Environ. Sci. Technol. 2002, 36, 943-948
Antimony Speciation in Alkaline Sulfide Solutions: Role of Zerovalent Sulfur GEORGE R. HELZ,* MELISSA S. VALERIO, AND NATHAN E. CAPPS Chemistry and Biochemistry Department and Water Resources Research Center, University of Maryland, College Park, Maryland 20742
Antimony is subject to a lower drinking water standard than arsenic, its notorious group 15 cohort in the periodic table. Both elements often co-vary in nature and are fairly soluble under reducing, alkaline conditions. Of the two, much less is known about the environmental chemistry of Sb. Measurements of Sb solubility in sulfidic solutions equilibrated with stibnite (Sb2S3) and orthorhombic sulfur reveal the existence of two new complexes that may control Sb behavior in many reducing environments. Formation reactions and stability constants (23 ( 2 °C) are HS- + S(s) + Sb2S3(s) a HSb2S5-, log K ) -1.47 ( 0.17; and HS+ 2S(s) + Sb2S3(s) a Sb2S62- + H+, log K ) -9.55 ( 0.07. The first complex is a mixed-valence SbIII,V complex; the second is an SbV complex. Their stability in sulfidic solutions may explain previously puzzling evidence of SbV in natural anoxic environments. Owing to these complexes, zerovalent S can enhance stibnite solubility up to 3 orders of magnitude. In neutral-to-alkaline, reducing environments, less than 7 µM HS- will transform O-coordinated, electrically neutral Sb(OH)3o to predominantly anionic S-coordinated complexes. This transformation could diminish the adsorption of Sb to negatively charged mineral surfaces, lowering retardation factors in anoxic aquifers.
Introduction Here, we examine incongruities in current knowledge of the speciation and solubility of antimony in sulfidic environments. Cases of human arsenic poisoning produced by consumption of groundwaters extracted from anaerobic aquifers are infamous (1-7). Drinking water standards are lower for Sb than for As (50 nM Sb vs 130 nM As in the United States). Geochemical similarities between As and Sb suggest that more needs to be known about the latter’s occurrence and behavior in anaerobic (including sulfidic) groundwaters. A decade ago, a consensus appeared to develop concerning the speciation of Sb in sulfidic waters. Within a few months of one another, three independent groups announced more or less congruent conclusions. Spycher and Reed (8) critically reviewed the existing, somewhat discordant literature and arrived at the view that the principal aqueous complexes in equilibrium with stibnite (Sb2S3) were the dimers H2Sb2S4o, HSb2S4-, and Sb2S42-. Krupp (9) published the results of an entirely new stibnite solubility study which reached the same * Corresponding author phone: (301) 405-1797; fax: (301) 3149121; e-mail:
[email protected]. Current address: 3101 Chemistry Bldg., University of Maryland, College Park, MD 20742. 10.1021/es011227c CCC: $22.00 Published on Web 01/26/2002
2002 American Chemical Society
conclusion. Wood (10) presented a Raman spectroscopic investigation of high pH, ∼1 M HS- solutions, concluding that the evidence was best interpreted in terms of two species, SbS33- and Sb2S42-. Wood’s evidence could be construed to support the conclusions of the other workers. In equilibrium with stibnite, SbS33- would decline in concentration with falling OH- and HS- more rapidly than Sb2S42-, making the latter anion likely to predominate in the solutions considered by the other workers. A hint that Sb speciation in sulfidic waters was actually more complicated was introduced by Tossell (11), who suggested on the basis of theoretical calculations that H2Sb2S4o was more consistent than Sb2S42- with Wood’s Raman data. Yet, solubility evidence excludes H2Sb2S4o as a stable species in Wood’s high pH solutions. Significantly, Tossell noted that 3-coordinate SbIII and 4-coordinate SbV have similar Sb-S stretching frequencies. Now, two