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As(V) and As(III) reactions on pristine pyrite and on surface-oxidized pyrite
Journal of Colloid and Interface Science 388 (2012) 170–175
Contents lists available at SciVerse ScienceDirect
Journal of Colloid and Interface Science www.elsevier.com/locate/jcis
As(V) and As(III) reactions on pristine pyrite and on surface-oxidized pyrite Fenglong Sun a, Brian A. Dempsey a,⇑, Kwadwo A. Osseo-Asare b a b
Department of Civil and Environmental Engineering, Penn State University, UP, United States Department of Materials Science and Engineering and of Geo-Environmental Engineering, Penn State University, UP, United States
a r t i c l e
i n f o
Article history: Received 21 February 2012 Accepted 9 August 2012 Available online 23 August 2012 Keywords: Pyrite Arsenic Arsenate Arsenite Adsorption Reduction XANES
a b s t r a c t Reactions of As(III) and As(V) with pyrite were investigated using pristine pyrite (produced and reacted in a rigorously anoxic environment with PO2 < 108 atm) and using surface-oxidized pyrite (produced under anoxic conditions, exposed to air, then stored and reacted under rigorously anoxic conditions). Results with surface-oxidized pyrite were similar to previously reported arsenic-pyrite results. However As(III) adsorbed over a broader pH range on pristine pyrite than on surface-oxidized pyrite, As(V) adsorbed over a narrower pH range on pristine pyrite than on surface-oxidized pyrite, and adsorbed As(V) on pristine pyrite was reduced to As(III) but adsorbed As(V) was not reduced with surface-oxidized pyrite. Reduction of As(V) with pristine pyrite was first-order in total As(V), Fe(II) was released, and sulfur was oxidized. The proposed mechanism for pyrite oxidation by As(V) was similar to the published mechanism for oxidation by O2 and rates were compared. The results can be used to predict the removals of As(V) and As(III) on pyrite in continuously anoxic environments or on pyrite in intermittently oxic/anoxic environments. Rigorous cleanup and continuous maintenance of strictly anoxic conditions are required if commercial or produced pyrites are to be used as surrogates for pristine pyrite. Ó 2012 Elsevier Inc. All rights reserved.
1. Introduction Arsenic has caused human health problems world-wide, notably in Bangladesh, India, and Taiwan [1]. Dissolved arsenic is typically high for intermediate redox potential that results in dissolution of iron (hydr)oxides, usually low in oxic surface or ground water when iron (hydr)oxides are present [2–5], and also usually low when iron sulfide minerals are present [6,7]. Pyrite is the most abundant sulfide mineral in the earth’s crust. Reactions of pyrite with environmental contaminants have been studied extensively [8] and it has been shown that pyrite affects the speciation and mobility of Mn(II), Cu(II), Sr(II), and As [9–11]. Several studies have focused on the reactions of As(III) or As(V) with pyrite and other iron sulfides [12–17]. Farquhar et al. [12] stated that As(V)/As(III) formed outer-sphere surface complexes with mackinawite and pyrite. Paikaray et al. [13] reported substantial adsorption of As(III) on pyritic shales and the extent of sorption was correlated with pyrite content. Bostick and Fendorf [14] studied As(III) adsorption on troilite (FeS) and pyrite (FeS2) and reported surface precipitation of arsenopyrite (FeAsS). Kim and Batchelor [16] found that As(III) and As(V) were removed by synthetic pyrite due to formation of As2S3 or As4S4 surface phases. Wolthers et al. [17] studied As(III) uptake by pyrite in a continuous-flow system and reported that arsenian pyrite was formed. ⇑ Corresponding author. E-mail address: [email protected] (B.A. Dempsey). 0021-9797/$ - see front matter Ó 2012 Elsevier Inc. All rights reserved. http://dx.doi.org/10.1016/j.jcis.2012.08.019
The variety of products reported in these studies may be due to differences in the characteristics of the pyrite materials. Pyrite is thermodynamically stable in sufficiently reducing environments but it is unstable in oxidizing environments. Exposure to O2 or other oxidants can result in formation of ferric oxide or other solid phases on the pyrite surface [18]. In this context Jeong et al. [19] showed that partial oxidation of mackinawite (FeS) affected the mobility of arsenic. Commercial sources of pyrite have been used in most studies involving arsenic species and pyrite and the surface of commercial pyrite can be oxidized due to grinding, washing, or storage [20–22]. Therefore the objectives of this work were to produce pristine pyrite and surface-oxidized pyrite and then to study the adsorption of As(III) and the adsorption and reduction of As(V) with the two pyrite materials. Pristine pyrite represented pyrite found in continuously anoxic sediments or groundwaters while surface-oxidized pyrite represented pyrite in intermittently oxic/anoxic sediments or groundwaters or commercial pyrite having been exposed to grinding or to air or other oxidants. 2. Material and methods 2.1. Materials and characterization Chemicals were reagent grade or synthesized from reagent grade materials. Solutions were prepared using Milli-Q water (>10.2 MX cm1 resistivity) that was deoxygenated by bubbling
F. Sun et al. / Journal of Colloid and Interface Science 388 (2012) 170–175
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Fig. 1. SEM image of (a) synthetic pyrite, (b) oxygen treated pyrite with, (c) EDS analysis of synthetic pyrite and (d) EDS analysis of oxygen treated pyrite.
