Applied Geochemistry 26 (2011) 2247–2259

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Applied Geochemistry

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Sphalerite oxidation pathways detected by oxygen and sulfur isotope studies

Claudia Heidel a,⇑, Marion Tichomirowa a, Cornelia Breitkopf b

a Institute of Mineralogy, TU Bergakademie Freiberg, Brennhausgasse 14, 09599 Freiberg, Germanyb Institute of Power Engineering, TU Dresden, Helmholtzstraße 14, 01069 Dresden, Germany

a r t i c l e i n f o

Article history:Received 28 September 2010Accepted 1 August 2011Available online 26 August 2011Editorial handling by R.R. Seal

0883-2927/$ - see front matter � 2011 Elsevier Ltd. Adoi:10.1016/j.apgeochem.2011.08.007

⇑ Corresponding author. Tel.: +49 3731 392656; faxE-mail address: claudia.heidel@gmx.net (C. Heidel

a b s t r a c t

Sphalerite oxidation is a common process under acid-mine drainage (AMD) conditions and results in therelease of SO2�

4 , Zn and potentially toxic trace metals, which can pollute rivers and oceans. However, thereare only a few studies on the mechanisms of aerobic sphalerite oxidation. Oxygen and S isotope investi-gations of the produced SO2�

4 may contribute to the understanding of sphalerite oxidation mechanisms sohelping to interpret field data from AMD sites. Therefore, batch oxidation experiments with an Fe-richsphalerite were performed under aerobic abiotic conditions at different initial pH values (2 and 6) for dif-ferent lengths of time (2–100 days). The O and S isotope composition of the produced SO2�

4 indicatedchanging oxidation pathways during the experiments. During the first 20 days of the experiments at bothinitial pH values, molecular O2 was the exclusive O source of SO2�

4 . Furthermore, the lack of S isotopeenrichment processes between SO2�

4 and sphalerite indicated that O2 was the electron acceptor fromsphalerite S. As the oxidation proceeded, a sufficient amount of released Fe(II) was oxidized to Fe(III)by O2. Therefore, electrons could be transferred from sphalerite S sites to adsorbed hydrous Fe(III) andO from the hydration sphere of Fe was incorporated into the produced SO2�

4 as indicated by decreasingd18OSO4 values which became more similar to the d18OH2O values. The enrichment of 32S in SO2�

4 relativeto the sphalerite may also result from sphalerite oxidation by Fe(III).

The incorporation of O2 into SO2�4 during the oxidation of sphalerite was associated with an O isotope

enrichment factor eSO4–O2 of ca. �22‰. The O isotope enrichment factor eSO4–H2O was determined to be64.1‰. A comparison with O and S studies of other sulfides suggests that there is no general oxidationmechanism for acid-soluble sulfides.

� 2011 Elsevier Ltd. All rights reserved.

1. Introduction

Sulfide oxidation investigations commonly focus on pyrite dueto its ubiquitous occurrence and its tendency to produce acidityand high Fe and SO2�

4 concentrations (known as acid mine drain-age – AMD). Likewise, sphalerite (ZnS) oxidation results in the re-lease of Zn and SO2�

4 , which can pollute rivers and oceans. Forexample, oxidation of the remaining sphalerite in the abandonedmining district of Freiberg (Germany) contributes 37% of the Znpollution in the River Elbe, which drains into the North Sea (Martinet al., 1994). Furthermore, Fe and other potentially toxic trace met-als (e.g., Ag, Cd and Pb; De Giudici et al., 2002) are released fromsphalerite during its oxidation. Schippers and Sand (1999) pro-posed different oxidation mechanisms for acid-insoluble (e.g., pyr-ite) and acid-soluble sulfides (e.g., sphalerite). However, there areonly a few studies on the mechanisms of aerobic sphalerite oxida-tion that can aid interpretation of geochemical field data fromAMD sites.

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: +49 3731 394060.).

Sphalerite can be oxidized either by dissolved molecular O2 orFe(III) (e.g., Seal and Hammarstrom, 2003; Malmström and Collin,2004):

ZnSþ 2O2 ! Zn2þ þ SO2�4 ; ð1Þ

ZnSþ 8Fe3þ þ 4H2O! Zn2þ þ 8Fe2þ þ SO2�4 þ 8Hþ: ð2Þ

Reaction (1) implies that the resulting SO2�4 contains only

atmospheric-derived molecular oxygen, whereas all O in SO2�4 pro-

duced from reaction (2) should originate from water molecules.In addition, sphalerite can be dissolved non-oxidatively under

acid pH conditions (e.g., Seal and Hammarstrom, 2003; Malmströmand Collin, 2004):

ZnSþ 2Hþ ! Zn2þ þH2SðaqÞ: ð3Þ

Dissolved H2S may escape as gaseous H2S or may be oxidized toSO2�

4 by O2 (e.g., De Giudici et al., 2002; Seal and Hammarstrom,2003):

H2SðaqÞ ! H2SðgÞ; ð4Þ

H2SðaqÞ þ 2O2 ! 2Hþ þ SO2�4 : ð5Þ

Table 1Chemical composition of sphalerite determined by IR (S), ICP-OES (Zn), and ICP-MS(Al, Ca, Cd, Cu, Fe, Mg, Mn).

Element Zn S Fe Cu Cd Ca Mg Al Mn

Content (wt.%) 53.5 36.0 7.5 0.1 0.1 0.1 0.1 0.1 0.1Content (at.%) 39.1 53.5 6.4 0.1 0.4 0.1 0.1 0.1 0.0

2248 C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259

Schippers and Sand (1999) suggested that Fe(III) rather than O2

acts as sulfide oxidant. However, Weisener et al. (2004) proposedthat the Fe-rich sphalerite used in their study was oxidized byO2. They did not consider the potential role of Fe(III) as an oxidant.Furthermore, the formation of intermediate S species and the loca-tion of these reactions are still under discussion.

Sulfur isotopes of H2S produced by sphalerite oxidation shouldreflect the S source (sphalerite) and may reveal S isotope fraction-ation processes associated with the formation of intermediate Sspecies. The O isotope composition of SO2�

4 results from the relativecontribution of both potential O sources (O2 and water accordingto reactions (1) and (2)) and can be described as follows (Lloyd,1967):

d18OSO4 ¼ Xðd18OH2O þ eSO4�H2OÞ þ ð1� XÞðd18OO2 þ eSO4�O2Þ; ð6Þ

where d18OSO4, d18OH2O, and d18OO2 are the O isotope compositionsof SO2�

4 , water (d18OH2O < 0‰ for meteoric waters), and O2

(d18OO2 = 23.5‰, Kroopnick and Craig, 1972), respectively. X is theproportion of O in SO2�

4 which derives from water, (1 � X) is theresulting proportion of O2 in SO2�

4 . The eSO4–H2O value is the O iso-tope enrichment factor for the incorporation of O from water intoSO2�

4 , and eSO4–O2 is the O isotope enrichment factor for the incorpo-ration of O2 into SO2�

4 . The value of eSO4–H2O = 7.9‰ was determinedfor abiotic sphalerite oxidation (Balci et al., 2003), which differs sig-nificantly from eSO4–H2O = 0.0–4.0‰ for pyrite oxidation (Tayloret al., 1984; Balci et al., 2007; Mazumdar et al., 2008; Heidel andTichomirowa, 2011b). A value for eSO4–O2 has not been publishedfor sphalerite oxidation yet, whereas eSO4–O2 = �4.3‰ to �9.8‰

has been determined for abiotic pyrite oxidation (Taylor et al.,1984; Balci et al., 2007; Heidel and Tichomirowa, 2010).