new X-ray spectrographic studies both suggest that SbS43- is important in concentrated sulfide solutions (12, 13). The implication is that sulfidic solutions, which form only in reducing environments, need not mobilize Sb only in the lowest of its two common oxidation states (III and V). Indeed, both aqueous SbV and SbIII have been reported in sulfidic marine environments (14-17), but SbV has not been satisfactorily explained. To shed new light on Sb speciation in sulfidic solutions, we have measured the solubility of the thermodynamically stable solid assemblage, Sb2S3 + S (minerals stibnite and orthorhombic sulfur). We compare our results to those of previous workers who investigated the solubility of Sb2S3 alone. A specific goal is to establish, for the first time, stability constants for any SbV sulfide species. The role of solid S in these experiments is to fix the activity of zerovalent sulfur, aS, and to provide oxidizing capacity. Sulfidic solutions equilibrated with S contain mixtures of polysulfide ions (HxSnx-2, x ) 0-2, n g 2; see ref 18). Polysulfides can function as oxidizing agents, as shown by the following conceptual reaction
SbIII + Sn2- + H+ a SbV + HS- + Sn-12-
(1)
Elemental S is often disregarded in natural waters. Ordinarily, it is a relatively insoluble, hydrophobic substance. Yet, when transformed by HS- into polysulfides, it profoundly affects chemical behavior of diverse contaminants that could cause concern in drinking water (18-24).
Methods Orthorhombic sulfur (Aldrich, 99.98%; 0.1 g) and stibnite (Sb2S3; Alfa Aesar, Aldrich, 99.999%; 0.3 g) powders were weighed into 20-mL glass ampules. Stock solutions of NaHS were prepared by bubbling high-purity H2S through solutions of deoxygenated NaOH for approximately 45 min. Inside a N2-filled glovebox, appropriate amounts of deoxygenated NaOH and NaHS solutions were transferred volumetrically into the ampules. The ampules were removed from the glovebox after being capped with rubber balloons to prevent air contact. They were fusion-sealed and then tumbled on a rotator for an equilibration period of 2-6 weeks. Preliminary experiments to determine the time needed for equilibration indicated that solubility peaked in about 6 days and then declined by about 5% during the following ∼10 days. The slight decline is attributed to slow annealing of the solids. Equilibration occurred at room temperature (23 ( 2 °C). Ampules were opened inside the glovebox and their solutions syringe-filtered (Gelman Acrodisc, 0.45 µm) into VOL. 36, NO. 5, 2002 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Experimental Dataa,b pH
I (M)
∑Sb (M)
∑S-II (M)
aS
HS(M)
“SbS43-” (M)
7.57 7.62 7.64 7.82 7.91 8.01 8.16 8.16 8.59 8.68 8.70 8.77 8.92 9.02 9.09 9.41 9.64 9.71 9.73 9.80 9.98 10.11 8.16 8.38 9.00 8.85 9.20 9.48 10.13
0.09 6 0.097 0.095 0.025 0.007 8 0.008 5 0.010 2 0.089 6 0.012 3 0.002 43 0.002 71 0.003 37 0.000 13 0.004 68 0.018 6 0.000 09 0.001 11 0.002 3 0.006 5 0.027 1 0.014 8 0.015 3 0.096 2 0.122 0.069 4 0.084 2 0.096 0 0.080 6 0.090 3
0.010 0 0.010 5 0.013 5 0.002 17 0.001 01 0.001 41 0.001 18 0.001 07 0.001 73 0.000 62 0.000 45 0.000 52 0.000 40 0.000 75 0.002 29 0.000 13 0.000 42 0.001 20 0.001 55 0.004 04 0.003 29 0.002 56 0.007 32 0.004 32 0.002 57 0.002 67 0.001 68 0.001 72 0.001 01
0.114 0.113 0.113 0.027 1 0.008 63 0.009 37 0.010 1 0.008 90 0.010 5 0.002 43 0.002 37 0.002 80 0.000 63 0.003 54 0.012 4 0.000 23 0.001 09 0.002 80 0.004 95 0.016 5 0.010 6 0.009 55 0.104 0.127 0.070 8 0.086 7 0.097 8 0.082 2 0.091 2
1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 1 0.512* 0.374* 0.248* 0.220* 0.102* 0.079* negligible
0.074 1 0.073 6 0.069 0 0.018 9 0.005 58 0.005 55 0.006 51 0.005 74 0.005 11 0.000 86 0.001 05 0.001 19 small† 0.001 11 0.003 56 small† 0.000 062† 0.000 083† 0.000 34 0.001 12 0.000 41 0.000 34 0.083 5 0.113 0.064 3 0.079 9 0.093 8 0.078 5 0.089 1