99.9% nitrogen gas for at least 5 h and then storage for at least 2 days in an anaerobic chamber (Coy Laboratory Products, Inc., 5% H2 in N2 with Pd catalyst). Pyrite was synthesized [23] by combining equal volumes of 0.033 M FeCl3 (J.T. Baker) and 0.067 M Na2S (Alfa-Aesar) solutions at pH 3.7 in an anaerobic chamber. The mixture was aged for at least 3 days prior to experiments. XRD (PANalytical Theta-2-Theta Powder Diffractometer) showed that the initial products contained mackinawite, which was removed by washing with 1 M HCl followed by three rinses with deoxygenated water. Subsequent XRD results showed that mackinawite was removed and the purified sample appeared to be single-crystalline pyrite. Pyrite was freshly prepared before each experiment. Surface-oxidized pyrite was prepared as described by Fornasiero et al. [24]. Pristine pyrite was removed from the anaerobic chamber and bubble aerated for 12 h at pH 4–5. The suspension was centrifuged, transferred back to the anaerobic chamber, resuspended in de-oxygenated water, and stored. The morphology and elemental composition of pristine and surface-oxidized pyrites were characterized with scanning electron microscopy (SEM) (FEI Quanta 200) and energy dispersive X-ray spectroscopy (EDS). Surface areas were determined by BET N2. All pyrite stock suspensions were 1 g/L. 2.2. Iron, arsenic and sulfate analysis Dissolved concentrations were measured after centrifugation (13.4 krpm, 5 min). Ferrous and total iron (after treatment with hydroxylamine chloride) concentrations were measured with
ferrozine [25]. As(V) and total arsenic (after oxidation with potassium iodate at pH 1) concentrations were measured using molybdenum blue at 865 nm [26]. Adsorbed arsenic was extracted (24 h) with 5 M NaOH which recovers >95% of adsorbed arsenic without change in valence [27]. Sulfate was analyzed by Dionex DX-100 ion chromatograph with an AS4A-SC analytical column connected to a conductivity detector. 2.3. Experimental procedure Experiments were conducted in an anaerobic chamber at room temperature. A stock solution of As(V) or As(III) was added to a pyrite suspension and NaOH or HCl was added to adjust the pH. For variable pH experiments, 40 mL Pyrex glass tubes were used as batch reactors with 48 h reaction time, which was shown to be sufficient for equilibration. For reaction rate experiments, 250 mL Pyrex glass bottles were connected to an oxygen trap as described by Jeon et al. [28] and 10 mL of suspension was removed for centrifugation at desired times. 2.4. X-ray absorption analysis Samples for XANES (X-ray absorption near-edge structure) analyses were prepared by reacting 200 lM As(V) with 1 g/L synthetic pyrite or surface oxidized pyrite at pH 3.5 or 6.0 for 5 days. The samples were rinsed, dried, and packed in plastic sample holders sealed with Kapton tape in the anaerobic chamber. Arsenic K edge XANES spectra were collected at the GSECARS’s beamline (sector-13D) at the Advanced Photon Source in Argonne, IL. Si(1 1 1) crystal was
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Fig. 2. XRD patterns of pristine pyrite and surface oxidized pyrite.
used as a monochromator. The XANES spectra were collected from 11,800 to 12,200 eV in the fluorescence mode, using a Ge 16 element detector. Na2HAsO47H2O and NaAsO2 powders were used as As(V) and As(III) standards. XANES spectra were processed using the Iffefit software package (Athena and Artemis) [29]. 3. Results and discussions 3.1. Characterization of synthetic pyrite and surface-oxidized pyrite Fig. 1a shows the SEM image of pristine pyrite after purification. Elemental analysis from EDS in Fig. 1c indicated peaks of iron (Fe) and sulfur (S). The average ratio of Fe to S) was 1:2.09 based on six randomly selected locations, which was close to the stoichiometry in pyrite (FeS2). The surface area was 12.6 m2/g. Since the density of pyrite is 5.01 g/cm3 therefore the primary particle size (assuming spheres) was 0.048 lm. The average particle size from SEM was 1.0 lm, suggesting that pristine pyrite was an aggregate of smaller particles. Surface-oxidized pyrite has a core–shell structure [24]. Fig. 1b shows the SEM image of surface-oxidized pyrite and indicates morphological changes occurred due to aeration and surface oxidation. Based on similar analyses as reported in the prior paragraph, the average ratio (Fe/S) was 1:1.40, the surface area was 7.7 m2/g, primary particle diameter was 0.078 lm, and SEM particle size was 1.5 lm. The elemental analysis in Fig. 1d showed a small oxygen peak. These results indicated formation of an iron oxide phase on the pyrite surface but pyrite was the only XRD-identified mineral phase for either pristine or surface-oxidized samples (Fig. 2). 3.2. Reactions of As(III) on pristine and on surface-oxidized pyrite As(III) was adsorbed on pristine pyrite from pH 2 to 12 as shown in Fig. 3. Adsorption decreased for pH > 12 perhaps related to pH approaching the pKa3 of H3AsO3. As(III) adsorbed over a narrower pH range on surface-oxidized pyrite, with maximum adsorption between pH 3.5 and 10. The narrower adsorption envelope on surface-oxidized pyrite was similar to that reported for As(III) on goethite or ferrihydrite [2,3]. Fig. 3 also shows reported results for As(III) adsorption on commercial pyrite that had been washed in a 0.01 M sulfide solution, used in a 10/90 H2/N2 environment, and using a similar adsorbate to adsorbent ratio as in the other curves in Fig. 3 [14]. Those authors found evidence of surface oxidation of sulfur and iron, possibly due to material preparation prior to analysis by X-ray photoelectron spectroscopy (XPS). They noted
Fig. 3. As(III) removal using pristine pyrite, surface-oxidized pyrite, and reduced commercial pyrite. Conditions: 45 lM As, room temperature, and 48 h reaction time.
that ‘‘limited surface oxidation may influence As(III) adsorption, potentially altering the extent of sorption or the redox state of sorbed As.’’ The similarity of the As(III) adsorption envelope on the sulfide-treated commercial pyrite [14] and our results on surface-oxidized pyrite suggest that the pyrite surfaces in both experiments were oxidized, while the broader adsorption envelope for As(III) on our pristine pyrite suggests that the surface was different, perhaps a true pyrite surface.
