Environ. Sci. Technol. 2000, 34, 4714-4720
Kinetics of Arsenate Reduction by Dissolved Sulfide ELIZABETH A. ROCHETTE,† BENJAMIN C. BOSTICK,‡ GUANGCHAO LI,‡ AND S C O T T F E N D O R F * ,‡ Department of Natural Resources, University of New Hampshire, Durham, New Hampshire 03824, and Department of Geologic and Environmental Sciences, Stanford University, Stanford, California 94305-2115
Arsenic toxicity and mobility in soil and aquatic environments depends on its speciation, with reducing environments generally leading to more hazardous conditions with respect to this element. Aqueous sulfide (H2S or HS-) is a strong reductant and often occurs at appreciable concentrations in reduced systems. Consequently, it may play an integral part in arsenic redox chemistry. Therefore, reactions between arsenic and sulfide may strongly influence water quality in arsenic-contaminated systems. To evaluate this possibility, we investigated the kinetics and reaction pathways of arsenate with sulfide. Arsenate reduction by hydrogen sulfide is rapid and conforms to a second-order kinetic model, having a rate constant, k ) 3.2 × 102 M-1 h-1, that is more than 300 times greater at pH 4 than at pH 7. However, arsenite is not the direct reaction product. Rather, arsenic-sulfide complexes develop, including the formation of a trimeric species (HxAs3S6x-3), that persist in solution for several days, ultimately dissociating and leading to the production of dissolved arsenite. The precipitation of orpiment is dominant only at high (20:1) S:As ratios, considering the reaction conditions used in this study (133 µM As, pH 4). Hence, models of arsenic behavior in the environment should consider abiotic reduction of arsenate by sulfide, at least under moderately acidic conditions, and the possibility of dissolved arsenic-sulfide complexes.
Introduction Agricultural soils as well as stream and lake sediments may contain arsenic from historic arsenical pesticide applications, mining practices, and industrial activities. Various pathways have been proposed for arsenate (HxAsO4x-3) reduction, such as reduction and alkylation (1) and through respiration coupled with the oxidation of organic molecules or hydrogen (2-5). Microbial activity is also thought to be indirectly or directly responsible for arsenate reduction to arsenite (HxAsO3x-3) in flooded soils and sediments (6-8), with microbially stimulated reduction demonstrated in these natural materials (9-11). Nonetheless, abiotic arsenate reduction processes should not be dismissed. Sulfide in the form of H2S or HS- is frequently present at appreciable concentrations in reducing environments and is a strong reductant (see, for example, ref 12). On the basis of standard* Corresponding author phone: (650)723-5238; fax: (650)725-2199; e-mail:
[email protected]. † University of New Hampshire. ‡ Stanford University. 4714
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state redox conditions, hydrogen sulfide (H2S) should be capable of reducing arsenate (H2AsO4- or HAsO42-) to arsenite (H3AsO3):
H2AsO4- + H2S + H+ S H3AsO3 + 1/8S8 + H2O (1) On the basis of the free energies of formation provided by Wagman et al. (13), ∆G°rxn ) -100.5 kJ/mol. Newman et al. (4) noted that the kinetics of arsenate reduction by sulfide at pH 6.8 are slow (second-order rate constant, kpH6.8 of 1.04 M-1 h-1). However, other pH values were not considered, and the conversion of sulfide from HSto H2S below pH 7 could vastly change the interactions of sulfide with arsenate. For instance, within 33 h nearly complete conversion of arsenate to arsenite occurs at pH 4, while less than 10% occurs at pH 7 with initial sulfide:arsenate molar ratios of roughly 104:1 (14). Furthermore, arsenate reduction to arsenite has been observed in fjord waters at a depth where H2S is dominant over O2 (15). Once reduced, As(III) may also form soluble and insoluble arsenic-sulfide complexes. Several soluble arsenic-sulfide complexes, such as di- and trithioarsenite monomers as well as dimeric and trimeric arsenic-sulfur complexes have been described or suggested during orpiment (As2S3) dissolution (16-24). Because of the environmental significance of arsenic redox reactions and the potential for it reacting with sulfide in acidic systems, we conducted this study to determine the kinetics and reaction pathways of arsenate with dissolved sulfide.