Oxygen and S isotopes of SO2�4 have rarely been used for sphal-

erite oxidation investigations. Toran (1986) performed sphaleriteoxidation experiments under aerobic abiotic conditions, but couldnot determine O isotope compositions for SO2�

4 from a submersedexperiment. Gould et al. (1989) observed that SO2�

4 from bioticsphalerite oxidation experiments contained 57% water-derived Oand 43% O2 which was attributed to microbial activity and isotopeexchange processes. Balci et al. (2003) performed anaerobic sphal-erite oxidation studies. They observed that the O isotope signatureof SO2�

4 and water differed by 7.9‰ and 8.5‰ for abiotic and bioticconditions, respectively. Balci et al. (2003) concluded that atmo-spheric-derived O2 was incorporated into SO2�

4 . However, there isstill a lack of comprehensive O and S isotope data and an eSO4–O2

value for abiotic aqueous sphalerite oxidation.This study conducted batch oxidation experiments with an Fe-

rich sphalerite under aerobic abiotic conditions at different initialpH values (2 and 6) for different lengths of time (2–100 days).The O and S isotope composition of the produced SO2�

4 was inves-tigated in order to determine O and S sources and isotope enrich-ment factors. This information should give more detailed insightsinto sphalerite oxidation pathways. A comparison with isotopestudies of further sulfide minerals may reveal the presence (or ab-sence) of a general oxidation mechanism for acid-soluble sulfides.In addition, results may help to interpret hydrochemical and iso-tope field data from AMD.

2. Material and methods

2.1. Material

Sphalerite pieces was obtained from the Mineralogical Collec-tion, TU Bergakademie Freiberg. The sulfide material originatedfrom Lengefeld, Ore Mountains, Germany. Sphalerite was groundand sieved to obtain a grain size 63–100 lm. Visible intergrowthswith quartz were separated under the binocular microscope.

Before starting the experiments, ultrafine particles were removedby rinsing with acetone. Subsequently, the grains were dried in abox under a N2 atmosphere and used for the experiments. Moseset al. (1987) observed that a pre-treatment procedure with boilingof sphalerite in HCl resulted in the loss of S by degassing as H2S(according to reactions (3) and (4)). Therefore, sphalerite is com-monly pre-treated without washing in acid (e.g., Weisener et al.,2004; Abraitis et al., 2004; Malmström and Collin, 2004).

A specific surface area of 0.289 m2 g�1 was determined by BET(Brunauer–Emmett–Teller) measurements of the pre-treatedsphalerite at the Granulometric Laboratory of the Institute of Elec-tronic and Sensor Materials, TU Bergakademie Freiberg. The sphal-erite composition was determined by combustion with infrareddetection (IR), inductively coupled plasma-optical emission spec-trometry (ICP-OES), and inductively coupled plasma-mass spec-trometry (ICP-MS) at the Activation Laboratories Ltd., Ancaster,Ontario. Analyses showed that the sphalerite was relatively Fe-rich(Table 1: 7.5 wt.% Fe).

The bulk sphalerite composition of the pre-treated (after wash-ing in acetone) sphalerite was investigated by X-ray diffractometry(XRD) (URD 6 from Seifert-FPM) at the Mineralogical Laboratory ofthe Institute of Mineralogy, TU Bergakademie Freiberg. XRD inves-tigations showed that traces of quartz were still present, i.e., quartzwas not completely excluded. Peaks of in the XRD spectra were as-signed to only one other mineral: sphalerite. The spectra did notreveal any peaks associated with pyrite or other sulfide minerals.XRD spectra of Fe-poor and Fe-rich sphalerite are similar. Boththese observations implied that Fe occurred as a transition metalimpurity in sphalerite and not in a separate phase.

2.2. Experiments

Aerobic abiotic sphalerite batch oxidation experiments werecarried out in duplicate with two different waters: arctic rainfall(d18OH2O = �17.7‰) and ultrapure (18.2 MX cm) deionized water(d18OH2O = �8.7‰). Initial pH values of 2 and 6 were adjusted byadding 1 M HCl and 1 M NaOH, respectively. No pH buffer wasadded. Waters were filtered through 0.2 lm cellulose nitrate filters(Macherey–Nagel) prior to the experiments.

Water (250 mL) and 50 g sphalerite were loaded in an acid-washed 530 mL flask. The relatively high mass of 50 g sphaleritewas chosen to obtain a sufficient amount of SO2�

4 to allow repeatedO and S isotope measurements. The headspace of the flask (ca.280 mL) was initially filled with ambient air; but the experimentswere performed in a closed system (i.e., flasks were sealed) toavoid evaporation effects. Sphalerite was allowed to oxidize in adark air-conditioned room at 21 �C for various lengths of time(2–100 days).

2.3. Morphological and mineralogical investigations

After finishing the experiments, sphalerite was dried in a desic-cator. The grain morphology was observed by scanning electronmicroscopy (SEM) (JSM-6400 from JEOL) and the chemical compo-sition was determined by energy-dispersive X-ray analysis (EDX)(Noran Vantage from Noran Instruments) at the SEM laboratoryof the Institute of Geology, TU Bergakademie Freiberg. Selectedsamples were investigated by XRD.

C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259 2249

Investigations of the sphalerite surface composition of the pre-treated sphalerite and samples from experiments finished after 2and 20 days were performed by X-ray photoelectron spectroscopy(XPS) at the Institute of Technical Chemistry, University of Leipzig.XPS analysis was carried out using a hemispherical energy analyzerPhoibos 150 (Specs GmbH) equipped with an X-ray source XR 50(Specs GmbH). All spectra were analyzed using the software pack-age CasaXPS Version 2.2.24.

2.4. Hydrochemical analyses

Before finishing the experiments (i.e., opening the flasks), dis-solved O2 (DO) concentrations were measured non-invasively fromoutside through the wall of the flasks by an oxygen meter con-nected to fiber-optic oxygen minisensors which were glued insidethe flasks (Fibox 3 with planar oxygen-sensitive spots from Pre-Sens). An error of better than 1% is given by the instrumentmanufacturer.

After opening the flasks, pH values were measured immedi-ately by means of a pH electrode (pH 340 with a SenTix 41 sen-sor from WTW). Simultaneously, dissolved Fe2+ and total Feconcentrations were determined spectrophotometrically usingthe 1, 10 phenanthroline method (photoLab S12 from WTW).Sulfate, SO2�

3 , and S2O2�3 concentrations were determined by

ion chromatography (IC) (DX-120 from Dionex) at the Geochem-ical-Analytical Laboratory of the Institute of Mineralogy, TUBergakademie Freiberg. Zinc concentrations were measured byinductively coupled plasma-optical emission spectrometry (ICP-OES) (Plasma 2000 from Perkin Elmer) at the ICP laboratory ofthe Institute of Analytical Chemistry, TU Bergakademie Freiberg.Additionally, control measurements of total Fe concentrationswere performed by ICP-OES for the experiments that lasted100 days. Errors were better than 7% for IC, 5% for ICP-OES,and 10% for spectrophotometry based on repeated measure-ments of standards.

2.5. Isotope analyses

Oxygen isotopes of SO2�4 , O2, and water and S isotopes of SO2�

4

and sphalerite were determined with a mass spectrometer, Deltaplus (Finnigan MAT), at the Geochemical Isotope Laboratory ofthe Institute of Mineralogy, Freiberg. The common d-notation isused for reporting isotope ratios:

d ¼ Rsample

Rstandard

� �� 1

� �� 1000ð‰Þ; ð7Þ

Fig. 1. SEM images of sphalerite which was oxidized for 1

where R = 18O/16O and R = 34S/32S, relative to the standard Vienna-Standard Mean Ocean Water (V-SMOW) for d18O und Vienna-Can-yon Diablo Troilite (V-CDT) for d34S.

The d18O values of dissolved O2 were measured as d18OO2 valuesof the headspace gas (Oba and Poulson, 2009). The measuring tech-nique was modified after Wassenaar and Koehler (1999). An ali-quot of the headspace gas flowed into an evacuated capillary andthrough a sample loop into a He flow, where it was carried througha water and CO2 trap and a gas chromatography column into themass spectrometer. Isotope analyses were compared against air(d18OO2 = 23.5‰, Kroopnick and Craig, 1972). The d18OH2O valueswere measured by isotope ratio mass spectrometry with a dual in-let system after the solution was equilibrated with CO2 (Epsteinand Mayeda, 1953).