3.3. Reactions of As(V) on pristine pyrite and on surface-oxidized pyrite Fig. 4a shows that As(V) was adsorbed on pristine pyrite for acidic pH, and that all of the adsorbed As(V) at pH 3 and 3.5 was reduced to As(III). As(V) adsorption and the extent of reduction to As(III) decreased at higher pH values (6.3, 7.1, 8.3, 10.0, 11.0). None of the As(III) produced by reduction of As(V) was released into the dissolved phase. The apparent incorporation of arsenic into the solid phase during reduction of As(V) with pristine pyrite indicates that arsenic removal in anaerobic environments could be greater than would be predicted based on simple adsorption of As(III). The reaction of As(V) with surface-oxidized pyrite exhibited an entirely different reaction mechanism, as shown in Fig. 4b. There was significant removal of As(V) up to pH 9 using 1 g/L surface-oxidized pyrite but removal of As(V) peaked around pH 6 with 0.1 g/L pyrite. There was no reduction of As(V) to As(III). The reactions of As(V) with surface oxidized pyrite were similar to those reported for adsorption of As(V) onto hydrous ferric oxides [3,30]. XANES spectra are shown in Fig. 5 for As(V) and As(III) standards, for As(V) reaction with pristine pyrite (pH = 3.5), and for As(V) reaction with surface-oxidized pyrite (pH = 6.0). Background removal and normalization were performed for the pyrite samples. A white line (first most intense absorption peak) for As(V) was at 11874.8 eV and for As(III) at 11871 eV. The XANES results confirmed that As(V) was reduced to As(III) with pristine pyrite but not with surface-oxidized pyrite. Some As(V) was found in the pristine pyrite with XANES which the wet chemistry analyses showed that all As(V) adsorbed onto pristine pyrite was reduced to As(III). The measurement of some As(V) with XANES could have been due to oxidation of As(III) by the synchrotron beam as observed by Bostick and Fendorf [14]. The rate of As(V) reduction by pristine pyrite is shown in Fig. 6, as a function of pH and initial As(V) concentration. The arsenic removal rate increased with decreasing pH. The inset in Fig. 6 shows
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Fig. 5. As K-edge XANES spectra of As(V)/As(III) standard samples, and arsenic on the solid phase from As(V) reaction on pristine pyrite (pH = 3.5) and surface oxidized pyrite (pH = 6).
Fig. 4. As(V) removal using (a) pristine pyrite, showing adsorbed and dissolved As(III) and As(V) and (b) surface-oxidized pyrite (0.1 g/L or 1 g/L) which did not reduce As(V) to As(III). Conditions: 45 lM arsenic, room temperature, and 48 h reaction time.
log[As(V)dissolved] versus time for several of the conditions and demonstrates that the reaction was first order with respect to the concentration of dissolved As(V) at pH 3.5 and 3.7. However the reaction stopped short of equilibration at pH 6.7 possibly due to production of an impermeable surface layer of FeOOH that inhibited contact of As(V) with pyrite. 3.4. Stoichiometry and mechanisms for As(V) reaction with pristine pyrite The experimental stoichiometry of the As(V) and pristine pyrite reaction at pH 3.5 is shown in Fig. 7. The incremental Fe(II) concentration was determined after compensating for the dissolution of pyrite in the absence of arsenate. Thiosulfate was not detected and sulfate appeared to be the only oxidized sulfur species. Williamson and Rimstitdt [31] reported that thiosulfate was detected in very small amounts and that polysulfides and sulfite were not detected during oxidation of pyrite by O2. Fe(III) was not produced, which is consistent with previous studies using relatively weak oxidants. Naveau et al. [10] reported
Fig. 6. Total As(V) concentration (dissolved + adsorbed) versus time at different pH values. The inset figure shows log (total As(V)) versus time. Initial pyrite was 1 g/L.
that Cu(II) was reduced by pyrite to Cu(I) and only sulfur was oxidized. Naveau et al. [32] reported that Se(IV) was reduced by pyrite and only sulfur was oxidized. However Demoisson et al. [33] found that Cr(VI) oxidized pyrite to sulfate and Fe(III). Plots of the logarithms of dissolved As(V), Fe(II), and sulfate concentrations versus time in Fig. 7b produced R2 values of 0.99, 0.89, 0.91. The slopes were 0.110 s1 for As(V), 0.031 s1 for sulfate, and 0.014 s1 for Fe(II). The ratio of these slopes was 7/2/0.9, which is close to the theoretical stoichiometric ratio of 7/2/1 for the reaction.
7H2 AsO4 þ FeS2 þ 5Hþ þ H2 O ¼ Fe2þ þ 2SO2 4 þ 7H3 AsO3
DG0
¼ 42:99 kJ=mol
ð1Þ
The reaction rate constant did not change over time at pH 3.5. Williamson and Rimstidt [31] reported a decrease in the rate constant for pyrite oxidation by O2 with time, possibly due to
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Fig. 8. The rate of pristine pyrite oxidation by As(V) (solid lines) is compared with oxidation by O2 and dissolved Fe(III) (data from Williamson and Rimstidt [31]). Rates were normalized using the pyrite surface area (mol m2 s1). Conditions: 0.2 atm O2, dissolved Fe(III) in equilibrium with ferrihydrite, 45 lM or 0.133 lM As(V).
