anoxic conditions for experiments #2 and #3,it was necessary to set up the system in a glovebag undernitrogen. Each experimental charge consisted of 20 mL ofdeaerated NaHCO3 solution and 2.4 g of rock sample in aSchlenk tube. A total of 24 bottles was shaken using a platformshaker, and one bottle of each experimental solution wasremoved from the shaker after each of the time intervals(from 4 h to about 4 days). The chemical characteristics ofthe test solution including temperature, pH, DO, and redoxpotential (Eh) were measured. The mixtures were filteredthrough a 0.45 µm membrane, and the arsenic concentrationsin the filtrates were determined by GFAAS.

Arsenic Leaching at Different pH. The effect of H+ on thearsenic leaching was tested in the pH range of 2-10. Sincethe pH of 0.04 M NaHCO3 solution is 8.5, the pH of the solutionwas adjusted using concentrated HCl or NaOH. The coresample (2.4 g) and 20 mL of each pH-adjusted 0.04 M NaHCO3

solution were placed in polyethylene bottles, and the bottleswere shaken using a platform shaker. One bottle of eachsolution was removed from the shaker after each of the timeintervals (4 h, 1 day, and 3 days). All leaching tests were rununder air at room temperature.

Relation between Total Arsenic Concentration in CoreSamples and Arsenic Leaching. Eight core samples withdifferent total arsenic concentrations, ranging from 1.6 to70.7 mg/kg, were used to investigate the release of arsenicfrom MSS samples. Six grams of sample were reacted with50 mL of 0.04 M NaHCO3; 20 g of samples was used whenarsenic concentrations were low (<12 mg/kg). Schlenk tubescontaining 6 g or 20 g rock sample and deaerated NaHCO3

solutions were prepared, and the entire leaching experimentwas run under nitrogen using the same procedure describedabove. After filtration through a 0.45 µm membrane, the

TABLE 1. Average Concentrations of Major and Minor IonsFound in Groundwater of Southeast Michigan

ion concn (M) ion concn (M)

HCO3- 5.0 × 10-3 NO3

- 1.4 × 10-5

Ca2+ 1.8 × 10-3 Fe2+ 1.2 × 10-5

Na+ 1.7 × 10-3 NH4+ 1.1 × 10-5

Mg2+ 1.0 × 10-3 Br- 1.2 × 10-6

Cl- 8.7 × 10-4 HS- 8.8 × 10-7

SO42- 3.9 × 10-4 PO4

3- 8.8 × 10-7

K+ 1.6 × 10-4 Mn 6.0 × 10-7

F- 3.0 × 10-5 H+ 4.0 × 10-8

Fe (total) 1.8 × 10-5

TABLE 2. pH of Deionized Water, Groundwater, and 0.1 MSolutions Used in Leaching Experiments

solution pH solution pH

deionized water 6.08 CaSO4 8.76groundwater 7.03 NaHCO3 8.50KCl 6.04 KHCO3 8.26Na2SO4 6.51 FeCl3 2.16MgSO4 6.21

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concentrations of arsenic in the leachates were determinedby GFAAS. Arsenic species in the leachates were separatedusing an ion exchange method (21).

Stability of Arsenic Species in Rock-NaHCO3 SolutionMixture. The influence of NaHCO3 and MSS samples on thestability of arsenic species was evaluated under differentconditions. Six sets of experiments were designed as shownin Table 3. In every experiment, 40 mL of 0.04 M NaHCO3

was used. In experiments on rock samples, 6 g of sample(total As concentration ) 9.1 mg/kg) was suspended in aNaHCO3 solution, and arsenic was spiked to bring theconcentration to about 170 µg/L As. In experiments undernitrogen, the method described above was used. For experi-ments under air, the solution was completely aerated. Aftershaking the suspended mixtures for 3 days, they were filteredthrough a 0.45 µm membrane. The pH of the filtrate wasadjusted with HCl, and the arsenic species were separatedusing an ion exchange method (21). The concentrations oftotal arsenic and arsenic species were determined by GFAAS.