Sulfate for O and S isotope measurements was prepared as fol-lows: If necessary, Fe was precipitated by adjusting the pH value ofthe aqueous solution to greater than 7 with 1 M NaOH. After filtra-tion of the aqueous solution, the pH value was adjusted to about3.5 with 1 M HCl. Barium sulfate was precipitated by adding BaCl2

into the heated solution. Afterwards, precipitated BaSO4 was col-lected on a 0.45 lm cellulose nitrate filter (Sartorius) and rinsedwith deionized water to remove Cl�. The d18OSO4 values andd34SSO4 values were measured by continuous flow isotope ratiomass spectrometry (CF-IRMS) using pyrolysis (e.g., Kornexl et al.,1999) and an elemental analyzer EA 1110 from Carlo Erba (e.g.,Giesemann et al., 1994), respectively. Oxygen isotope ratios werenormalized to NBS 127, IAEA-SO-5, and IAEA-SO-6 (d18O = 8.7‰,d18O = 12.0‰, and d18O = �11.0‰, Kornexl et al., 1999) and S iso-tope ratios were normalized to an internal Ag2S standard andIAEA-SO-2 and IAEA-SO-3 (d34S = 22.7‰ and d34S = �32.3‰, Dinget al., 2001). All samples were measured at least in triplicate. Thereproducibility was better than 0.2‰ for d18OH2O values, 0.5‰ ford18OO2 and d18OSO4 values, and 0.3‰ for d34SSO4 and d34SZnS values.

3. Results

3.1. Morphology and mineralogy

SEM images (Fig. 1) showed that very fine material was stillpresent at the end of the longest experiments (100 days). Some pitsand etch features were observed on the sphalerite surfaces (Fig. 1).SEM/EDX analyses identified an Fe-rich sphalerite with traces ofquartz, chlorite, and mica. Oxygen- and S-bearing precipitatescould not be detected by SEM/EDX investigations.

XRD measurements did not show differences between the start-ing material and selected sphalerite samples. As mentioned in

00 days at initial pH 2 (left) and initial pH 6 (right).

Table 2Total concentrations of S, Zn, Fe and S species on sphalerite surfaces determined by XPS measurements. Carbon and O have been normalized out.

Experiment duration(days)

S(2p)(at.%)

Zn(2p)(at.%)

Fe(2p)(at.%)

MolarratioS/(Zn + Fe)

S2� (at.%)161.4–161.8 eV

S2�2 (at.%)

162.5–162.7 eV

S2�n>2 (at.%) 163.2 eV SO2�

4 (at.%) 169.1 eV

0 (pre-treated sphalerite) 52.7 41.0 6.3 1.1 31.1 17.4 4.3

Initial pH 22 55.3 38.7 6.0 1.2 30.9 24.520 47.6 37.7 14.7 0.9 34.3 13.3

Initial pH 62 52.2 41.5 6.3 1.1 22.9 29.320 50.2 43.0 6.8 1.0 30.0 20.2

2250 C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259

Section 2.1, sphalerite and quartz could be identified by XRD spec-tra, whereas XRD spectra of the investigated material did not re-veal any peaks associated with pyrite or other sulfide minerals.XPS measurements indicated that minor amounts of SO4 appearedon the surface of the pre-treated sphalerite. However, SO4 was notdetected on the surface of sphalerite which was oxidized for 2 and20 days in experiments at both initial pH values (Table 2). Theobserved binding energies from the S2p spectra suggested thatmonosulfide (at 161.1–161.8 eV) and disulfide species (at162.5–162.7 eV) occurred on sphalerite surfaces from theseexperiments. A slightly higher binding energy (163.2 eV) fromthe experiment at initial pH 2 which was finished after 20 daysmay indicate that polysulfide species appeared instead of disulfide

Table 3Chemical composition of the solutions. Total Fe concentrations after 100 days were measure(SO2�

4 , SO2�3 , and S2O2�

3 ).

Experiment duration (days) pH Concentrations (mmol L�1)

DO Zn Fe(II)

Initial pH 2, initial d18OH2O = �17.7‰

2 3.1 0.21 0.54 0.575 5.3 0.18 0.52 0.43

10 6.1 0.18 0.59 0.1620 6.8 0.19 0.61 nd57 5.9 0.10 1.09 nd76 6.6 0.14 0.97 nd

100 7.1 0.12 1.03 nd

Initial pH 2, initial d18OH2O = �8.7‰

2 3.2 0.19 0.40 0.325 5.6 0.20 0.47 0.33

10 6.2 0.18 0.47 0.1520 6.9 0.21 0.54 nd57 6.2 0.11 1.09 nd76 6.3 0.11 1.16 nd

100 6.5 0.08 1.33 nd

Initial pH 6, initial d18OH2O = �17.7‰

2 7.3 0.20 0.07 nd5 7.5 0.17 0.07 nd

10 7.7 0.20 0.03 nd20 7.9 0.22 0.02 nd58 7.4 0.16 0.08 nd76 7.4 0.16 0.08 nd

100 7.7 0.14 0.11 nd

Initial pH 6, initial d18OH2O = �8.7‰

2 7.3 0.21 0.09 nd5 7.6 0.16 0.11 nd

10 7.8 0.24 0.03 nd20 7.9 0.23 0.03 nd58 7.3 0.17 0.07 nd76 7.3 0.16 0.06 nd

100 7.6 0.15 0.08 nd

nd = Not detectable.

(Table 2). The surface composition of sphalerite from experimentsat initial pH 6 did not show obvious changes from 0 to 20 days (Ta-ble 2). Considering that the binding energy of elemental S is164.1 eV (Weisener et al., 2004), XPS data from the experimentsdid not indicate the formation of elemental S during the first20 days of oxidation. Iron2p spectra indicated that Fe occurred asFe(III) (binding energy of 710.8–711.3 eV) not only after 2 and20 days of oxidation but also before starting the experiments.The observed binding energy is assigned to Fe oxide or Fe oxyhy-droxide (Cai et al., 2009). The atomic weight of Fe on the sphaleritesurface (Table 2: 6–7 at.%) agreed with those from the bulk sphal-erite (Table 1: 6.4 at.%) except of the sample which was oxidized atan initial pH 2 for 20 days (Table 2: 14.7 at.%).

d by ICP-OES. The molar ratio S/Zn was calculated with cumulative dissolved S species

Molar ratio S/Zn

Fe (total) SO2�4 SO2�

3 S2O2�3

0.54 0.37 nd 0.002 0.70.42 0.56 nd 0.003 1.10.16 0.53 0.0004 nd 0.9nd 0.52 nd 0.002 0.9nd 1.27 nd nd 1.2nd 0.67 0.0013 0.038 0.70.003 0.67 nd 0.052 0.7

0.32 0.33 nd 0.001 0.80.33 0.52 nd 0.003 1.10.15 0.55 0.0015 0.002 1.2nd 0.72 nd 0.001 1.3nd 1.46 nd 0.009 1.4nd 1.52 nd nd 1.30.005 1.79 nd nd 1.3

nd 0.41 nd 0.008 6nd 0.47 nd 0.017 7nd 0.51 0.0006 0.053 18nd 0.61 0.0010 0.097 33nd 0.73 0.0025 0.096 10nd 0.71 0.0034 0.151 100.003 0.82 0.0013 0.162 9

nd 0.58 nd 0.008 7nd 0.43 nd 0.017 4nd 0.51 0.0003 0.060 22nd 0.74 0.0023 0.086 33nd 0.72 0.0028 0.123 12nd 0.67 0.0044 0.193 150.002 0.68 0.0014 0.171 11

C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259 2251

3.2. Hydrochemistry

The results of the hydrochemical measurements (Table 3, andFig. 2) indicated that an initial fast oxidation was followed by aslower oxidation. Experiments at both initial pH values showed anincipiently sharp pH increase which was followed by slight pH vari-ations (Fig. 2). After 100 days, a pH value of about 6–7 was achievedin experiments at initial pH 2 and a pH of about 7–8 was measured inexperiments at initial pH 6. Dissolved O2 concentrations showed acontinuous decrease in experiments at both initial pH values (Table3). However, the decrease in DO concentrations and the pH increasewere stronger in experiments at initial pH 2 compared with initialpH 6.

Zinc concentrations increased sharply at the beginning of theexperiments at initial pH 2. Afterwards, the Zn release decreased(Fig. 2). At initial pH 6, dissolved Zn concentrations remained lowthroughout the experiments. Dissolved Fe (occurring as Fe(II)) wasdetected only at the beginning of the experiments at initial pH 2. La-ter on (and throughout the experiments at initial pH 6), dissolved Feconcentrations were below the detection limit of 0.005 mmol L�1

(spectrophotometric measurements). Control measurements per-formed by ICP-OES after 100 days showed that small amounts of dis-solved Fe were present (Table 3: 0.002–0.005 mmol L�1).