4. Conclusions and environmental implications
Fig. 7. (a) Concentrations of dissolved As(V), incremental dissolved Fe(II) (after subtraction of pyrite self-dissolution), and dissolved sulfate versus time during the reaction between As(V) and pristine pyrite at pH 3.5; (b) logarithm of As(V), sulfate, and incremental Fe(II) concentrations plotted versus time. Initial As(V) was 200 lM and initial pyrite 1 g/L.
formation of an oxidized shell. Their initial rate constant was not restored after cleaning with concentrated HNO3 (an oxidizing agent) or EDTA solution (relatively weak chelator at acidic pH). They noted that similar reductions in oxidation rate had been previously observed after weathering or polishing of pyrite. As(V) was a weak oxidant for pyrite compared to O2 or Fe(III). This is illustrated in Fig. 8, showing reaction rates (normalized by pyrite mass) for oxidation by 45 lM As(V), 10 lg/L As(V) (the USEPA maximum contaminant level for arsenic in potable water), 0.2 atm O2, or dissolved Fe(III). The O2 and Fe(III) data are from Williamson and Rimstidt [31]. They suggested that the observed pH effects were due to changes in adsorbed O2 and Fe(II) and on solubility of Fe(III). Similarly the decreasing rate of pyrite oxidation by As(V) with increasing pH could be due to decreased adsorption of As(V) at higher pH, consistent with our observation that only adsorbed As(V) was able to reduce pyrite. The slope of log(r) versus pH for As(V) was 0.5. Fig. 8 illustrates that O2 must be rigorously excluded from anaerobic chambers that are used to study reactions involving pyrite and As(V) or other weak oxidants.
Pristine pyrite was produced, cleaned, and stored in a rigorously anoxic environment (PO2 < 108 atm) and ‘‘surface-oxidized’’ pyrite was produced from pristine pyrite that was exposed to air and subsequently stored and reacted under anoxic conditions. Both materials appeared to be pure pyrite based on XRD analyses, but the two materials exhibited substantially different reactivity towards As(III) and As(V). As(III) adsorbed over a broader pH range on pristine pyrite than on surface-oxidized pyrite. Adsorption of As(III) on surface-oxidized pyrite was similar to reported adsorption on pyrite that had been prepared and stored in less-rigorously anoxic conditions than in our experiments or on iron oxides. As(V) adsorbed over a narrower pH range on pristine pyrite than on surface-oxidized pyrite. Adsorption of As(V) on surfaceoxidized pyrite was similar to previously reported adsorption on iron oxides. All adsorbed As(V) on pristine pyrite was reduced to As(III) at pH < 4 but As(V) was not reduced at all to As(III) when using surface-oxidized pyrite. Half of adsorbed As(V) was reduced to As(III) at pH 6.3 after 2 days. Reduction of As(V) to As(III) with pristine pyrite was confirmed using XANES analyses. As(III) that was produced by reaction of As(V) with pristine pyrite was incorporated into the solid phases and was not released back to solution. This result suggested that reduction of As(V) by pristine pyrite could result in better immobilization than would be predicted based on experiments in which As(III) is adsorbed onto pyrite. Reduction of As(V) with pristine pyrite was first-order in total As(V), the rate of As(V) reduction decreased with increasing pH, and only sulfur was oxidized. The proposed mechanisms of pyrite oxidation by As(V) were similar to previously published mechanisms for the oxidation of pyrite by O2 and the logarithms of the rates of the two reactions were proportional to the free energy of the two reactions. Therefore oxidation of pyrite by As(V) is unlikely to be observed in the presence of very low partial pressures of O2 that could be present in glove bag environments without using extra precautions for removal of trace concentrations of O2. The results can be used to predict reactivity of arsenic with pyrite in either consistently anoxic environments or in intermittently
F. Sun et al. / Journal of Colloid and Interface Science 388 (2012) 170–175
anoxic environments. The results also are pertinent for laboratory studies in which it may be necessary to produce and maintain pyrite in rigorously anoxic environments in order to avoid experimental artifacts. Acknowledgment This research was partially supported by the Center for Environmental Kinetics Analysis at Penn State University (NSF Grant No. Che-0431328). References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10]
P.L. Smedley, D.G. Kinniburgh, Appl. Geochem. 17 (2002) 517. B.A. Manning, S.E. Fendorf, S. Goldberg, Environ. Sci. Technol. 32 (1998) 2383. C.P. Raven, A. Jain, R.H. Loeppert, Environ. Sci. Technol. 32 (1998) 344. S. Dixit, J.G. Hering, Environ. Sci. Technol. 37 (2003) 4182. J. Gimenez, J. de Pablo, M. Martinez, M. Rovira, C. Valderrama, J. Colloid Interface Sci. 348 (2010) 293. K.A. Sullivan, R.C. Aller, Geochim. Cosmochim. Acta 60 (1996) 1465. H.A. Lowers, G.N. Breit, A.L. Foster, J. Whitney, J. Yount, N. Uddin, A. Muneem, Geochim. Cosmochim. Acta 71 (2007) 2699. R. Murphy, D.R. Strongin, Surf. Sci. Rep. 64 (2009) 1. M.A. Huerta-Diaz, J.W. Morse, Geochim. Cosmochim. Acta 56 (1992) 2681. A. Naveau, F. Monteil-Rivera, E. Guillon, J. Dumonceau, J. Colloid Interface Sci. 303 (2006) 25.