Determination of Arseno-Carbonate Complexes UsingIon Chromatography (IC). Formation of arseno-carbonatecomplexes was evaluated with pure arsenic minerals byplacing 0.25 g each of orpiment (As2S3) and arsenic trioxide(As2O3) in 100 mL of N2-saturated 0.04 M NaHCO3 solutionto maintain anaerobic conditions. Other orpiment solutionwas prepared by reacting 0.25 g of As2S3 in 100 mL of 3 × 10-4

M NaOH solution, under anaerobic condition. After all themixtures were stirred for 1 day, they were filtered througha 0.45 µm membrane. In addition, As(III) and As(V) standardsolutions (2 × 10-3 M) were prepared from NaAsO2 and Na2-HAsO4‚7H2O, respectively.

The separation of As(III), As(V), and other arseniccomplexes was performed using an Ion Chromatography (IC,Waters) equipped with IC-Pak Anion HC (150 × 4.6 mm,10-µm particle size) analytical anion-exchange columnpreceded by a guard column. The stationary phase consistedof trimethylammonium functionalized groups on poly-methacrylate. The 0.85 mM NaHCO3 and 0.9 mM Na2CO3

(HPLC grade) (pH ) 9.9) were used as mobile phase at a flowrate of 1.4 mL/min. Each sample was injected into a 1 mLloop directly using a syringe. Samples that were analyzedusing IC included As2S3 leachate in NaHCO3; As2O3 leachatein NaHCO3; As2S3 leachate in NaOH; As(III) solution; andAs(V) solution.

ResultsArsenic was not significantly leached out of MSS samplesafter 3 days using either deionized water or groundwater (<5µg/L). By contrast, NaHCO3, KHCO3, and FeCl3 solutionsextracted arsenic from the rock samples most efficiently; 5.9%of arsenic in rock sample was leached with 0.1 M NaHCO3,4.6% with 0.1 M KHCO3, and 1.9% with 0.1 M FeCl3 after 3

days incubation (Figure 1). These results indicate thatbicarbonate ion and ferric ion were most effective in arsenicleaching. The leaching rate with NaHCO3 was about 25%higher than that with KHCO3 (Figure 1). The arsenic leachingrate increased with reaction time in NaHCO3, KHCO3, CaSO4,and MgSO4 solutions. On the other hand, in leachingexperiment with other reagents, arsenic concentrations inthe leachate remained fairly constant (( 20% of reportedvalues) after 1-day incubation (Figure 1).

A number of researchers have tried to extract arsenic inpolluted soils using various aqueous media including NaH-CO3 (23, 24). Pantsar-Kallio and Manninen extracted about10% of the total arsenic from As2O5-contaminated soil using0.1 M NaHCO3 (23). McLaren et al. (24) used NaHCO3 toextract arsenic in a sequential fractionation scheme anddescribed NaHCO3 extractable arsenic as “nonexchangeablebut readily labile As associated with soil mineral surfaces”.They found that a relatively small proportion of arseniccontent was extracted using NaHCO3. The studies with soilswere done under aerobic conditions.

Experiments using pH-adjusted NaCl and MgSO4 solutionsdemonstrated that bicarbonate ion in 0.1 M NaHCO3 solution(pH 8.5) was primarily involved in arsenic dissolution andthat OH- ion did not affect arsenic dissolution to anysignificant degree. Even though the pH of NaCl (pH ) 10.1)and MgSO4 (pH ) 8.8) solutions were higher than that ofNaHCO3 solution, the amount of arsenic leached was small(<10 µg/L) compared to what was obtained using NaHCO3

(78.2 µg/L). The result using CaSO4 at pH 8.8 (<5.5 µg/L) alsoshowed that OH- itself was not an effective leaching agentunder moderately alkaline conditions.