0.013 7 0.015 8 0.012 6 0.003 24 0.001 17 0.001 65 0.001 41 0.001 42 0.001 85 0.000 61 0.000 51 0.000 69 0.000 15 0.000 94 0.002 60 0.000 02 0.000 23 0.000 69 0.001 27 0.004 52 0.002 62 0.002 94 0.009 57 0.007 10 0.003 55 0.003 40 0.002 78 0.004 61 0.001 39
a The HS- values marked (†) are estimated to have >15% analytical uncertainty; the two “small” entries are samples for which eq 5 yielded negative HS- values. b aS values denoted (*), from eqs 5 and 6 in text; in the order given in table, ∑So for these samples is 0.005 92, 0.004 68, 0.002 96, 0.001 87, 0.000 94, and 0.000 312 M.
15-mL centrifuge tubes with plastic screw caps. Final pH measurements were obtained using an Orion 420A meter equipped with an Orion 8103 semi-micro combination electrode calibrated at pH 7 and 10 with commercial buffers. Sample aliquots were pipetted into screw-cap glass jars containing 50-60 mL of deoxygenated NaOH, which produced pH ∼ 13. After being removed from the glovebox, the solutions were potentiometrically titrated for total sulfide using Hg2+ and an Orion silver/sulfide ion selective electrode combined with an Orion Ag/AgCl double junction reference electrode. UVvisible absorption data on filtered samples in quartz cuvettes were obtained with a Hewlett-Packard 8452A diode array spectrophotometer. When dilution was necessary to overcome saturated optical absorption, 1 mM NaOH was used as the diluent to prevent Sb precipitation. Total Sb concentrations were determined by flame (air-acetylene) atomic absorption spectroscopy using a PerkinElmer 5000 spectrophotometer at a wavelength of 217.6 nm. Calibration curves were produced over the range 1-50 ppm using a commercial antimony standard (EM Science). Sample aliquots were diluted with deionized water to obtain absorbance values within the linear range of the standard plots. To prevent chemical interferences from sulfate, phosphate, and iron (25), 1000 mg/L Al was introduced to all standards and samples following the addition of BrCl (0.9 M). The use of BrCl was required to avoid the precipitation of aluminum sulfide. X-ray powder diffraction patterns were acquired over 1 h between 3° and 90° 2θ using a step width of 0.02°.
Results Data. Table 1 presents the experimental data. The values in the last three columns require comment. Initially, we made an unsuccessful attempt to determine HS- by spectropho944
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FIGURE 1. Optical absorption spectra of 3 samples from Table 1. Measurements made in 0.1 cm path length cells: pH 9.98 sample diluted 1:10, others diluted 1:100. tometry. Figure 1 shows spectra obtained in a 0.1-cm cell using solutions that had been diluted 10-100-fold in order to avoid saturated optical absorption. In the solution that contained no So (other than traces that might have formed by oxidation), a clean HS- peak is seen at 230 nm. However, in the two solutions that were saturated with So during the equilibration period, the HS- peak at 230 nm is buried by other absorptions. Some of this interference might arise from small amounts of colloidal S precipitated during dilution. Additionally, there are probably contributions from thioantimony and from residual polysulfides in this region. Because of these interferences, we chose to obtain the HSconcentrations (Table 1) from the titration (∑S-II) data in the following way. At high pH, Hg2+ titrant reacts in a 1:1 ratio with polysulfides and free sulfide; we find that it also reacts in a 4:1 ratio with an SbS43- standard, suggesting that all S2bonded to Sb also forms HgS during the titration. This information is summarized in the following relation between the molar concentration of titrated sulfide, ∑S-II (column 4, Table 1), and molar concentrations of individual species, where m represents the average number of reactive S2- ions coordinated to each Sb atom.