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[11] P. Bonnissel-Gissinger, M. Alnot, J.J. Ehrhardt, P. Behra, Environ. Sci. Technol. 32 (1998) 2839. [12] M.L. Farquhar, J.M. Charnock, F.R. Livens, D.J. Vaughan, Environ. Sci. Technol. 36 (2002) 1757. [13] S. Paikaray, S. Banerjee, S. Mukherji, Curr. Sci. 88 (2005) 1580. [14] B.C. Bostick, S. Fendorf, Geochim. Cosmochim. Acta 67 (2003) 909. [15] E.J. Kim, B. Batchelor, Environ. Eng. Sci. 26 (2009) 1785. [16] E.J. Kim, B. Batchelor, Environ. Sci. Technol. 43 (2009) 2899. [17] M. Wolthers, I.B. Butler, D. Rickard, P.R.D. Mason, in: P.A. O’Day, D. Vlassopoulos, X. Meng, L.G. Benning, American Chemical Society, 2005, p. 60. [18] P.C. Singer, W. Stumm, Science 167 (1970) 1170. [19] H.Y. Jeong, Y.S. Han, S.W. Park, K.F. Hayes, Geochim. Cosmochim. Acta 74 (2010) 3182. [20] K.M. Rosso, D.J. Vaughan, Rev. Miner. Geochem. 61 (2006) 557. [21] D.M.C. Huminicki, J.D. Rimstidt, Appl. Geochem. 24 (2009) 1626. [22] J.D. Rimstidt, D.J. Vaughan, Geochim. Cosmochim. Acta 67 (2003) 873. [23] D.W. Wei, K. OsseoAsare, Colloid Surf. A-Physicochem. Eng. Asp. 118 (1996) 51. [24] D. Fornasiero, V. Eijt, J. Ralston, Colloid Surface 62 (1992) 63. [25] E. Viollier, P.W. Inglett, K. Hunter, A.N. Roychoudhury, P.V. Cappellen, Appl. Geochem. 15 (2000) 785. [26] J. Murphy, J.P. Riley, Anal. Chim. Acta 27 (1962) 31. [27] J.-H. Jang, B.A. Dempsey, Environ. Sci. Technol. 42 (2008) 2893. [28] B.-H. Jeon, B.A. Dempsey, R.A. Royer, W.D. Burgos, J. Environ. Eng. (2004) 1407. [29] B. Ravel, M. Newville, J. Synchrotron Radiat. 12 (2005) 537. [30] J. Gimenez, M. Martinez, J. Hazard. Mater. 141 (2007) 575. [31] M.A. Williamson, J.D. Rimsidt, Geochim. Cosmochim. Acta 58 (1994) 5443. [32] A. Naveau, F. Monteil-Rivera, E. Guillon, J. Dumonceau, Environ. Sci. Technol. 41 (2007) 5376. [33] F. Demoisson, M. Mullet, B. Humbert, J. Colloid Interface Sci. 316 (2007) 531.
Contents lists available at SciVerse ScienceDirect
Journal of Colloid and Interface Science www.elsevier.com/locate/jcis
As(V) and As(III) reactions on pristine pyrite and on surface-oxidized pyrite Fenglong Sun a, Brian A. Dempsey a,⇑, Kwadwo A. Osseo-Asare b a b
Department of Civil and Environmental Engineering, Penn State University, UP, United States Department of Materials Science and Engineering and of Geo-Environmental Engineering, Penn State University, UP, United States
a r t i c l e
i n f o
Article history: Received 21 February 2012 Accepted 9 August 2012 Available online 23 August 2012 Keywords: Pyrite Arsenic Arsenate Arsenite Adsorption Reduction XANES
a b s t r a c t Reactions of As(III) and As(V) with pyrite were investigated using pristine pyrite (produced and reacted in a rigorously anoxic environment with PO2 < 108 atm) and using surface-oxidized pyrite (produced under anoxic conditions, exposed to air, then stored and reacted under rigorously anoxic conditions). Results with surface-oxidized pyrite were similar to previously reported arsenic-pyrite results. However As(III) adsorbed over a broader pH range on pristine pyrite than on surface-oxidized pyrite, As(V) adsorbed over a narrower pH range on pristine pyrite than on surface-oxidized pyrite, and adsorbed As(V) on pristine pyrite was reduced to As(III) but adsorbed As(V) was not reduced with surface-oxidized pyrite. Reduction of As(V) with pristine pyrite was first-order in total As(V), Fe(II) was released, and sulfur was oxidized. The proposed mechanism for pyrite oxidation by As(V) was similar to the published mechanism for oxidation by O2 and rates were compared. The results can be used to predict the removals of As(V) and As(III) on pyrite in continuously anoxic environments or on pyrite in intermittently oxic/anoxic environments. Rigorous cleanup and continuous maintenance of strictly anoxic conditions are required if commercial or produced pyrites are to be used as surrogates for pristine pyrite. Ó 2012 Elsevier Inc. All rights reserved.
1. Introduction Arsenic has caused human health problems world-wide, notably in Bangladesh, India, and Taiwan [1]. Dissolved arsenic is typically high for intermediate redox potential that results in dissolution of iron (hydr)oxides, usually low in oxic surface or ground water when iron (hydr)oxides are present [2–5], and also usually low when iron sulfide minerals are present [6,7]. Pyrite is the most abundant sulfide mineral in the earth’s crust. Reactions of pyrite with environmental contaminants have been studied extensively [8] and it has been shown that pyrite affects the speciation and mobility of Mn(II), Cu(II), Sr(II), and As [9–11]. Several studies have focused on the reactions of As(III) or As(V) with pyrite and other iron sulfides [12–17]. Farquhar et al. [12] stated that As(V)/As(III) formed outer-sphere surface complexes with mackinawite and pyrite. Paikaray et al. [13] reported substantial adsorption of As(III) on pyritic shales and the extent of sorption was correlated with pyrite content. Bostick and Fendorf [14] studied As(III) adsorption on troilite (FeS) and pyrite (FeS2) and reported surface precipitation of arsenopyrite (FeAsS). Kim and Batchelor [16] found that As(III) and As(V) were removed by synthetic pyrite due to formation of As2S3 or As4S4 surface phases. Wolthers et al. [17] studied As(III) uptake by pyrite in a continuous-flow system and reported that arsenian pyrite was formed. ⇑ Corresponding author. E-mail address: [email protected] (B.A. Dempsey). 0021-9797/$ - see front matter Ó 2012 Elsevier Inc. All rights reserved. http://dx.doi.org/10.1016/j.jcis.2012.08.019
The variety of products reported in these studies may be due to differences in the characteristics of the pyrite materials. Pyrite is thermodynamically stable in sufficiently reducing environments but it is unstable in oxidizing environments. Exposure to O2 or other oxidants can result in formation of ferric oxide or other solid phases on the pyrite surface [18]. In this context Jeong et al. [19] showed that partial oxidation of mackinawite (FeS) affected the mobility of arsenic. Commercial sources of pyrite have been used in most studies involving arsenic species and pyrite and the surface of commercial pyrite can be oxidized due to grinding, washing, or storage [20–22]. Therefore the objectives of this work were to produce pristine pyrite and surface-oxidized pyrite and then to study the adsorption of As(III) and the adsorption and reduction of As(V) with the two pyrite materials. Pristine pyrite represented pyrite found in continuously anoxic sediments or groundwaters while surface-oxidized pyrite represented pyrite in intermittently oxic/anoxic sediments or groundwaters or commercial pyrite having been exposed to grinding or to air or other oxidants. 2. Material and methods 2.1. Materials and characterization Chemicals were reagent grade or synthesized from reagent grade materials. Solutions were prepared using Milli-Q water (>10.2 MX cm1 resistivity) that was deoxygenated by bubbling
F. Sun et al. / Journal of Colloid and Interface Science 388 (2012) 170–175
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Fig. 1. SEM image of (a) synthetic pyrite, (b) oxygen treated pyrite with, (c) EDS analysis of synthetic pyrite and (d) EDS analysis of oxygen treated pyrite.