The amounts of arsenic leached out of the rock sampleswere noticeably dependent on NaHCO3 concentration andincreased with reaction time for each concentration (Figure2). It should be noted that starting NaHCO3 solutions thathad the initial pH values adjusted to 8.5 ( 1.0 showed arelatively constant pH during incubation process. After a3-day incubation, 1.5-14.8% of the arsenic content of therock samples was leached: 1.5% in a 0.02 M solution and14.8% in a 0.6 M solution.

To evaluate arsenic leaching under oxic/anoxic conditions,experimental conditions including temperature, pH, DO, andredox potential were kept relatively constant ((20% ofreported values). For all three experimental sets, temperatureand pH were 20 °C and 8.5, respectively. On the other hand,DO and redox potential values were different depending onoxic/anoxic conditions: DO values were 8.2 (experiment #1)

TABLE 3. Experimental Design and Results to Assess theStability of As(III) and As(V) in 0.04 M NaHCO3 Solution afterIncubation

experimentresulta

experi-ment

spikedAs rock

N2 orair Asb As(T)c As(V)c

#1 no yes air 15.3 ( 2.9 10.0 ( 3.1#2 no yes N2 19.2 ( 3.5 16.8 ( 2.6#3 As(III) no N2 167.0 167.4 ( 2.7 43.7 ( 2.3#4 As(V) no N2 170.2 169.8 ( 2.5 148.1 ( 4.2#5 As(III) yes N2 167.0 76.4 ( 5.1 65.6 ( 5.8#6 As(V) yes N2 170.2 91.3 ( 5.3 80.4 ( 4.9a Arsenic concentration unit ) µg/L. b Spiked arsenic concentration.

c Arsenic concentration after 3 days of reaction; As(T) ) total Asconcentration, mean ( SD is shown, As(III) ) As(T) - As(V).

FIGURE 1. Arsenic leaching rate in several solutions which includethe major ions in groundwater (DI ) deionized water, GW )groundwater).

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and 0.3 (experiments #2 and #3), and redox potential valueswere 145 mV (experiment #1), 12 mV (experiment #2), and-80 mV (experiment #3). The highest arsenic leaching ratewas obtained from the air-saturated solution (experiment#1), followed by the anoxic solution (experiment #2), and thelowest rate was from anoxic groundwater (experiment #3)(see Figure 3). The arsenic concentration increased sharplywithin 1 day and then more slowly in subsequent days.

The results of arsenic leaching at different pH values areshown in Figure 4. Significant arsenic leaching (2.6-16.4%of total arsenic) was found in the extreme pH ranges of <1.9and 8.0-10.4. Considering that HCO3

- is the major carbonatespecies between pH 6.3 and pH 10.3 and CO3

2- dominatesabove pH of 10.3 (the pKa values for carbonate system are6.3 and 10.3), the observation suggests that CO3

2- is the likelyprimary arsenic leaching agent. Even though CO3

2- seemedto be much more efficient than HCO3

- in arsenic leaching,HCO3

- was used for all leaching experiments in order toapproximate groundwater conditions. The pH range ofgroundwater in the newly drilled well was 7.1-7.3, and themost abundant carbonate species is expected to be HCO3

-.Relatively high arsenic leaching at low pH (<1.9) can beexplained by acidic leaching of arsenic rather than by therole of carbonate species.

After 3 days of incubating eight rock samples with differentarsenic contents (1.6-70.7 mg/kg) in 0.04 M NaHCO3

solutions under deaerated conditions, 0.3-4.0% of arseniccontents originally present was leached from the core

samples. A linear relation was found between arsenic contentsin core samples and those in leachates. These results suggestthat dissolved arsenic concentration in groundwater is relatedto the amount of arsenic in aquifer rocks.

The ion exchange separation method showed that themajor arsenic species in leachates obtained under nitrogenwas retained on the resin, i.e., was negatively charged. Thisspecies accounted for 95-100% of total arsenic in theleachates, regardless of the arsenic content of the rocksamples.