∑S
-II
) HS- + H2S +
∑S
2-
n
+
∑HS
-
n
+m
∑Sb
(2)
Fully protonated polysulfides (H2Sn) are neglected because they are insignificant in the alkaline solutions under consideration. Using polysulfide formation constants for the following general reactions, eq 2 can be transformed to eq 5, permitting HS- to be calculated.
(n - 1)S + HS- a Sn2- + H+ KRn ) aSn2-aH+/aHS-aS(n-1) (3) (n - 1)S + HS- a HSn-
∑S
-
HS ) 1+
γ′10-pH KH2S
+
Kβn ) aHSn-/aHS-aS(n-1) (4) -II
-m
∑Sb
γ′KRnaSn-1
∑ γ′′10
-pH
+
(5)
∑K
n-1
βnaS
Activity coefficients for singly and doubly charged ions (γ′ and γ′′) were calculated from the Davies equation. Summations in the denominator of eq 5 are for polysulfides having
FIGURE 2. X-ray diffraction pattern of solids recovered from the pH 7.64 sample in Table 1. Cu Kr radiation. n ) 2-5. Values for KRn and Kβn are tabulated in ref 23. KH2S (10-7.01) is the first ionization constant of H2S. To solve eq 5, experimental values must be supplied for ∑S-II (from titration), ∑Sb (from atomic absorption spectroscopy), H+ (from pH measurement), as well as for aS and m. In most of our samples, sufficient elemental S was added to ensure that the solutions remained saturated; for these, aS was defined to be unity (column 5, Table 1). Figure 2 shows an X-ray diffraction pattern obtained from the solids recovered from an ampule after a solubility experiment. Sharp peaks corresponding to stibnite and orthorhombic S demonstrate that both solids persisted through the experiment in well-crystallized forms. In several experiments, gathered at the bottom of Table 1, we deliberately charged the ampules with insufficient elemental S to maintain saturation. For these cases, aS was calculated by assuming that all of the elemental S dissolved, giving a known ∑So (footnote, Table 1), which is related to polysulfide concentrations by the following equation:
∑S ) ∑(n - 1)S + ∑(n - 1)HS ) (n - 1)γ′K a (HS ) ∑ + ∑(n - 1)K γ′′10 o
-
(
-
2-
n
n
n-1
Rn S
-pH
n-1
βnaS
)
(6)
Because ∑S-II, ∑So, ∑Sb, and pH are known from the experiment, eqs 5 and 6 can be combined and solved numerically for both HS- and aS. Equations 5 and 6 operationally define aS, which is an important property of sulfidic waters, even though not directly measurable. Influenced by previous EXAFS results (12, 13), we began this investigation with the belief that the principal dissolved Sb species in S-saturated solutions would prove to be SbS43-. Tests with solutions prepared from commercial Na3SbS4‚ 9H2O (Pfalz & Bauer) showed that all four sulfurs in this anion are titrated by Hg2+, implying that m in eq 5 should be 4. However, this produced many negative HS- concentrations from eq 5. Furthermore, computing the charge imbalance in solution revealed a significant excess of negative charge if all dissolved Sb was assumed to be present as SbS43-. Test computations using eq 5 showed that HS- was positive in all but two cases if m ) 2, the value used to calculate HS- in Table 1. This choice is further justified later. Note that m ) 2 is consistent with known SbIII complexes, HxSb2S4x-2, but excludes SbS43- as a major component in solution. Two drawbacks to quantifying HS- as previously described are that (a) the denominator of eq 5 is subject to systematic error related to possible inaccuracies in the polysulfide