99.9% nitrogen gas for at least 5 h and then storage for at least 2 days in an anaerobic chamber (Coy Laboratory Products, Inc., 5% H2 in N2 with Pd catalyst). Pyrite was synthesized [23] by combining equal volumes of 0.033 M FeCl3 (J.T. Baker) and 0.067 M Na2S (Alfa-Aesar) solutions at pH 3.7 in an anaerobic chamber. The mixture was aged for at least 3 days prior to experiments. XRD (PANalytical Theta-2-Theta Powder Diffractometer) showed that the initial products contained mackinawite, which was removed by washing with 1 M HCl followed by three rinses with deoxygenated water. Subsequent XRD results showed that mackinawite was removed and the purified sample appeared to be single-crystalline pyrite. Pyrite was freshly prepared before each experiment. Surface-oxidized pyrite was prepared as described by Fornasiero et al. [24]. Pristine pyrite was removed from the anaerobic chamber and bubble aerated for 12 h at pH 4–5. The suspension was centrifuged, transferred back to the anaerobic chamber, resuspended in de-oxygenated water, and stored. The morphology and elemental composition of pristine and surface-oxidized pyrites were characterized with scanning electron microscopy (SEM) (FEI Quanta 200) and energy dispersive X-ray spectroscopy (EDS). Surface areas were determined by BET N2. All pyrite stock suspensions were 1 g/L. 2.2. Iron, arsenic and sulfate analysis Dissolved concentrations were measured after centrifugation (13.4 krpm, 5 min). Ferrous and total iron (after treatment with hydroxylamine chloride) concentrations were measured with
ferrozine [25]. As(V) and total arsenic (after oxidation with potassium iodate at pH 1) concentrations were measured using molybdenum blue at 865 nm [26]. Adsorbed arsenic was extracted (24 h) with 5 M NaOH which recovers >95% of adsorbed arsenic without change in valence [27]. Sulfate was analyzed by Dionex DX-100 ion chromatograph with an AS4A-SC analytical column connected to a conductivity detector. 2.3. Experimental procedure Experiments were conducted in an anaerobic chamber at room temperature. A stock solution of As(V) or As(III) was added to a pyrite suspension and NaOH or HCl was added to adjust the pH. For variable pH experiments, 40 mL Pyrex glass tubes were used as batch reactors with 48 h reaction time, which was shown to be sufficient for equilibration. For reaction rate experiments, 250 mL Pyrex glass bottles were connected to an oxygen trap as described by Jeon et al. [28] and 10 mL of suspension was removed for centrifugation at desired times. 2.4. X-ray absorption analysis Samples for XANES (X-ray absorption near-edge structure) analyses were prepared by reacting 200 lM As(V) with 1 g/L synthetic pyrite or surface oxidized pyrite at pH 3.5 or 6.0 for 5 days. The samples were rinsed, dried, and packed in plastic sample holders sealed with Kapton tape in the anaerobic chamber. Arsenic K edge XANES spectra were collected at the GSECARS’s beamline (sector-13D) at the Advanced Photon Source in Argonne, IL. Si(1 1 1) crystal was
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Fig. 2. XRD patterns of pristine pyrite and surface oxidized pyrite.