Experiments to identify the arseno-carbonate complexesusing IC showed the likelihood of their existence (Figure 5).In the chromatograms from As2S3 and As2O3 leachates inNaHCO3 solution (Figure 5a,b), the two peaks at elution timesof 8.7 and 10.4 min are unique and can be attributed toAs(CO3)2

- and As(CO3)(OH)2-. Two chromatograms (Figure

5c,d) show an identical sharp peak at elution time of 6.9 min,suggesting that As(III) is the major leached species from As2S3

in NaOH solution. As(V) was eluted at 47.2 min: thechromatogram for As(V) is not shown. The sharp peak atelution time of 14.6 min in Figure 5a is not identified, butit is not an arsenic peak. After collecting eluting samplesfrom IC every minute, they were analyzed for arsenic usingGFAAS: the results showed corresponding arsenic peaks at

FIGURE 2. Arsenic leaching rate in different NaHCO3 concentrationunder air.

FIGURE 3. Arsenic leaching rate over time under oxic/anoxicconditions: (9) 0.04 M NaHCO3 made in deionized water under air,([) 0.04 M NaHCO3 made in deionized water under nitrogen, and(2) 0.04 M NaHCO3 made in groundwater under nitrogen.

FIGURE 4. Arsenic leaching in solutions with different pH exposedto air.

FIGURE 5. IC chromatograms for (a) As2S3 leachate in NaHCO3, (b)As2O3 leachate in NaHCO3, (c) As2S3 leachate in NaOH, and (d)As(III) solution. * Y axis is not to scale. Chromatogram of each ofthe four solutions has been adjusted and is shown separately forclearer display of retention time.

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similar elution times. Further studies using other types ofinstruments such as C NMR and Raman spectroscopy toinvestigate the forms of arseno-carbonate complexes areunderway.

DiscussionIn soils and sediments, the behavior of arsenic is largelycontrolled by Fe, Mn, organic matter, redox reactions,hydrolysis, and diffusive transport (25-27). In the reducingenvironment of Marshall Sandstone (16), the concentrationof arsenic in groundwater should be controlled by mineralphases other than the oxyhydroxides of Fe and Mn. Arsenic-rich pyrite grains have been identified in Marshall Sandstone(22), and it is widely believed that the occurrence of arsenicin groundwater of southeastern Michigan and other parts ofthe world is associated with arsenian pyrite (7, 10, 28, 29).This view is an oversimplification. In the presence of sulfideion (common in groundwater in Michigan) (16), both As(V)and As(III) are quickly and reductively converted to insolublearsenic sulfide precipitates:

In anaerobic environments, arsenic can also form a largenumber of stable sulfosalts with differing S:As ratios whichcan be represented by the following generalized formula(30): mMI

2S‚nMIIS‚oMIAsS2‚pMII(AsS2)2 where MI and MII

refer to monovalent and divalent metals. The oxidation statefor arsenic in these minerals is +3 (as in trechmannite, AgAsS2

and in sartorite, Pb(AsS2)2). We posit that the role of arsenicsulfides and sulfosalts in controlling the levels of arsenic ingroundwater is much more important than is generallyrealized.

Experiments in aerated deionized water show increasedarsenic leaching from MSS samples in the presence ofbicarbonate ion (Figure 1) and that the amount of arsenicleached depends on the HCO3

- concentration (Figure 2).Nickson et al. (7, 8) reported that arsenic concentration wassignificantly correlated with bicarbonate concentration inanoxic groundwater. Other studies have shown that pyriteoxidation is affected (enhanced) by the presence andconcentration of HCO3

- (31, 32). The increase in pyriteoxidation is attributed to the formation of a pyrite surface-Fe(II)CO3 complex which facilitates the transfer of electronfrom Fe(II) to O2 or other electron acceptor (32). Unlike Fe3+

which functions as an electron acceptor, the role of HCO3-

in the oxidation of pyrite is indirect. Figure 1 shows that therate of oxidative leaching of arsenic out of the MSS samplesis higher in the presence of bicarbonate ion compared toferric ions.