Sulfate concentrations indicated that experiments with thesame pH and d18OH2O value but different experiment durationproceeded non-uniformly, e.g., experiments at initial pH 2 andd18OH2O = �17.7‰ showed the highest SO2�

4 concentration after57 days. Likewise, SO2�

4 concentrations in experiments at initialpH 6 and d18OH2O = �8.7‰ did not increase continuously with

Fig. 2. Hydrochemical measurements from experiments at initial pH 2 (left) and initial pline indicates the change from the initial fast reaction to the slower reaction.

increasing experiment duration but the highest SO2�4 concentration

was obtained after 20 days. In contrast, SO2�4 concentrations in

experiments at initial pH 2 and d18OH2O = �8.7‰ and experimentsat initial pH 6 and d18OH2O = �17.7‰ increased with progressiveexperiment duration (Fig. 2).

The increase in SO2�4 concentrations was strongest within the

first 2 and 5 days for experiments at initial pH 6 and initial pH 2,respectively. Afterwards, the SO2�

4 release into solution sloweddown especially for experiments at initial pH 6 (Fig. 2). Concentra-tions of the intermediate S species thiosulfate were low in experi-ments at initial pH 2, but increased continuously in experiments atinitial pH 6 (Table 3). Sulfite was detected only in a few experi-ments at initial pH 2. Experiments at initial pH 6 – where sphaler-ite was allowed to oxidize for more than 10 days – showed low butdetectable sulfite concentrations (Table 3). Dissolved H2S concen-trations were not measured, but a slight H2S odor was observedduring the opening of the flasks from experiments at initial pH 2lasting 2–100 days. The molar ratio between dissolved S species(sulfate, sulfite, thiosulfate) and Zn ranged from 0.7 to 1.4 in exper-iments at initial pH 2 (Table 3). Thus, considering that the atomicratio S/Zn was 1.4 in sphalerite (from Table 1), aqueous solutionsshowed a slight Zn excess or S deficit. In contrast, the molar ratioS/Zn ranged from 4 to 33 in experiments at initial pH 6 indicatinga large Zn deficit or S excess in solution.

Oxidation rates were estimated based on the produced amountof dissolved S species:

RZnS ¼a

ABET �mðmol m�2 s�1Þ; ð8Þ

H 6 (right) with d18OH2O = �17.7‰ (top) and d18OH2O = �8.7‰ (bottom). The dotted

2252 C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259

where RZnS is the sphalerite oxidation rate, a is the slope (mol s�1)from the linear regression of dissolved S concentration (sulfate, sul-fite, thiosulfate) vs. time, ABET is the sphalerite specific surface area(0.289 m2 g�1), and m is the oxidized sphalerite mass (50 g). Sphal-erite oxidation rates were 8.2 ± 0.3 � 10�11 mol m�2 s�1 (1r, n = 2)for the first 5 days of the experiments at initial pH 2 and2.0 ± 0.5 � 10�10 mol m�2 s�1 (1r, n = 2) for the first 2 days of theexperiments at initial pH 6. Afterwards, oxidation rates decreasedto 7.0 ± 6.2 � 10�12 mol m�2 s�1 and 3.4 ± 0.9 � 10�12 mol m�2 s�1

in experiments at initial pH 2 and initial pH 6, respectively. Presum-ably, oxidation rates were somewhat higher due to the formation offurther intermediate S species (e.g., polysulfides) which were notconsidered in the calculations. In addition, the dissolved S deficitin experiments at initial pH 2 and the dissolved S excess in experi-ments at initial pH 6 suggested that the released species underwentphysico-chemical reactions (e.g., degassing and re-dissolution ofH2S, oxidation of gaseous H2S) which may modify the calculationof the oxidation rates. Nonetheless, a strong decrease of the oxida-tion rates was obvious for all experiments.

3.3. Oxygen and sulfur isotope chemistry

The d18OH2O values did not change during the experiments (i.e.,mostly agreed within ±0.2‰) whereas d18OO2 values increasedfrom 23.5‰ (Kroopnick and Craig, 1972) to up to 31.6‰ (Table 4,Fig. 3).

The d18OSO4 values remained more or less constant (rangedfrom 0.4‰ to 4.0‰) at the beginning of the experiments (days 2–20) at both initial pH values and were similar regardless of the ini-tial d18OH2O value. Afterwards (days 57–100), the d18OSO4 valuesdecreased by up to 14.5‰ and 8.5‰ for experiments with

Table 4Oxygen and S isotope measurements. The initial d18OO2 value of ambient air was 23.5‰ (

Experiment duration (d) d18OH2O (‰)

Initial pH 2, initial d18OH2O = �17.7‰

2 �17.65 �17.6

10 �17.520 �17.757 �17.676 �18.0

100 �18.0

Initial pH 2, initial d18OH2O = �8.7‰

2 �8.85 �8.8

10 �8.820 �8.957 �9.276 �9.1

100 �9.2

Initial pH 6, initial d18OH2O = �17.7‰

2 �17.65 �17.6

10 �17.220 �17.858 �17.976 �17.7

100 �18.1

Initial pH 6, initial d18OH2O = �8.7‰

2 �8.85 �8.9

10 �8.820 �9.058 �9.176 �9.0

100 �9.0

na = Not analyzed.

d18OH2O = �17.7‰ and d18OH2O = �8.7‰, respectively, and becamemore similar to the d18OH2O values (Fig. 3). Thus, d18OSO4 valueschanged significantly after 20 days, whereas hydrochemicalparameters changed markedly after only 2 and 5 days in experi-ments at initial pH 6 and 2, respectively (see Section 3.2).

The relative proportion of water-derived O in SO2�4 (and thus,

the relative proportion of O2 in SO2�4 ) can be estimated indepen-

dent of eSO4–H2O and eSO4–O2 by rearrangement of Eq. (6):

d18OSO4 ¼ Xðd18OH2OÞ þ ½ð1� XÞðd18OO2 þ eSO4�O2Þþ XðeSO4�H2OÞ�: ð9Þ

Accordingly, the slope of a regression line of d18OSO4 values vs.d18OH2O values (multiplied by 100) equals the approximatewater-derived proportion of O in SO2�

4 (e.g., Gould et al., 1989; Balciet al., 2007; Nordstrom et al., 2007). Because only two d18OH2O val-ues were used in this study, the regression line was formed by onlytwo data points. Thus, results should be regarded as rough estima-tions. During the first 20 days, SO2�

4 from experiments at both ini-tial pH values contained about 0% water-derived O (Fig. 4). Fromdays 57 to 100, water-derived O proportions of SO2�

4 differed forexperiments at initial pH 2 and 6. Experiments at initial pH 6showed an increasing proportion of water-derived O in SO2�

4 (fromca. 34% to 76%) with increasing experiment duration (Fig. 4). Incontrast, water-derived O proportions of SO2�

4 in experiments atinitial pH 2 showed no systematic trend with increasing experi-ment duration; but implausible negative values were calculatedfor the experiments lasting 76 and 100 days.

The d34SSO4 values remained constant and agreed within error(±0.3‰) with d34SZnS = 0.4‰ during the first 20 days (Table 4,Fig. 5). Subsequently (days 57–100), the d34SSO4 values decreased

Kroopnick and Craig, 1972). The d34S value of the oxidized sphalerite was 0.4‰.

d18OO2 (‰) d18OSO4 (‰) d34SSO4 (‰)

na 2.0 0.1na 2.2 0.7na 1.6 0.0na 2.1 0.2na �8.7 �0.827.1 �1.2 �1.229.6 �2.5 �2.0

na 1.1 0.0na 3.4 0.3na 1.1 0.1na 2.0 �0.327.9 �2.9 �0.729.2 �2.8 �0.7na �5.1 �1.1

na 4.0 0.2na 1.2 0.5na 0.4 0.8na 1.5 �0.126.6 �7.9 �0.627.9 �8.5 �1.131.6 �10.5 �2.1

na 2.3 0.4na 2.0 0.3na 1.4 0.3na 1.4 0.027.2 �4.9 �1.028.3 �3.7 �1.730.7 �3.6 �2.3

Fig. 3. Oxygen isotope composition of SO2�4 , O2, and water in experiments with d18OH2O = �17.7‰ (left) and d18OH2O = �8.7‰ (right).