used as a monochromator. The XANES spectra were collected from 11,800 to 12,200 eV in the fluorescence mode, using a Ge 16 element detector. Na2HAsO47H2O and NaAsO2 powders were used as As(V) and As(III) standards. XANES spectra were processed using the Iffefit software package (Athena and Artemis) [29]. 3. Results and discussions 3.1. Characterization of synthetic pyrite and surface-oxidized pyrite Fig. 1a shows the SEM image of pristine pyrite after purification. Elemental analysis from EDS in Fig. 1c indicated peaks of iron (Fe) and sulfur (S). The average ratio of Fe to S) was 1:2.09 based on six randomly selected locations, which was close to the stoichiometry in pyrite (FeS2). The surface area was 12.6 m2/g. Since the density of pyrite is 5.01 g/cm3 therefore the primary particle size (assuming spheres) was 0.048 lm. The average particle size from SEM was 1.0 lm, suggesting that pristine pyrite was an aggregate of smaller particles. Surface-oxidized pyrite has a core–shell structure [24]. Fig. 1b shows the SEM image of surface-oxidized pyrite and indicates morphological changes occurred due to aeration and surface oxidation. Based on similar analyses as reported in the prior paragraph, the average ratio (Fe/S) was 1:1.40, the surface area was 7.7 m2/g, primary particle diameter was 0.078 lm, and SEM particle size was 1.5 lm. The elemental analysis in Fig. 1d showed a small oxygen peak. These results indicated formation of an iron oxide phase on the pyrite surface but pyrite was the only XRD-identified mineral phase for either pristine or surface-oxidized samples (Fig. 2). 3.2. Reactions of As(III) on pristine and on surface-oxidized pyrite As(III) was adsorbed on pristine pyrite from pH 2 to 12 as shown in Fig. 3. Adsorption decreased for pH > 12 perhaps related to pH approaching the pKa3 of H3AsO3. As(III) adsorbed over a narrower pH range on surface-oxidized pyrite, with maximum adsorption between pH 3.5 and 10. The narrower adsorption envelope on surface-oxidized pyrite was similar to that reported for As(III) on goethite or ferrihydrite [2,3]. Fig. 3 also shows reported results for As(III) adsorption on commercial pyrite that had been washed in a 0.01 M sulfide solution, used in a 10/90 H2/N2 environment, and using a similar adsorbate to adsorbent ratio as in the other curves in Fig. 3 [14]. Those authors found evidence of surface oxidation of sulfur and iron, possibly due to material preparation prior to analysis by X-ray photoelectron spectroscopy (XPS). They noted
Fig. 3. As(III) removal using pristine pyrite, surface-oxidized pyrite, and reduced commercial pyrite. Conditions: 45 lM As, room temperature, and 48 h reaction time.
that ‘‘limited surface oxidation may influence As(III) adsorption, potentially altering the extent of sorption or the redox state of sorbed As.’’ The similarity of the As(III) adsorption envelope on the sulfide-treated commercial pyrite [14] and our results on surface-oxidized pyrite suggest that the pyrite surfaces in both experiments were oxidized, while the broader adsorption envelope for As(III) on our pristine pyrite suggests that the surface was different, perhaps a true pyrite surface.
3.3. Reactions of As(V) on pristine pyrite and on surface-oxidized pyrite Fig. 4a shows that As(V) was adsorbed on pristine pyrite for acidic pH, and that all of the adsorbed As(V) at pH 3 and 3.5 was reduced to As(III). As(V) adsorption and the extent of reduction to As(III) decreased at higher pH values (6.3, 7.1, 8.3, 10.0, 11.0). None of the As(III) produced by reduction of As(V) was released into the dissolved phase. The apparent incorporation of arsenic into the solid phase during reduction of As(V) with pristine pyrite indicates that arsenic removal in anaerobic environments could be greater than would be predicted based on simple adsorption of As(III). The reaction of As(V) with surface-oxidized pyrite exhibited an entirely different reaction mechanism, as shown in Fig. 4b. There was significant removal of As(V) up to pH 9 using 1 g/L surface-oxidized pyrite but removal of As(V) peaked around pH 6 with 0.1 g/L pyrite. There was no reduction of As(V) to As(III). The reactions of As(V) with surface oxidized pyrite were similar to those reported for adsorption of As(V) onto hydrous ferric oxides [3,30]. XANES spectra are shown in Fig. 5 for As(V) and As(III) standards, for As(V) reaction with pristine pyrite (pH = 3.5), and for As(V) reaction with surface-oxidized pyrite (pH = 6.0). Background removal and normalization were performed for the pyrite samples. A white line (first most intense absorption peak) for As(V) was at 11874.8 eV and for As(III) at 11871 eV. The XANES results confirmed that As(V) was reduced to As(III) with pristine pyrite but not with surface-oxidized pyrite. Some As(V) was found in the pristine pyrite with XANES which the wet chemistry analyses showed that all As(V) adsorbed onto pristine pyrite was reduced to As(III). The measurement of some As(V) with XANES could have been due to oxidation of As(III) by the synchrotron beam as observed by Bostick and Fendorf [14]. The rate of As(V) reduction by pristine pyrite is shown in Fig. 6, as a function of pH and initial As(V) concentration. The arsenic removal rate increased with decreasing pH. The inset in Fig. 6 shows
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Fig. 5. As K-edge XANES spectra of As(V)/As(III) standard samples, and arsenic on the solid phase from As(V) reaction on pristine pyrite (pH = 3.5) and surface oxidized pyrite (pH = 6).
Fig. 4. As(V) removal using (a) pristine pyrite, showing adsorbed and dissolved As(III) and As(V) and (b) surface-oxidized pyrite (0.1 g/L or 1 g/L) which did not reduce As(V) to As(III). Conditions: 45 lM arsenic, room temperature, and 48 h reaction time.
log[As(V)dissolved] versus time for several of the conditions and demonstrates that the reaction was first order with respect to the concentration of dissolved As(V) at pH 3.5 and 3.7. However the reaction stopped short of equilibration at pH 6.7 possibly due to production of an impermeable surface layer of FeOOH that inhibited contact of As(V) with pyrite. 3.4. Stoichiometry and mechanisms for As(V) reaction with pristine pyrite The experimental stoichiometry of the As(V) and pristine pyrite reaction at pH 3.5 is shown in Fig. 7. The incremental Fe(II) concentration was determined after compensating for the dissolution of pyrite in the absence of arsenate. Thiosulfate was not detected and sulfate appeared to be the only oxidized sulfur species. Williamson and Rimstitdt [31] reported that thiosulfate was detected in very small amounts and that polysulfides and sulfite were not detected during oxidation of pyrite by O2. Fe(III) was not produced, which is consistent with previous studies using relatively weak oxidants. Naveau et al. [10] reported
Fig. 6. Total As(V) concentration (dissolved + adsorbed) versus time at different pH values. The inset figure shows log (total As(V)) versus time. Initial pyrite was 1 g/L.
that Cu(II) was reduced by pyrite to Cu(I) and only sulfur was oxidized. Naveau et al. [32] reported that Se(IV) was reduced by pyrite and only sulfur was oxidized. However Demoisson et al. [33] found that Cr(VI) oxidized pyrite to sulfate and Fe(III). Plots of the logarithms of dissolved As(V), Fe(II), and sulfate concentrations versus time in Fig. 7b produced R2 values of 0.99, 0.89, 0.91. The slopes were 0.110 s1 for As(V), 0.031 s1 for sulfate, and 0.014 s1 for Fe(II). The ratio of these slopes was 7/2/0.9, which is close to the theoretical stoichiometric ratio of 7/2/1 for the reaction.