Experimental Section Chemicals. Sodium sulfide (Na2S‚9H2O, 99%), sodium borohydride (NaBH4, 98-99%), ferric ammonium sulfate [FeNH4(SO4)2‚12H2O, 99%], and N,N-dimethyl-p-phenylenediamine sulfate [(C2H5)2NC6H4NH2‚H2SO4, 97%] were obtained from Sigma-Aldrich Chemical Co. (St. Louis, MO). Sulfidic solutions were prepared using degassed doubly deionized water in an O2-free environment (i.e., within a N2-purged glovebox). Sodium arsenate (NaH2AsO4‚7H2O), sodium arsenite (NaAsO2), glacial acetic acid (CH3COOH), sodium acetate (CH3COONa‚ 3H2O), Tris [(HOCH2)3CNH2], sodium hydroxide (NaOH), and concentrated hydrochloric acid (HCl) were reagent-grade chemicals. Acetate buffers (0.1 M) were used to maintain a constant pH of 4.0 or 5.0 and were prepared by combining acetic acid and sodium acetate to obtain the desired pH. Tris buffer was used to maintain a pH of 7 and was prepared by adjusting the pH of 0.1 M Tris with the addition of concentrated hydrochloric acid. Kinetic Experiments. We first explored the effect of sulfide and arsenate concentrations on the ensuing redox reaction. Sodium arsenate was added to 0.1 M acetate buffer to obtain a concentration of 133 µM arsenate. Buffered (pH 4, 5, or 7) arsenate solutions in glass bottles were placed in a nitrogenpurged glovebox along with a concentrated solution of sodium sulfide; dissolved sulfide was added to the buffered arsenate solutions in the glovebox to obtain initial sulfide concentrations of 13.3, 26.6, 133, and 266 µM sulfide, yielding sulfide:arsenate ratios of 1:10, 1:5, 1:1, and 2:1, respectively. Solution pH and Eh were determined, and then the bottles were capped and placed on a shaker at 150 rpm for ca. 96 h, after which time the samples were filtered through 0.22µm pore filters. These experiments were performed in duplicate at 24 ( 2 °C. To obtain greater detail on the reduction of arsenate by sulfide, we investigated the reaction at pH 4 and an initial sulfide:arsenate ratio of 2:1 (266:133 µM) while sampling more 10.1021/es000963y CCC: $19.00
2000 American Chemical Society Published on Web 09/30/2000
frequently in triplicate. Separate experiments, differing only with respect to the sample times and speciation methods, were performed for arsenate and sulfide analysis. Subsamples of the buffered arsenate/sulfide solutions were filtered and analyzed at times from approximately 1 to 73 h; a simple fourth-order polynomial (r 2 > 0.99) was used to extrapolate the sampling times and quantities of arsenate for comparison to those of sulfide. Two experiments were performed with higher S:As ratios. In one case the S:As ratio was 100:1 (initial SII- 13.3 mM:AsV 133 µM) at pH 7 (Tris-HCl); samples were withdrawn after 7 days. The other experiment was performed at pH 4 (0.1 M acetate buffer) and S:As 20:1 (initial SII- 2660 µM:AsV 133 µM). Samples were withdrawn at 28.5 and 32 h. Arsenic Speciation and Analysis. Arsenic speciation was performed with an ion chromatographic (IC)-hydride gas generation (HGG)-inductively coupled plasma (ICP) spectrometric method, based on the IC-HGG-atomic absorption spectrometric (AAS) method of Manning and Martens (25). The IC-HGG-ICP apparatus consisted of a Dionex DQP-1 pump, a 6-port Rheodyne injection valve with a 50-µL sample loop, Dionex AG-11 and AS-11 guard and analytical columns, a Varian gas-liquid separator, and a Thermo Jarrell-Ash IRIS ICP spectrometer. Flow rates of the IC, sample, and HGGICP were equal at 2 mL min-1. Sample, 6 N HCl, and sodium borohydride solution (0.16 M NaBH4/0.12 M NaOH) were combined after the ICP pump prior to a mixing coil and were followed by a gas-liquid separator. Arsine gas was swept from the separator into the ICP torch by argon carrier gas. Arsenite and arsenate retention times were monitored in time-scan mode; peak areas were determined and compared to those of calibration standards for quantification. The detection level for arsenite (As III) was 50 h), arsenite begins to form appreciably (Figure 2), and a more extensive reaction series that leads to the formation of this species is required. Identifying Intermediate Species. During the initial ligand displacement, arsenate concentrations should be depleted with the concurrent buildup of H2AsO3S- and H2AsO2S2-. Using ion chromatography with conductivity detection, we note that an additional peak at a retention time of 15.2 min is present. This retention time is greater than that of sulfite, sulfate, thiosulfate, thiocyanate, and arsenate. The concentration of this anion is greatest at approximately 20 h and
FIGURE 2. Dissolved arsenic species and total dissolved arsenic as a function of reaction time at pH 4. The initial arsenate and sulfide concentrations were 133 and 266 µM, respectively. The concentration of As-S complex(es) was(were) calculated by subtracting arsenate and arsenite concentrations from the total dissolved As concentration.