Fig. 4. Determination of the relative proportion of water-derived O in SO2�4 from the slope of the regression line of d18OSO4 values vs. d18OH2O values (multiplied by 100)

according to Eq. (9) from experiments at initial pH 2 (left) and initial pH 6 (right). Error bars represent the the analytical uncertainty of the d18OSO4 values (0.5‰).

Fig. 5. Sulfur isotope composition of SO2�4 in experiments with d18OH2O = �17.7‰ (left) and d18OH2O = �8.7‰ (right).

C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259 2253

2254 C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259

to �2.3‰ which reflected an enrichment of 32S in SO2�4 relative to

sphalerite.

4. Discussion

4.1. Changing reaction rates during sphalerite oxidation experiments

The obtained reaction rates were in the range of previous sphal-erite dissolution studies (e.g., Acero et al., 2007). Furthermore,sphalerite oxidation rates decreased during the experiments whichhas already been observed during several sphalerite dissolutionstudies at different pH conditions, temperatures, and Fe contentof sphalerite (e.g., Weisener et al., 2003; Abraitis et al., 2004;Malmström and Collin, 2004; Cama and Acero, 2005; Acero et al.,2007). Acero et al. (2007) suggested that altered layers wereformed on sphalerite surfaces during grinding. Therefore, initiallyhigh reaction rates were caused by the dissolution of these alteredlayers. Buckley et al. (1989) and Weisener et al. (2004) proposedthat the formation of a polysulfide layer during the initial sphaler-ite oxidation caused a passivation of the sphalerite surface result-ing in slower oxidation rates. Constant oxidation rates during thelater course of sphalerite oxidation experiments were explainedby a decreasing proportion of polysulfides and the developmentof a porous layer of elemental S on the sphalerite surface (Weiseneret al., 2004). Malmström and Collin (2004) and Acero et al. (2007)suggested that initially high oxidation rates were caused by thedissolution of microparticles. Possible effects of the release of ini-tial SO2�

4 , the formation of polysulfides and Fe(III) oxyhydroxides,and the oxidation of ultrafine sphalerite grains are discussed inthe following sections.

4.1.1. Release of initial sulfateXPS investigations showed that minor amounts of SO2�

4 werepresent on the surface of pre-treated sphalerite (Table 2) be-cause the pre-treatment procedure was not performed understrict anaerobic conditions (i.e., grinding and sieving occurredin the presence of air) and sphalerite surfaces could not be com-pletely cleaned from potential SO2�

4 coatings (see Section 2.1).The release of initial SO2�

4 formed from pre-experimental air oxi-dation may explain high SO2�

4 concentrations at the beginning ofthe experiments. If so, d18OSO4 values should be independent ofthe d18OH2O value during the first 2 and 5 days (in experimentsat initial pH 6 and initial pH 2, respectively). Afterwards,decreasing reaction rates implied that the produced SO2�

4 origi-nated only from aqueous sphalerite oxidation. Therefore,d18OSO4 values should have changed because aqueous sphaleriteoxidation became the dominant SO2�

4 ‘‘source’’. However,d18OSO4 values did not change during the first 20 days of theexperiments although significant amounts of SO2�

4 were newlyformed from sphalerite oxidation (Fig. 3). Thus, the release ofinitial SO2�

4 may contribute to changing reaction rates but isnot the dominating factor for them.

4.1.2. Surface passivation by polysulfidesXPS spectra indicated that polysulfide species were formed on

the surface of sphalerite which was oxidized at initial pH 2(Fig. 2). Weisener et al. (2004) proposed that polysulfides areformed via disulfide by the following sequence of reactions:

4S2�ðsÞ þ 4Hþ þ O2 ! 2S2�2 ðsÞ þ 2H2O; ð10Þ

3S2�2 ðsÞ þ 4Hþ þ O2 ! S2�

3 ðsÞ þ 2H2O; ð11Þ

8S2�2 ðsÞ þ 12Hþ þ 3O2 ! 2S2�

4 ðsÞ þ 6H2O: ð12Þ

Accordingly, the large proton consumption (i.e., pH increase)during the first 5 days of the experiments at initial pH 2 also cor-roborates the formation of polysulfides which may passivate thesphalerite surface. In contrast, polysulfides could not be found onsphalerite grains from experiments at initial pH 6 although reac-tion rates also decreased during these experiments. At least forthe experiments at initial pH 6, decreasing reaction rates cannotbe attributed only to the formation of polysulfides.

4.1.3. Surface passivation by ferric oxyhydroxidesFerrous iron from the Fe-rich sphalerite was released into solu-

tion especially at the beginning of the experiments at initial pH 2(Table 3). Due to fast increase of the pH value, the Fe(II) oxidationrate increased rapidly (Singer and Stumm, 1970). Thus, Fe(III) oxy-hydroxides were formed as indicated by decreasing concentrationsof dissolved Fe (Table 3):

Fe2þ þ 0:25O2 þ 2:5H2O! FeðOHÞ3 þ 2Hþ: ð13Þ

Ferric oxyhydroxides also may be formed in experiments at ini-tial pH 6 as indicated by the large Zn deficit in solution which mayresult from the adsorption of Zn ions onto Fe(III) oxyhydroxides(e.g., Johnson, 1986; Smith, 1999; Seal and Hammarstrom, 2003).Despite the presence of Fe(III) oxyhydroxides, sphalerite surfaceswere probably not substantially passivated. Otherwise, the sphal-erite oxidation should be stopped after passivation of the sulfidesurface.

4.1.4. Initial oxidation of ultrafine sphalerite grainsUltrafine sphalerite grains should not have been present be-

cause a grain size of 63–100 lm was used for the experiments.However, small grains adhering on larger grains may not have beenremoved by the pre-treatment procedure (see Section 2.1). A rela-tively high proportion of very fine material was still present after100 days of oxidation (Fig. 1). Due to the lack of SEM images frompre-treated sphalerite, it remains speculative if more ultrafine par-ticles occurred before starting the experiments. However, the dis-solution of ultrafine material is a pH-independent process, whichmay explain decreasing reaction rates in experiments at both ini-tial pH values. Thus, the initial oxidation of ultrafine sphalerite par-ticles may be the main reason for decreasing oxidation rates. Therelease of initial SO2�

4 and the formation of polysulfides and Fe(III)oxyhydroxides may additionally contribute to decreasing reactionrates but they are not the dominating factors.

4.2. Relative proportions of molecular oxygen and water-derivedoxygen in sulfate

During the first 20 days, d18OH2O-independent d18OSO4 valuesindicated that no water-derived O was incorporated into the SO2�

4

produced during all experiments (Fig. 4). Accordingly, O in SO2�4

originated completely from O2. Afterwards, decreasing d18OSO4 val-ues indicated changing contributions of O from both potentialsources which may result from changing oxidation pathways. Fromdays 58 to 100, an increasing proportion of water-derived O in SO2�

4

(from 34% to 76%) was obtained from experiments at initial pH 6(Fig. 4). Negative proportions of water-derived O were calculatedfor experiments at initial pH 2 from days 76 to 100. These implausi-ble results are attributed to different oxidation rates (i.e., amounts ofproduced SO2�

4 ) and are not considered further in the discussion. Dif-ferent oxidation rates in experiments with the same initial pH andexperiment duration but different d18OH2O values and the generalnon-uniform SO2�

4 increase with increasing experiment time(Fig. 2) are attributed to a different proportion of ultrafine grains.The difficulty of creating a homogenous grain size distribution andthe impact of different grain sizes on the oxidation rates have

Fig. 6. Determination of eSO4–O2 from days 57 to 100 from the slope of theregression line of ln [(d18OO2 + 1000)/(d18Oinitial O2 + 1000)] vs. ln F according to Eqs.(14) and (15) from experiments at initial pH 2 (left) and initial pH 6 (right) (see textfor calculation).

C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259 2255

already been discussed by Tichomirowa and Junghans (2009) forpyrite oxidation experiments.

Both SO2�4 and dissolved intermediate S species (e.g., sulfite, thio-

sulfate) may exchange O isotopes with water which would result in apseudo-increase of the water-derived O proportion in SO2�

4 . How-ever, the occurrence of O isotope exchange reactions during theexperiments can be excluded for several reasons. The O isotope ex-change rate between SO2�

4 and water increases with increasing tem-perature and decreasing pH value (Hoering and Kennedy, 1957;Chiba and Sakai, 1985; Tichomirowa and Junghans, 2009). The ex-change rate is extremely low at moderate temperatures and (nearly)neutral pH conditions, i.e., the half-life of exchange should be�1 ka(Tichomirowa and Junghans, 2009). Therefore, O isotope exchangereactions between SO2�

4 and water should not occur during theexperiments.