In aerobic environments, there are therefore mechanismsfor oxidizing arsenian pyrite to release AsO3

3- or AsO43- or

their protonated species in the presence of bicarbonate orferric ions. Other arsenic sulfides can similarly be oxidizedand would add to the anomalous arsenic in the groundwater(15). Results of this study, however, show that bicarbonateion is equally effective in releasing arsenic from the rocksamples in deoxygenated deionized water and groundwater(Figure 3), suggesting that nonoxidative leaching process isinvolved. The oxidation state of arsenic in arsenopyrite is -1(33), and the major dilemma for proponents of arsenopyriteas the source of arsenic is that there is no mechanism forchanging the oxidation state to +3 in anaerobic groundwaterenvironment which generally contain sulfide and highconcentrations of ferrous and manganous ion. To overcomethis anachronism, some people have suggested that thearsenopyrite or arsenic-rich pyrite is oxidized first (in the

overlying or adjacent rock formations for instance), and thearsenate is then transported to and reduced in the aquifer.The problem with this explanation is that the decoupling ofarsenate and iron/manganese cycles in aerobic environmentsis not always feasible.

We propose that arsenic sulfides and sulfosalts, ratherthan arsenian pyrite, are the principal sources of arsenic inanaerobic groundwater environments of MSS. Besides themechanistic difficulties, support for the suggestion comesfrom the fact that there is no significant relationship betweenarsenic and iron concentrations in the rock samples (Figure6a) or between the arsenic and iron concentrations in theleachates (Figure 6b). We further suggest that the formationof stable arseno-carbonate complexes are responsible foreffective leaching of arsenic in anaerobic systems (Figures3 and 5). The dissolution reactions are envisaged as follows(using orpiment, As2S3, as an example):

That carbonate ion forms stable complex(es) with As(III) hasbeen known for a long time. An aqueous solution of As2O3

in K2CO3 was the Fowler’s solution, a common lotion in theold materia medica (30). Classical analytical chemistrytextbooks also pointed to the amazing solubility of arsenicsulfides at elevated pH in the presence of carbonate ions.

To confirm the forms and stability of arsenic in ourexperimental solutions, mixtures of MSS and 0.04 M NaHCO3

solution were spiked with either As(III) or As(V) and incubatedunder nitrogen or air. The results are shown in Table 3. Inexperiments #1 and #2, 1.1% and 1.4% of the original arseniccontent of the rock samples were leached under air and under

2H3AsO3 + 3H+ + 3HS- f As2S3(orpiment) + 6H2O (2)

2As2S3 + 2H+ + 4e- f 4AsS(realgar) + 2HS- (3)

FIGURE 6. Relation between Fe and As occurrence in MarshallSandstone and leaching solution.

As2S3 + HCO3- f As(CO3)+ + AsS2

- + HS- (4)

As2S3 + 2HCO3- f As(CO3)2

- + HAsS2 + HS- (5)

As2S3 + HCO3- + 2H2O f

As(CO3)(OH)2- + HAsS2 + HS- + H+ (6)

HAsS2 + H2O f HAsS2(OH)- + H+ (7)

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nitrogen respectively using 0.04 M NaHCO3. The major formof arsenic under both conditions had a negative charge(retained on ion-exchange resin). Formation of As(III)carbonate complex is suggested by the fact that the dominantform in leachate with deoxygenated water (experiment #2)was charged rather than the expected neutral H3AsO3. Theresults from experiments #3 and #4 show that arsenic specieswere relatively stable in 0.04 M NaHCO3 solution undernitrogen. Approximately 26% of As(III) was oxidized to As(V)and 13% of As(V) was reduced to As(III) under this conditionin 3 days of incubation. These observations suggest that theoxidation of As(III) could not have been primarily responsiblefor the charged species of experiment #2.