stability constants and (b) if the difference that is computed in the numerator of eq 5 is small, then the numerator’s value is vulnerable to imprecision in the ∑S-II and ∑Sb measurements. To explore the magnitude of the systematic error, we repeated all of the calculations discussed in this paper with an independent set of constants for reactions 3 and 4. The HS- concentrations in Table 1 were calculated using constants selected by Shea (23, 26). Shea tested various constants’ abilities to predict optical absorption (370 nm) for a series of polysulfide solutions that he prepared and characterized carefully. The alternate set of constants that we used were from ref 27 and were obtained by an interpolation procedure from published data. The alternate set of constants produced HS- concentrations systematically ∼2-fold lower than those in Table 1. A corresponding ∼2-fold increase arose in the values of the equilibrium constants derived in the following paragraphs. A key point, however, was that the conclusions regarding the stoichiometry of the dissolved Sb complexes were unchanged. Systematic error can never be fully quantified (otherwise corrections would be made). Nevertheless, this comparison of two sets of polysulfide stability constants gives a general sense of the likely magnitude of systematic error from this source. To explore the random error, we computed analytical uncertainty in the numerator of eq 5 for all cases in Table 1 by propagation of the following estimated measurement errors: 2% for ∑S-II and 5% for ∑Sb. With m ) 2, the numerator in eq 5 was precise to better than 10% except in the four cases flagged by daggers (†) in Table 1. In two of these four, the error was apparently large enough that the calculated HS- values were negative. These two cases were considered no further. The remaining two cases, though producing plausible HS- concentrations, nonetheless consistently gave trouble during data analysis, as will be noted in the following paragraphs. The final column in Table 1 shows concentrations of “SbS43-” calculated from optical absorption at 284 nm (Figure 1). Using Na3SbS4‚9H2O standards, we found that SbS43produces a peak at 284 nm with a molar extinction coefficient of 13 960 M-1 cm-1; this value was used to produce the data in column 7 of Table 1. Polysulfides also absorb near 284 nm, although they have lower molar extinction coefficients. Nonetheless, some of the absorbance attributed to “SbS43-” may be due to residual polysulfides. Curiously, there is fairly good correspondence between “SbS43-” and total Sb concentrations determined by atomic absorption spectroscopy (column 3, Table 1). Seemingly in contradiction to the aforementioned evidence but consistent with expectations based on the EXAFS studies, the optical absorption evidence suggests that dissolved Sb in these samples consists of SbS43-. Analysis of Speciation. Deduction of Sb speciation and derivation of stability constants were accomplished by a statistical fitting procedure. Reactions of the following form were postulated, permitting the calculation of provisional concentrations, MSb, for hypothetical complexes.
aHS- + bSο + c/2Sb2S3(s) a H(a-d)SbcS(a+b+3c/2)(-a-d) + dH+, MSb ) K(aHS-)a(aS)b(aH+)-d(γ(-a-d))-1 (7) (γ(-a-d) is the activity coefficient from the Davies equation for an ion having the charge of the hypothetical Sb complex.) Values of K were adjusted by a least-squares process to achieve the best match between logarithms of predicted and observed solubilities. The first step involved fitting only the aS ) 1 data and assuming that only a single Sb species VOL. 36, NO. 5, 2002 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 2. Statistical Description of Fits to the Data for Only Solutions Saturated with Orthorhombic Sulfura reaction
x)0
x)1
x)2
1.5HS- + 0.5Sb2S3(s) a HxSbS3x-3 + (1.5 - x)H+ 1.0HS- + Sb2S3(s) a HxSb2S4x-2 + (1.0 - x)H+ 0.5 HS- + 0.5Sb2S3(s) + H2O a HxSbOS2x-3 + (2.5 - x)H+ 0.5Sb2S3(s) + 2H2O a HxSbO2Sx-3 + (3.5 - x)H+ + 0.5HS-
0.451, 0.975 0.258, 0.992 1.40, 0.760
0.774, 0.927 0.710, 0.939 0.920, 0.897
1.53, 0.715 0.291, 0.989 1.34, 0.783
x)3
0.748, 0.932 0.944, 0.892
a The number in normal font is σ; the number in italic font is R2; it was assumed that only one Sb complex was present; x refers to the number of protons in the complex, as indicated in the reaction.