Although the half-life of the O isotope exchange between dis-solved SO2�

3 and water also increases with increasing pH (Bettsand Voss, 1970), the exchange rate is extremely high even underneutral pH conditions. At pH 7, SO2�

3 would exchange 50% of its O iso-topes with those of water within 18 ms (Hubbard et al., 2009). Brun-ner et al. (2006) observed that the isotope exchange is associatedwith an enrichment of 18O in SO2�

3 relative to water (expressed aseSO3–H2O). They determined a eSO3–H2O value of 11.5‰ at pH 7.2. IfSO2�

3 undergoes O isotope exchange reactions with water and is oxi-dized to SO2�

4 afterwards, an enrichment of 18O relative to water willalso occur in the produced SO2�

4 . However, d18OSO4 values in thepresent work became more similar to d18OH2O values after more than20 days of sphalerite oxidation (Fig. 3). Thus, the O isotope composi-tion of SO2�

4 has not been affected by O isotope exchange betweenSO2�

3 and water.Oxygen isotope exchange reactions between SO2�

4 and water orSO2�

3 and water cannot explain the observed d18OSO4 values, i.e., theincrease of the water-derived O proportion in sulfate. The O iso-tope exchange between other intermediate S species (e.g., thiosul-fate) and water molecules cannot be excluded but has not beenstudied yet due to extremely fast exchange rates.

4.3. Oxygen isotope enrichment factors eSO4–O2 and eSO4–H2O associatedwith sphalerite oxidation and e value associated with the reduction ofmolecular oxygen

From day 2 to 20, O2 was the exclusive O source of SO2�4 (Fig. 4).

Thus, the O isotope difference between SO2�4 and O2 (D18OSO4–O2 =

d18OSO4–d18OO2) resulted solely from O isotope enrichment pro-cesses that occurred during the incorporation of O2 into SO2�

4 .Therefore, the O isotope enrichment factor eSO4–O2 = �22‰ wasroughly estimated from the average D18OSO4–O2 value (calculatedwith a constant d18OO2 value of 23.5‰ and d18OSO4 values fromexperiments at both initial pH values from Table 4). Accordingly,an enrichment of 16O in SO2�

4 relative to O2 occurred during theoxidation of sphalerite, i.e., 16O was preferentially consumed dur-ing sphalerite oxidation.

Sulfate produced from day 57 to 100 contained water-derived O(Fig. 4). Therefore, eSO4–O2 could not be estimated from D18OSO4–O2

values. However, the e value associated with the reduction of O2

can be calculated from the observed increase in d18OO2 values fromday 57 to 100 (Table 4). This e value results not only from the pref-erential incorporation of 16O into SO2�

4 during the sphalerite oxida-tion but also from the preferential consumption of 16O during theoxidation of Fe(II) and intermediate S species (e.g., H2S). Thed18OO2 values and DO concentrations are expected to follow a Ray-leigh fractionation trend (Oba and Poulson, 2009):

lnd18OO2 þ 1000

d18OinitialO2 þ 1000

!¼ ða� 1ÞlnF; ð14Þ

where d18OO2 is the measured d18OO2 value, d18Oinitial O2 is thed18OO2 value at the beginning of the experiments, a is the O isotopefractionation factor, and F is the fraction of remaining DO. The slopeof the regression line in Fig. 6 equals (a � 1). The correspondingenrichment factor e was calculated as follows:

e ¼ 1000ða� 1Þ: ð15Þ

The resulting e value associated with the reduction of O2 was�26.9 ± 6.4‰ in the experiments at initial pH 6 from day 58 to100 (Fig. 6) and it agreed roughly with the eSO4–O2 value of ca.�22‰ estimated from D18OSO4–O2 values (for day 2–20). Thus,the dominant O-consuming processes (incorporation of O2 intoSO2�

4 during sphalerite oxidation, oxidation of Fe(II)) should havesimilar e values. Furthermore, the observed e value associated withthe reduction of O2 suggests that the enrichment of 16O in SO2�

4 rel-ative to O2 is larger for abiotic sphalerite oxidation than for abioticpyrite oxidation (e = �4.3‰ to �9.8‰ from Taylor et al., 1984; Bal-ci et al., 2007; Heidel and Tichomirowa, 2010). However, experi-ments were performed under different pH conditions, which mayinfluence the dominant O consumption process and, thus, the Oisotope enrichment factor.

The e value associated with the reduction of O2 was �6.4 ± 6.0‰

in the experiments at initial pH 2 from day 57–100 (Fig. 6) and dif-fers significantly from the eSO4–O2 value of ca.�22‰. Thus, the e va-lue associated with the reduction of O2 indicates that the dominantO consumption process differs for experiments at initial pH 2 andinitial pH 6. According to reaction (3), the observation of a slightH2S odor and the initially high Zn and Fe release may indicate thatsphalerite was additionally dissolved non-oxidatively at the begin-ning of the experiments at initial pH 2. Therefore, the e value asso-ciated with the reduction of O2 may increasingly reflect theoxidation of dissolved (reaction (5)) or gaseous H2S.

Oxygen isotope enrichment processes do not only occur duringthe consumption of O2, but also during the incorporation of water-derived O into SO2�

4 (eSO4–H2O). The eSO4–H2O value can be experi-mentally determined by anaerobic experiments with Fe(III) asoxidant (reaction (2)). Accordingly, O in SO2�

4 is derived exclusivelyfrom water, and the O isotope difference between SO2�

4 and water(D18OSO4–H2O = d18OSO4–d18OH2O) is solely caused by O isotopeenrichment processes between SO2�

4 and water. Balci et al. (2003)determined D18OSO4–H2O = eSO4–H2O = 7.9‰ from anaerobic sphaler-ite oxidation experiments. In contrast, D18OSO4–H2O values in aero-bic experiments do not only result from O isotope enrichmentprocesses between SO2�

4 and water (eSO4–H2O), but may also

2256 C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259

represent a proportion of O2 in SO2�4 . Therefore, the smallest

D18OSO4–H2O value from the present work (observed after 100 daysin the experiment at initial pH 2 and initial d18OH2O = �8.7‰) indi-cates that eSO4–H2O did not exceed 4.1‰ and may be even smaller.The value of 4.1‰ should be the maximum eSO4–H2O value, whichdiffers significantly from eSO4–H2O = 7.9‰ from Balci et al. (2003),but is similar to eSO4–H2O = 0.0–4.0‰ from pyrite and chalcopyriteoxidation (Taylor et al., 1984; Taylor and Wheeler, 1994; Balciet al., 2007; Mazumdar et al., 2008; Heidel and Tichomirowa,2011b; Thurston et al., 2010). Thus, it is assumed that eSO4–H2O liesbetween 0.0‰ and 4.1‰ for aerobic abiotic sphalerite oxidationunder neutral pH conditions.

4.4. Sphalerite dissolution mechanisms

Sulfide oxidation is an electrochemical process that consists ofanodic reactions (electron release from sulfide sulfur), electrontransfer (within the mineral), and cathodic reactions (electronacceptance by an oxidant). The changing d18OSO4 and d34SSO4 valuesin experiments at both initial pH values indicated a change in theoxidation mechanisms. During the first 20 days, the exclusiveincorporation of O2 into SO2�

4 may have resulted from sphaleriteoxidation by O2 (reaction (1)). Reaction (2) implies that onlywater-derived O is incorporated into SO2�

4 during sphalerite oxida-tion by Fe(III). Thus, an increasing proportion of water-derived O inSO2�

4 may indicate that sphalerite was oxidized by Fe(III) at leastfrom day 57 to 100.

Moreover, e values associated with the reduction of O2 differedfor experiments at initial pH 2 and 6 (day 57–100). This observa-tion may be a consequence of the initial non-oxidative dissolutionof sphalerite which occurred at the beginning of the experiments atinitial pH 2.