FIGURE 4. Raman spectra of (a) arsenate-sulfide solutions reacted at pH 4 for 50 h and (b) the resulting solids after 330 h of reaction. Spectra of the model compounds orpiment (As2S3) and realgar (AsS) are shown for comparison.
FIGURE 3. Peak area of the 15.2-min eluent species (IC with conductivity detection) compared against the concentration of As bound in the As-S complex(es) estimated using the data shown in Figure 2. correlates with the difference between total dissolved arsenic and measured arsenate and arsenite (∆As) during this early course of reaction (Figure 3). After 20 h, the peak area of this anion decreases, and a disparity between ∆As and the peak area occurs. On the basis of IC-HGG-ICP, the anion contained arsenic, but unfortunately, it could not be quantified without a suitable standard. We therefore speculate that the 15.2-min retention peak within the first 20 h of reaction represents a convolution of the thioarsenate (H2AsO3S-) and dithioarsenate (H2AsO2S2-) complexes (although we cannot rule out thiolated arsenite species). The difference between ∆As and the chromatographic peak area beyond 20 h leads us to speculate that one or more additional arsenic-sulfur complexes are formed prior to the formation of arsenite. Accordingly, we used ∆As concentrations and ∆S concentrations along with Raman spectroscopy of a solution sample at t ≈ 50 h to identify dissolved complexes at intermediate time periods (t ≈ 50 h). At the 50-h reaction mark, the solution concentration of sulfide decreased from an initial value of 266 to 20 µM with very little S present as sulfoxy species. Therefore, the As-S complex(es) coupled with polysulfides and/or elemental sulfur must account for the majority of dissolved S. The solution S:As ratio was approximately 2:1, and since the As-S complex accounts for the majority of dissolved As at 50 h (and possibly S given the levels of sulfide, elemental sulfur,
and polysulfides measured at this time), it appears that an approximate S:As ratio of the complex is 2:1. Considering that elemental sulfur or polysulfides may also account for (small) losses of dissolved sulfide, only S:As ratios lower than 2:1 are possible. This is consistent with the S:As ratio of a trimeric (HxAs3S6x-3) complex, though not proof of a trimer. Mirnova et al. (18) proposed a dimer (HAs2S4-) that would have a similar molar S:As ratio of 2:1 However, on the basis of thermodynamic considerations, dimeric (and tetrameric) arsenic-sulfide complexes were rejected by Helz (24) in favor of monomers in (highly) undersaturated solution and trimers in (near) saturated solutions with respect to orpiment. The mixed dithioarsenite (H2AsOS2-) or dithioarsenate (H2AsO2S2-) complexes would also have this stoichiometry; we dismiss the dithioarsenate complex on the basis of the decreased chromatographic peak (Figure 3) assigned to this species (along with thioarsenate). We also discount the dithioarsenite species since, based on thermodynamic arguments, it should lead to the formation of HxAs3S6x-3 in solution near saturation with respect to orpiment (24), which is the case for this system (vida infra). The thioarsenite monomer [AsS(SH)2-] proposed by Helz et al. (24) is a possible candidate but would have a molar S:As ratio of 3:1. To further aid in identifying the As-S complexes, we used Raman spectroscopy with the spectral assignments of Helz et al. (24). Raman spectra of our samples at t ) 50 h (Figure 4a) consist of a strong band at 456 cm-1, ruling out monomeric As-S complexes observed to have bands between 325 and 412 cm-1 and which are dominantly below 390 cm-1 (24). [Arsenic solutions prepared at pH 11 do however provide bands consistent with the monomer described by Helz et al. (24).] The higher shifts in our spectrum obtained at pH 4 are consistent with theoretically predicted bands in the 350380 cm-1 range (24) for the trimeric As-S species HxAs3S6x-3. VOL. 34, NO. 22, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 5. Schematic illustration of the