TABLE 3. Antimony Equilibriaa reaction 1. 2. 3. 4. 5. 6. 7. 8.
HS-
+ S + Sb2S3(s) a HSb2III,VS5HS- + 2S + Sb2S3(s) a Sb2VS62- + H+ HS- + Sb2S3(s) a HSb2IIIS4HS- + Sb2S3(s) a Sb2IIIS42- + H+ 3H2O + 1/2Sb2S3(s) a SbIII(OH)3o + 3/2H+ + 3/2HSS + HSb2IIIS4- a HSb2III,VS5S + HSb2III,VS5- a Sb2VS62- + H+ 3H+ + 4HS- + 2SbIII(OH)3o a HSb2IIIS4- + 6H2O
log K
ref
-1.47 ( 0.17* this work -9.55 ( 0.07* this work -2.89 -12.69
32 32
-27.53
32
1.42 -8.08
derived derived
52.16
derived
aAll corrected to zero ionic strength; values from this work, 23 ( 2 °C; other values, 25 °C. For (*), uncertainties (1σ) are derived from fitting the data; they do not allow for possible systematic errors discussed in the text.
accounted for solubilities. Two statistics, σ and R2, obtained at this stage are shown in Table 2. Smaller values of σ ({[∑(Ycalc - Yobs)2]/[data points - fitting parameters]}2) and larger values of R2 ({1 - [∑(Ycalc - Yobs)2]/[∑Yobs2]}) denote better fits. This first-stage analysis suggests that either Sb2S42or H2SbOS2- (highlighted in boldface in the table) could predominate in solution. Both are characterized by small values of a and d in eq 7. Because the first-stage analysis involved only aS ) 1 data, the true formulas and oxidation states of the Sb complexes are ambiguous. For example, adding one equivalent of So to H2SbIIIOS2- produces H2SbVOS3-. Likewise, adding one or two equivalents of So to SbIII2S42- produces SbIIISbVS52- and SbV2S62-, respectively. At constant aS, these cannot be distinguished because their solubilities vary in the same way with aHS- and aH+. Likewise, the statistics of fitting SbVS43- to the data would be identical to the fit of SbIIIS33-, shown in line 1 of Table 2. In the second stage of the analysis, the data obtained at varying aS were included in the fitting process. Solubility was found to scale approximately as aS2. This result suggests that the dimer SbV2S62- is the best 1-complex descriptor of solubility; H2SbOS2- is less attractive both because it yields slightly worse fitting statistics (Table 2) and because adding two equivalents of S(0) to it would produce an implausible Sb(VII) species (although structures could be postulated that would have a polysulfide ligand associated with Sb(III)). The final stage of the analysis involved graphical examination of the deviations, log(calculated solubility/observed solubility), versus pH and aHS-. This subjective process sometimes reveals trends not detected when goodness-offit is evaluated solely from statistical parameters. A slight trend toward underprediction of solubility at lower pH was noted and was removed by expanding the model to two complexes, HSbIIISbVS5- and SbV2S62-. Table 3 presents stability constants for these complexes. For purposes of later comparison, additional reactions taken from the literature are also tabulated. 946
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FIGURE 3. Plots of experimental data. Left panel shows relationship between log [10-pH(γ′′∑Sb/2)] and log [γ′(HS-)] above pH 9; the 1:1 slope is consistent with predominance of Sb2S62-. The relationship deteriorates at low HS-, probably because of random error in the HS- values, as discussed in the text. Right panel shows the 2:1 relationship between log solubility and log aS, implying that the dominant complex in this pH range incorporates 2 So atoms per molecule. This is also consistent with Sb2S62-. Figure 3 illustrates some of the characteristics of the data set that led to postulating these particular Sb complexes. The left-hand panel shows that an empirical estimate of the product, aH+aSb2S62-, varies as the first power of an empirical estimate of aHS- (i.e., with a 1:1 slope on a log-log plot). In other words, the coefficient, a, in reaction 7 is 1. This information rules out SbS43-, for example, as a major species; SbS43- requires that a ) 3/2, b ) 1, c ) 1, and d ) 3/2. Another point shown by the left-hand panel of Figure 4 is the scatter that develops at low aHS-. For reasons already given, we believe that this arises from random error in the numerator of eq 5 at low HS- and have chosen to ignore it. If one were to assume that the two aberrant points at aHS-< 10-4 were valid measurements, then one might chose to propose additional species whose activities scale as the zeroth or 1/2 power of aHS-. The right-hand panel of Figure 4 shows that the solubility above pH 9 scales as the second power of aS, indicating that 2 So atoms are consumed in producing the predominant complex in this region. This evidence rules out known, large antimony(III)-polysulfide dimers, such as Sb2S152- (28, 29), which would scale as a high power of aS; Sb2S152- also displays a strong optical absorption band at 322 nm (not seen in our solutions).