4.4.1. Non-oxidative dissolution of sphaleriteAs shown in reaction (3), the non-oxidative sphalerite dissolu-

tion results in the release of metal ions and the formation of H2Swhile consuming protons. Accordingly, an initial sharp pH in-crease, initially high concentrations of Zn and dissolved Fe(II) anda permanent slight H2S odor during the experiments at initial pH2 indicated that, in addition to the oxidation, sphalerite was dis-solved non-oxidatively at the beginning of these experiments.The S deficit in solutions from experiments at initial pH 2 also sug-gests that S was ‘‘lost’’ by degassing of H2S into the headspace(reaction (4)).

When similar amounts of SO2�4 were formed in experiments at

different initial pH, d18OSO4 values from experiments at initial pH2 suggested a higher proportion of O2 in SO2�

4 produced from day57 to 100 (Fig. 3). This observation may be a consequence of theinitial non-oxidative sphalerite dissolution. Produced H2S both inthe headspace and in solution could be oxidized by O2 which re-sults in the incorporation of O2 into intermediate S species andthe product SO2�

4 .Hydrogen sulfide may be oxidized via the intermediate S spe-

cies sulfite (Byerley and Scharer, 1992). Oba and Poulson (2009)observed that the e value associated with the reduction of O2

was �3.0‰ during the oxidation of dissolved SO2�3 to SO2�

4 . Thus,the observed e value associated with the reduction of O2

(�6.4 ± 6.0‰) in experiments at initial pH 2 may indicate thatthe oxidation of H2S via sulfite was an important O2 consumingprocess especially from day 57 to 100.

During the evasion of H2S, 32S is enriched in gaseous H2S rela-tive to dissolved H2S (Fry et al., 1986). Thus the dissolved H2Sremaining in solution should be enriched in 34S. The oxidation of34S-enriched dissolved H2S may explain decreasing d34S values.However, the re-dissolution of 32S-enriched H2S or intermediateS species may also affect the d34S value of the produced SO2�

4 .

Therefore, future work is necessary to investigate the multipleoxidation pathways of gaseous and dissolved H2S. In addition, iso-tope enrichment processes associated with the multiple oxidationreactions and evaporation and re-dissolution of intermediate Sspecies should be studied.

4.4.2. Sphalerite oxidation by molecular oxygenDuring the first 20 days of all experiments, only O2 was incorpo-

rated into the produced SO2�4 . Thus, the anodic reaction (i.e., elec-

tron release from sphalerite S) may be simply described by:

ZnSþ 4O2� ! Zn2þ þ SO2�4 þ 8e�: ð16Þ

Thurston et al. (2010) reviewed sulfide oxidation studies andsuggested that an enrichment of 32S in SO2�

4 relative to the sulfideoccurred during the non-biological sulfide oxidation by Fe(III).Therefore, the lack of S isotope enrichment in SO2�

4 relative tosphalerite during the first 20 days of the experiments indicatedthat sphalerite was not oxidized by Fe(III). Moses et al. (1987) pro-posed that sphalerite may be directly attacked by paramagnetic O2

due to sphalerite impurities of transition metals (e.g., Fe), whichchange the diamagnetism of sphalerite into paramagnetism.Accordingly, the relatively high Fe content of the sphalerite maypromote the attack by O2. Thus, O2 may be the predominant elec-tron acceptor from sphalerite S in the corresponding cathodicreaction:

2O2 þ 8e� ! 4O2�: ð17Þ

The combination of the reactions (16) and (17) results in theoverall reaction (1) for sphalerite oxidation by O2. However, oxida-tion reactions proceed stepwise via intermediate S species becauseonly one or two electrons can be transferred in one step (Basoloand Pearson, 1967). Weisener et al. (2004) proposed an oxidationmechanism for an Fe-rich sphalerite that begins with the initialsurface protonation of sphalerite S sites:

2ZnSþ 2Hþ ! 2Zn2þ þ 2HS�ðsÞ: ð18Þ

Afterwards, monosulfide species S2� (oxidation state: �2) onthe protonated sphalerite surface are oxidized to SO2�

4 (+6) viadisulfide species (�1) (Weisener et al., 2004):

2HS�ðsÞ þ 0:5O2 þ 2Hþ ! H2S2ðsÞ þH2O; ð19Þ

H2S2ðsÞ þ 3:5O2 þH2O! 4Hþ þ 2SO2�4 : ð20Þ

In general, these oxidation steps are in agreement with theobservations (XPS, dissolved S species). Single oxidation steps fromdisulfide to SO2�

4 and the location of these steps were not specifiedby Weisener et al. (2004). According to reaction (20), O in the pro-duced SO2�

4 consists of 87.5% O2 and 12.5% water-derived O. How-ever, the O isotope data from day 2 to 20 indicated that reaction(20) does not reflect the true proportion of both O sources in SO2�

4 .Schippers and Sand (1999) proposed an oxidation mechanism

for acid-soluble sulfides (‘‘polysulfide mechanism’’) where elemen-tal S, which is formed via intermediary polysulfides, is the mainoxidation product. Although the formation of polysulfides and ele-mental S on the sphalerite surface may be indicated by increasingpH values and decreasing reaction rates, polysulfides were de-tected only on the surface of one sample. Elemental S could notbe detected by XPS and EDX investigations maybe due to slowreaction rates at low temperatures. Sulfate is formed by side reac-tions via intermediate S species such as thiosulfate, polythionatesand sulfite (Schippers and Sand, 1999), which is in agreement withthe detection of thiosulfate and sulfite at least in experiments atinitial pH 6 (Table 3). In contrast to acid-insoluble sulfides (e.g.,pyrite), the metal-S bond of acid-soluble sulfides breaks before sul-fidic S is oxidized (Sand et al., 2001). Accordingly, the absence of

C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259 2257

SO2�4 on sphalerite surfaces after 2 and 20 days and the presence of

dissolved thiosulfate and sulfite may indicate that intermediate Sspecies are released from sphalerite surfaces into solution wherethey are finally oxidized to SO2�

4 .Much lower sulfite and thiosulfate concentrations in experi-

ments at initial pH 2 (compared with initial pH 6) indicated thatfurther oxidation mechanisms, which are probably associated withthe initial non-oxidative sphalerite dissolution, occurred duringthese experiments.

Schippers and Sand (1999) proposed that the ‘‘polysulfidemechanism’’ is valid for the oxidation of all acid-soluble sulfides(e.g., sphalerite, galena, chalcopyrite) whereas acid-insoluble sul-fides (e.g., pyrite) are oxidized by another mechanism (‘‘thiosulfatemechanism’’). Heidel and Tichomirowa (2011a) conducted abioticgalena oxidation experiments at both initial pH 2 and initial pH6. Although some methodological modifications are necessary toobtain a comprehensive isotope database, important findingsmay be discovered. The d18OSO4 values remained more or less con-stant, but were clearly dependent on d18OH2O values. Accordingly,O in SO2�

4 derived largely from water molecules (71 ± 25%, 1r,n = 6), but minor amounts of O2 could also be observed. Thed34SSO4 values indicated an enrichment of 32S in SO2�

4 (relative togalena), which increased with increasing pH. In contrast to thesphalerite oxidation, sulfite and thiosulfate concentrations weresimilar for experiments at both initial pH 2 and initial pH 6. Thus,data from this work and Heidel and Tichomirowa (2011a) indicatethat oxidation mechanisms are different for sphalerite and galena.Furthermore, Thurston et al. (2010) suggested that the oxidation ofthe acid-soluble sulfide chalcopyrite is more similar to pyrite oxi-dation than sphalerite oxidation. These results show that there isno general oxidation mechanism for acid-soluble sulfides.

4.4.3. Sphalerite oxidation by ferric ironFrom day 57 to 100, water-derived O was incorporated into the

produced SO2�4 , which may be simply described by the following

anodic reaction:

ZnSþ 4H2O! Zn2þ þ SO2�4 þ 8Hþ þ 8e�: ð21Þ

According to Thurston et al. (2010), the enrichment of 32S inSO2�

4 relative to sphalerite from day 57 to 100 (Fig. 5) may indicatethat Fe(III) was the dominant electron acceptor from sphalerite Ssites. In the course of the experiments, Fe(II) was released fromthe sphalerite lattice and oxidized to Fe(III) by molecular O2:

Fe2þ þ O2 þ 4Hþ ! 4Fe3þ þ 2H2O: ð22Þ

Thus, a sufficient amount of Fe(III) was present from day 57 to100 and may accept electrons from sphalerite S during the cathodicreaction:

8Fe3þ þ 8e� ! 8Fe2þ: ð23Þ

The combination of the reactions (21) and (23) results in theoverall reaction (2) for sphalerite oxidation by Fe(III). Although dis-solved Fe species were not detected spectrophotometrically inexperiments lasting longer than 10 days, ICP-OES control measure-ments showed that small amounts of dissolved Fe occurred inexperiments finished after 100 days (Table 3). Presumably, Fe spe-cies were adsorbed on sphalerite surfaces where they could not bedetected by the measurements.