7H2 AsO4 þ FeS2 þ 5Hþ þ H2 O ¼ Fe2þ þ 2SO2 4 þ 7H3 AsO3
DG0
¼ 42:99 kJ=mol
ð1Þ
The reaction rate constant did not change over time at pH 3.5. Williamson and Rimstidt [31] reported a decrease in the rate constant for pyrite oxidation by O2 with time, possibly due to
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Fig. 8. The rate of pristine pyrite oxidation by As(V) (solid lines) is compared with oxidation by O2 and dissolved Fe(III) (data from Williamson and Rimstidt [31]). Rates were normalized using the pyrite surface area (mol m2 s1). Conditions: 0.2 atm O2, dissolved Fe(III) in equilibrium with ferrihydrite, 45 lM or 0.133 lM As(V).
4. Conclusions and environmental implications
Fig. 7. (a) Concentrations of dissolved As(V), incremental dissolved Fe(II) (after subtraction of pyrite self-dissolution), and dissolved sulfate versus time during the reaction between As(V) and pristine pyrite at pH 3.5; (b) logarithm of As(V), sulfate, and incremental Fe(II) concentrations plotted versus time. Initial As(V) was 200 lM and initial pyrite 1 g/L.
formation of an oxidized shell. Their initial rate constant was not restored after cleaning with concentrated HNO3 (an oxidizing agent) or EDTA solution (relatively weak chelator at acidic pH). They noted that similar reductions in oxidation rate had been previously observed after weathering or polishing of pyrite. As(V) was a weak oxidant for pyrite compared to O2 or Fe(III). This is illustrated in Fig. 8, showing reaction rates (normalized by pyrite mass) for oxidation by 45 lM As(V), 10 lg/L As(V) (the USEPA maximum contaminant level for arsenic in potable water), 0.2 atm O2, or dissolved Fe(III). The O2 and Fe(III) data are from Williamson and Rimstidt [31]. They suggested that the observed pH effects were due to changes in adsorbed O2 and Fe(II) and on solubility of Fe(III). Similarly the decreasing rate of pyrite oxidation by As(V) with increasing pH could be due to decreased adsorption of As(V) at higher pH, consistent with our observation that only adsorbed As(V) was able to reduce pyrite. The slope of log(r) versus pH for As(V) was 0.5. Fig. 8 illustrates that O2 must be rigorously excluded from anaerobic chambers that are used to study reactions involving pyrite and As(V) or other weak oxidants.
Pristine pyrite was produced, cleaned, and stored in a rigorously anoxic environment (PO2 < 108 atm) and ‘‘surface-oxidized’’ pyrite was produced from pristine pyrite that was exposed to air and subsequently stored and reacted under anoxic conditions. Both materials appeared to be pure pyrite based on XRD analyses, but the two materials exhibited substantially different reactivity towards As(III) and As(V). As(III) adsorbed over a broader pH range on pristine pyrite than on surface-oxidized pyrite. Adsorption of As(III) on surface-oxidized pyrite was similar to reported adsorption on pyrite that had been prepared and stored in less-rigorously anoxic conditions than in our experiments or on iron oxides. As(V) adsorbed over a narrower pH range on pristine pyrite than on surface-oxidized pyrite. Adsorption of As(V) on surfaceoxidized pyrite was similar to previously reported adsorption on iron oxides. All adsorbed As(V) on pristine pyrite was reduced to As(III) at pH < 4 but As(V) was not reduced at all to As(III) when using surface-oxidized pyrite. Half of adsorbed As(V) was reduced to As(III) at pH 6.3 after 2 days. Reduction of As(V) to As(III) with pristine pyrite was confirmed using XANES analyses. As(III) that was produced by reaction of As(V) with pristine pyrite was incorporated into the solid phases and was not released back to solution. This result suggested that reduction of As(V) by pristine pyrite could result in better immobilization than would be predicted based on experiments in which As(III) is adsorbed onto pyrite. Reduction of As(V) with pristine pyrite was first-order in total As(V), the rate of As(V) reduction decreased with increasing pH, and only sulfur was oxidized. The proposed mechanisms of pyrite oxidation by As(V) were similar to previously published mechanisms for the oxidation of pyrite by O2 and the logarithms of the rates of the two reactions were proportional to the free energy of the two reactions. Therefore oxidation of pyrite by As(V) is unlikely to be observed in the presence of very low partial pressures of O2 that could be present in glove bag environments without using extra precautions for removal of trace concentrations of O2. The results can be used to predict reactivity of arsenic with pyrite in either consistently anoxic environments or in intermittently
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anoxic environments. The results also are pertinent for laboratory studies in which it may be necessary to produce and maintain pyrite in rigorously anoxic environments in order to avoid experimental artifacts. Acknowledgment This research was partially supported by the Center for Environmental Kinetics Analysis at Penn State University (NSF Grant No. Che-0431328). References [1] [2] [3] [4] [5] [6] [7] [8] [9] [10]
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