Results from the experiment with arsenic spike into rock-solution mixtures (experiments #5 and #6) were ratherunexpected (Table 3). Total arsenic concentration decreasedfrom 167.0 µg/L to 76.4 µg/L for As(III) spike and from 170.2µg/L to 91.3 µg/L for As(V) spike after 3 days of incubationunder nitrogen. This suggests that both As(III) and As(V)oxyanions are reactive with MSS, resulting in 46-54%retention of the applied arsenic. The inference argues againstthe two-site theory in which arsenic minerals are first oxidizedat one site and oxyanions transferred and reduced in theaquifer. The result of experiment #5, more importantly, showsthat the incubation of As(III) with MSS in the presence of0.04 M NaHCO3 converted most of As(III) to a charged speciesbelieved to be As(CO3)2

-, As(CO3)(OH)2-, or As(CO3)+.

We are not aware of any published stability constants forarseno-carbonate complexes. From extrapolation of thelinear line obtained by plotting the ionic radii against thepublished constants for trivalent element carbonate com-plexes (lanthanide and actinide series) (Figure 7), the stabilityconstants for arsenic complexes are estimated to be about107 for â1 and 1012.5 for â2 (34, 35). According to these estimates,which need to be confirmed experimentally, the carbonatecomplexes may be the most stable inorganic arsenic speciesin the aquatic environment. Subsequent conversion of thecarbonate ion pair to the common oxyanions of arsenic mostlikely occurs in groundwater:

Based on experimental results and above reactions, it isexpected that once formed As(III) carbonate complexes canpersist under anaerobic groundwater conditions at acidic to

neutral pH. The persistence of the carbonate complexes ingroundwater is of interest in terms of the toxicity of the arsenicand the removal of arsenic from drinking water.

The presence of arseno-carbonate complexes may ex-plain some of the reported inconsistencies in the geochem-istry of arsenic in groundwater. For instance, various studieshave reported highly variable concentrations of inorganicarsenic species in groundwater (2, 3, 6, 10, 36). Some of theso-called As(V) may, in fact, be As(III) carbonate complexessince the resin commonly used to separate As(III) from As(V)cannot readily distinguish between HAsO4

2-, H2AsO4-,

As(CO3)2-, and As(CO3)(OH)2

-. On the other hand, some ofthe so-called As(III) may be AsCO3

+ which would not beretained by the anion-capturing resin.

When orpiment is treated with HCO3-, both sulfide and

thioarsenite are also formed (eqs 4-7). Sulfide ion is acommon constituent in groundwater from MSS (16) andpresumably is derived from carbonation of arsenic sulfidesand dissimilatory sulfate reduction is possible. The fate ofthioarsenite in anaerobic groundwater is unclear. In morealkaline environments, reaction with OH- to form HAsS2(OH)-

may be possible (37). Under circumneutral to acidic condi-tions, the reaction of the sulfide and thioarsenite withabundant ferrous ion in groundwater can result in theformation of a number of sulfosalts:

Pyrite and arsenopyrite are common diagenetic minerals inMSS (22). It would thus appear that arsenopyrite is a productrather than the cause of arsenic accumulation in groundwater.

AcknowledgmentsThe authors gratefully acknowledge the U.S. GeologicalSurvey in Lansing, MI for providing core samples of MarshallSandstone and geological information. We thank the Rack-ham Graduate School, University of Michigan for awardingPredoctoral Fellowship to the first author.

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FIGURE 7. Predicted stability constants for arseno-carbonatecomplexes from constants for other metal complexes (La, Ce, Pr,Nd, Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, Lu, Y).

As(CO3)2- + 3H2O f H3AsO3 + 2HCO3

- + H+ (8)

As(CO3)+ + 3H2O f H3AsO3 + HCO3- + 2H+ (9)

As(CO3)2- + 4H2O f HAsO4

2- + 2HCO3- + 5H+ + 2e-

(10)

As(CO3)(OH)2- + H2O f H3AsO3 + HCO3

- (11)

Fe2+ + 2AsS2- f Fe(AsS2)2 (12)

4FeS + Fe(AsS2)2 f

2FeAsS (arsenopyrite) + 3FeS2 (pyrite) (13)

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Received for review August 12, 1999. Revised manuscriptreceived May 2, 2000. Accepted May 4, 2000.

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