Sphalerite oxidation with Fe(III) as oxidant may proceed simi-larly to the mechanism described for pyrite oxidation at circumneutral pH values by Moses and Herman (1991). Accordingly, re-leased dissolved Fe(II) from sphalerite is immediately adsorbedon S sites of sphalerite and oxidized to Fe(III) by O2 (reaction(19)). An electron is then transferred from the sphalerite S site tothe adsorbed hydrous Fe(III). Thus, adsorbed hydrous Fe(III) is

reduced to Fe(II) and the cycle starts again. According to Mosesand Herman (1991), O is transferred from the hydration sphereof the adsorbed Fe to sphalerite S (anodic reaction). Therefore, Oin the oxidized S species originates from water. The cycle is re-peated until oxidized S species (e.g., thiosulfate, polythionates, sul-fate) dissociate from the surface due to a higher stability insolution than remaining at the surface (Moses and Herman,1991). The accumulation of thiosulfate in solution suggests thatthiosulfate (and also small amounts of sulfite) was released fromthe sphalerite surface and finally oxidized to SO2�

4 in solution.Sphalerite oxidation with Fe(III) as oxidant may also proceed as

a heterogeneous reaction between the solids sphalerite and Fe(III)oxyhydroxides. This mechanism should be more important inexperiments at initial pH 6 where higher pH values were achieved.The observed pH values in these experiments (Table 3: 7.3–7.9) re-sulted in a decreased solubility of Fe(III), which can be precipitatedas Fe(III) oxyhydroxides (reaction (13)). If so, heterogeneous reac-tions between the sphalerite and ferric oxyhydroxides may occur.Heterogeneous reactions are generally characterized by lowerreaction rates. Accordingly, lower SO2�

4 concentrations were ob-served in experiments at initial pH 6 (compared with initial pH2) especially after more than 20 days of oxidation.

4.5. Application to field data

Sulfate in AMD sites is formed by oxidation of pyrite and othersulfides. AMD is typically associated with high Fe concentrationsdue to Fe(II) release from pyrite and additional release from othersulfides such as sphalerite. Because Fe(II)-oxidizing microorgan-isms often occur in AMD sites, Fe(II) is usually available as a sulfideoxidant even at acid pH. The present data indicated that sphaleriteoxidation by Fe(III) results in the predominant incorporation ofwater-derived O into the produced SO2�

4 . Therefore, sphalerite oxi-dation produces SO2�

4 with d18OSO4 values slightly higher thand18OH2O values (in the presence of high Fe concentrations andmicroorganisms at acid pH) which may be indistinguishable fromSO2�

4 produced by oxidation of other sulfides (e.g., pyrite). Forexample, SO2�

4 from small AMD pools in the abandoned mine inFreiberg (Germany) results from the oxidation of different sulfides(pyrite, sphalerite, galena, minor arsenopyrite, chalcopyrite, Haub-rich and Tichomirowa, 2002). The d18OSO4 values slightly higherthan d18OH2O values suggest that sulfides are oxidized by Fe(III)which results in the predominant incorporation of water-derivedO into the produced SO2�

4 . Thus, d18OSO4 values alone do not allowconclusions regarding the oxidized sulfides.

However, Hubbard et al. (2009) observed relatively high d18OSO4

values (5.5–8.1‰) from AMD field studies. Considering that thed18OH2O values ranged from �7.0‰ to �4.8‰, they concluded thatsuch high d18OSO4 values cannot be explained by the commonlyproposed pyrite oxidation mechanisms. Although pyrite oxidationaccounts for more than 93% of the dissolved SO2�

4 , sphalerite alsooccurred in the study area (Hubbard et al., 2009). The presentsphalerite oxidation experiments have shown that d18OSO4 valuesmay be relatively high (0.4–4.0‰) if sphalerite is exclusively oxi-dized by O2. Although significant amounts of Fe(III) were detectedby Hubbard et al. (2009), sphalerite oxidation by O2 may (at leastin part) contribute to the observed high d18OSO4 values. However,sphalerite should be predominantly oxidized by Fe(III) if sufficientamounts of it are present.

5. Conclusions

Aerobic abiotic sphalerite oxidation experiments were per-formed at different initial pH values (2 and 6) for different lengthsof time (2–100 days). Oxygen and S isotope investigations of the

2258 C. Heidel et al. / Applied Geochemistry 26 (2011) 2247–2259

produced SO2�4 contribute to the understanding of sphalerite oxi-

dation mechanisms, which helps to interpret field data.An initially rapid oxidation followed by a slower oxidation was

observed during all experiments. Changing reaction rates aremainly attributed to the initial dissolution of ultrafine sphaleriteparticles. Furthermore, the release of initial SO2�

4 and the passiv-ation of the sphalerite surface by polysulfides and Fe(III) oxyhy-droxides may contribute to decreasing reaction rates. In additionto changing oxidation rates, O and S isotopes of the producedSO2�

4 changed during the experiments. However, the time delayindicated that the change of the oxidation rate (occurred betweenday 5 and 2 in experiments at initial pH 2 and 6, respectively) andthe change of the isotopic composition of SO2�

4 (occurred after20 days) proceeded independently.

Oxygen and S isotope investigations indicated that sphaleritemay be oxidized by both O2 and Fe(III). During the first 20 daysof the experiments at both initial pH values, O2 was the exclusiveO source of SO2�

4 . Furthermore, the lack of S isotope enrichmentprocesses between SO2�

4 and sphalerite indicated that O2 was theelectron acceptor from sphalerite S. As the oxidation proceeded, asufficient amount of released Fe(II) was oxidized to Fe(III) by O2.Thus, electrons could be transferred from sphalerite S sites to ad-sorbed hydrous Fe(III). Decreasing d18OSO4 values indicated that Ofrom the hydration sphere of Fe was incorporated into the pro-duced SO2�

4 . The enrichment of 32S in SO2�4 relative to the sphalerite

may also result from sphalerite oxidation by Fe(III).Thiosulfate and sulfite are important dissolved intermediate S

species during sphalerite oxidation independent of the oxidant.The absence of SO2�

4 on sphalerite surfaces after 2 and 20 days ofoxidation confirmed that intermediate S species were releasedfrom the sphalerite surface into solution where they are finally oxi-dized to SO2�

4 (as proposed by Sand et al, 2001).The incorporation of O2 into SO2�

4 during the oxidation of sphal-erite was associated with an O isotope enrichment factor eSO4–O2 ofca. �22‰. The e value associated with the reduction of O2 was�6.4 ± 6.0‰ in the experiments at initial pH 2 from 57–100 days(Fig. 6) and differs significantly from the eSO4–O2 value of ca.�22‰. This observation may be a consequence of the initial non-oxidative dissolution of acid-soluble sphalerite and the subsequentoxidation of both dissolved and gaseous H2S.

Oxygen 18 was preferentially incorporated from water mole-cules into SO2�

4 , which resulted in an O isotope enrichment factorof eSO4–H2O 6 4.1‰ which differs from the previously determinedvalue of eSO4–H2O = 7.9‰ (Balci et al., 2003). However, eSO4–

H2O 6 4.1‰ is similar to eSO4–H2O = 0.0–4.0‰ from pyrite and chal-copyrite oxidation experiments. Results from this work, Heidel andTichomirowa (2011a), and Thurston et al. (2010) suggest that thereis no general oxidation mechanism for acid-soluble sulfides.

Acknowledgements

We thank Rositta Liebscher for supporting us with the experi-ments and for carrying out the isotope measurements. We aregrateful for the XPS measurements provided by Michal Luteckifrom the University of Leipzig. This study was funded by the Ger-man Research Foundation (DFG). We thank an anonymous re-viewer, the Associate Editor Robert R. Seal and the ExecutiveEditor Ron Fuge for their helpful comments which substantiallyimproved the manuscript.

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