Discussion An interesting characteristic of the two new Sb complexes identified in this work is that they vary with pH and aHS- in the same fashion as the known SbIII complexes, HSb2S4- and Sb2S42-. Yet equilibrium constants for reactions 1 versus 3 and reactions 2 versus 4 in Table 3 reveal that, at aS ) 1, the higher-valent complexes are 1-3 orders of magnitude more soluble.
FIGURE 4. Proposed structural relationships of thioantimony complexes. Clockwise from left: two So atoms are added successively to nonbonding electron pairs on the predominant SbIII dimer in reducing solutions, producing a mixed-valence, SbIII,V dimer and then a SbV dimer. Figure 4 illustrates our concept of the structures of the new complexes and their relationships to the established SbIII dimer, HSb2S4-. We envision these higher-valent dimers forming when polysulfide ions donate So atoms to nonbonding electron pairs associated with SbIII atoms. The proposed structures suggest explanations for some of the seeming inconsistencies uncovered during our work. Kunkely and Volger (30) have noted that optical absorption spectra of SbIIIS33- and SbVS43- are quite similar (λmax 279 nm, ∈ 8000 M-1 cm-1 vs λmax 285, ∈ 13 700, respectively). Because the new dimers are very similar structurally to these two monomers, it is not surprising that they also display strong absorption in the same wavelength range. This accounts for the fair agreement between ∑Sb and “SbS43-” in Table 1. As previously described, the number of titratable S2- ions per Sb was set to m ) 2 in our data set. If all of the S atoms in HSb2S5- and Sb2S62- were titratable, then m would be larger, ranging from 2.5-3, depending upon pH. However, it is plausible that reaction of these complexes with Hg2+ during titration regenerates the So incorporated during complex formation (i.e., that SbV reverts to SbIII during titration). If true, m ) 2 would be appropriate for both HSb2S5- and Sb2S62-. The small average charge per Sb atom in these complexes, negative 0.5-1.0, also eliminates the apparent excess negative charge noted when it was assumed that SbS43was the principal species. Our work provides no support for the suggestion in ref 13 that Sb(SH)4+ is important in high pH, concentrated HSsolutions. Although our solutions did not match the compositions studied in ref 13, we nevertheless doubt that Sb(SH)4+ exists under those conditions. Pentavalent Sb would be unlikely to form such a weak acid that it could retain protons at high pH. Another X-ray spectrographic study (12), on the other hand, includes several observations consistent with the results of this work. For example, in many cases, observed X-ray absorption edges for dissolved Sb were intermediate between those of SbIII and SbV compounds. This may signal the presence of mixed-valent complexes such as HSb2S5-. It is tempting to identify the group 1B species in that paper, “probably multimeric, ... possibly Sb(V),” with our new
FIGURE 5. Stability fields of dissolved Sb species as a function of log aS and log aHS-. Diagram constructed for assumed conditions: pH 7.5 and Sb(OH)3o 10-9 M. At this pH, Sb2S62- would predominate only in solutions supersaturated with S. At higher pH, both the Sb2S62- and Sb(OH)3o fields expand. The hatched region denotes conditions below 150 m in the Black Sea, based on estimates in ref 24. complex, Sb2S62-. On the other hand, the evidence in that paper for SbS43- over a range of pH and HS- concentrations remains inconsistent with our work. Possibly, SbS43- is stable beyond the pH and aHS- ranges that we covered. Another possibility is that SbS43- is metastable and was produced in the X-ray spectrographic experiments by the dynamic redox conditions in the X-ray beam (31). Despite similarities between (12) and this work, it is not possible to reconcile them in detail, indicating that there is still more to learn about Sb speciation in