postulated reaction mechanism for the reduction of arsenate by sulfide under acidic conditions (pH 4) and S:As ratio of 2:1. Our observations suggest the development (oligomerization) of H2As3S6- upon reduction of arsenate by excess sulfide, leading us to include the following reactions in the sequence that may be operational beyond 20 h (t > 20 h)sas summarized in Figure 5. Following the formation of dithioarsenate (reaction 4), we propose an electron transfer reaction resulting in thioarsenite (reaction 5a); the S(0) group would subsequently dissociate (reaction 5b) providing the potential for polymerization and association with H2S to form polysulfides (addressed below in reaction 9) or precipitation as elemental sulfur (S8,solid):
H2AsO2S2- + H+ S H3AsO2S‚‚‚S
(5a)
H3AsO2S‚‚‚S S H3AsO2S + 1/8S8
(5b)
The production of thioarsenite is a pivotal point in the reaction series; it can either undergo ligand exchange with hydrogen sulfide to form dithioarsenite or it can hydrolyze and produce arsenite. Given (i) that we have reasonable evidence for a As-S trimer, (ii) that arsenite is not produced appreciably until late in the reaction sequence (t > 100 h), (iii) and that it (arsenite) does not form appreciably until dissolved sulfide is less than 20 µM, direct formation of arsenite from thioarsenite is rejected. Rather, we believe that initially thioarsenite complexes with H2S to form dithioarsenite:
H3AsO2S + H2S S H2AsOS2- + H+ + H2O
(6)
Oligermization of dithioarsenite is favorable under these reaction conditions (see ref 24) and thus results in the trimeric 4718
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complex (reaction 7; log K ) 18.2):
3H2AsOS2- + 2H+ S H2As3S6- + 3H2O
(7)
However, we have yet to account for the production of arsenite that is observed to increase extensively beyond 50 h and become the dominant aqueous As species beyond 140 h (Figure 2). Additionally, if only reactions 3-7 are considered, the S:As reactant ratio (3:1) would defy mass balance. We turn to the thermodynamic data of Helz et al. (24) in order to try and predict expected final concentrations given sufficient reaction time. Obviously the system is not at equilibrium within the time frame of our experiments, and thus such an analysis only serves to describe a driving force and not reaction rates. The balance between H3AsO30 and H2AsOS2- is described by reaction 8a (log K ) -4.1), and the balance between H3AsO30 and H2As3S6- is described by reaction 8b (log K ) -30.5):
H2AsOS2- + H+ + 2H2O S H3AsO30(aq) + 2H2S(aq) (8a) H2As3S6 - + H+ + 9H2O S 3H3AsO30(aq) + 6H2S(aq) (8b) At pH 4, formation of arsenite is thermodynamically favored over dithoarsenite (reaction 8a) at (H2S) < 87 µMsa condition satisfied only at early time points in the reaction (t < 15 h). Because of the reaction stoichiometry, predicted equilibrium concentrations of the trimeric species are dependent on the absolute activities of H3AsO30 and H2As3S6- along with H2S. Unfortunately, this negates making a prediction of their
relative concentrations at equilibrium, or driving force, based solely on pH and H2S concentrations; it is unwarranted to choose either H3AsO30 or H2As3S6- concentrations a priori. It is obvious, however, from reaction 8b that changes in sulfide concentration will have a dramatic effect on the reaction path, with diminished sulfide favoring arsenite production. On the basis of the rate at which arsenite is produced (Figure 2), we therefore predict that arsenate reacts with dissolved sulfide dominantly to form the trimeric thioarsenite complex (reactions 3-7) until [H2S] is less than approximately 20 µM (occurring at t ) 50 h), at which time arsenite production becomes the dominant pathway. As depicted in Figure 5, we postulate that under such conditions the trimer would deoligermize forming dithioarsenite, which would then undergo two ligand exchange reactions with water to form arsenite. Finally, both ion chromatography and sulfide analyses indicate that at approximately 50 h very little sulfide (