sulfidic solutions. In previous studies of stibnite solubility, aS has not been constrained. Small amounts of unrecognized sulfide oxidation during an experiment can have a large effect on aS, especially at near-neutral pH. For example, Figure 1 in ref 23 indicates that adventitious oxidation of only 1% of the sulfide in a pH 7 solution will raise aS nearly to saturation (specifically, log aS ≈ -0.2). This would be sufficient to measurably influence Sb solubility. Unrecognized oxidation may explain why some early workers (reviewed in 8) reported much higher Sb2S3 solubilities in the near neutral pH range than more recent workers (9, 32). Figure 5 shows calculated relationships among reduced Sb species at pH 7.5 based upon Table 3. Conditions in the deep Black Sea, which is often viewed as the classic example of a natural sulfidic environment, are shown by a hatched box. This figure suggests explanations for several previously observed aspects of Sb behavior in sulfidic environments. A key point is that the mixed-valence complex, HSb2S5-, would predominate in the Black Sea, according to this diagram. The redox state of Sb in a number of anoxic marine environments reveals the following pattern (14-17): at depth, where HS- first appears, SbIII concentrations rise, but SbV never disappears completely. This observation has intrigued previous workers. Cutter (17) devotes considerable discussion to alternate hypotheses for why SbV exists in the Black Sea, even though believed to be unstable. However, the results in this paper imply that SbV, in fact, can be stable in the Black Sea and probably in other anoxic marine basins. On the other hand, SbIII and SbV do not occur in a 1:1 ratio in the Black Sea, as would be predicted from Figure 5. This suggests (a) that we require still-better and more complete equilibrium information about Sb speciation, (b) that characterization of Sb’s redox state in Black Sea needs to be VOL. 36, NO. 5, 2002 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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improved, or (c) that dynamic processes in the deep Black Sea prevent complete redox equilibration of Sb, as has been proposed (17). Possibly, all are true. For example, fairly small changes in the stability constants used to construct Figure 5 could alter the predicted redox state. In this circumstance, even a temperature correction, which cannot be made with present data, might be large enough to alter predictions. Similarly, the true oxidation state of Sb in the deep Black Sea might be misrepresented by the BH4 reduction method, which is calibrated primarily with O-coordinated Sb standards. Information in this paper should be helpful in improving calibrations. Under the conditions represented by Figure 5, less than 7 µM of HS- will initiate transformation of Sb from Ocoordination to S-coordination over a broad range of aS. Below the transition, Sb in reducing environments occurs as uncharged molecules at near-neutral pH. Above the transition, Sb will be anionic. Because many mineral surfaces carry negative charges, transformation of Sb’s molecular charge from neutral to negative could promote desorption and release of dissolved Sb. How Sb interacts with mineral surfaces is an important question with respect to Sb mobility in reducing environments in aquifers. The data in Table 1 show that stibnite solubility in neutral to alkaline waters is large as compared to the drinking water standard of 50 nM. Therefore, precipitation of stibnite will not usefully restrain mobility of Sb in groundwater.
(9) (10) (11) (12)
Acknowledgments
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This work supported by the National Science Foundation (EAR9980532).
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Received for review August 21, 2001. Revised manuscript received December 17, 2001. Accepted December 17, 2001. ES011227C
