Environ. Sci. Technol. 1997, 31, 2148-2153
Inorganic Speciation of Mercury in Sulfidic Waters: The Importance of Zero-Valent Sulfur KRISTINA E. PAQUETTE† AND GEORGE R. HELZ* Department of Chemistry and Biochemistry and Water Resources Research Center, University of Maryland, College Park, Maryland 20742
Chemical speciation governs the biologic behavior of Hg. Because sulfide-producing microoganisms accomplish methylation, Hg speciation in sulfidic habitats is of scientific interest. Here, speciation is determined from solubilities of cinnabar (HgS) only and of cinnabar + sulfur in 0.7 M KCl medium at 298 K over the ranges, pH 1-12, and total dissolved sulfide, 10-3-10-1 M. Without dissolved zerovalent sulfur (S0), solubilities can be explained by the following: HgS(cinn) + H2S a Hg(SH)20, pK ) 5.36 ( 0.10; HgS(cinn) + SH- a HgS(SH)-, pK ) 5.34 ( 0.30; and HgS(cinn) + 2SH- / HgS22- + H2S, pK ) 7.14 ( 0.16 (all at I ) 0.7 M). Unlike ZnS and CdS, HgS solubility is increased by S0, which promotes formation of bidentate polysulfide ligands. Additional Hg solubility in S0 saturated solutions up to pH 9.5 can be explained by: HgS(cinn) + SH- + (n - 1)S0(rhom) a Hg(Sn)SH-, pK ) 3.97 ( 0.17; n could not be determined but is probably 4-6. In nearneutral, sulfidic natural waters, previously unknown Hg(Sn)SH- might commonly exceed other inorganic Hg(II) species because dissolved S0 occurs widely, especially near redox fronts where it is generated by biotic and abiotic oxidation of sulfide.
Introduction Discovery that numerous northern lakes contain fish with hazardous Hg concentrations (1-3) has sparked new interest in the environmental chemistry of Hg. Because most of the Hg in fish is methylated (4, 5) and because sulfate-reducing microorganisms are important producers of methylmercury (6-9), there is particular interest in understanding Hg chemistry in sulfidic waters where these microorganisms live. Mason et al. (10) show how knowledge of Hg speciation improves understanding of mercury’s transport through food chains. Schwarzenbach and Widmer (11), who studied the solubility of precipitated, “black HgS”, have contributed the benchmark work on Hg speciation in sulfidic solutions. However, for environmental chemists, their work suffers three shortcomings: (a) the nature of the solid phase is uncertain because later work has shown that pure metacinnabar, the black polymorph of HgS, is unstable below 315 °C (12) and inverts to cinnabar (red HgS) rapidly (13); (b) the sulfide concentration was fixed at 0.02 M, which is unrealistically high for most natural environments; and (c) the possible influence of zero-valent sulfur (S0), present as polysulfides, on the solubility and speciation of Hg was neglected; * Corresponding author telephone: (301)405-1797; fax: (301)3149121; e-mail:
[email protected]. † Present address: Division of Product Manufacture and Use, U.S. Food and Drug Administration, HFS-247, Washington, DC 20204.
2148
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 31, NO. 7, 1997
FIGURE 1. Variation of the mole fraction of HgS (left-hand scale), the activity of HgS (left-hand scale), and the logarithm of the activity of elemental Hg (right-hand scale) in cinnabar, a non-stoichiometric solid phase, as a function of the logarithm of the activity of elemental sulfur. Standard states are aHgS ) 1 for Hg1.000S, aHg0 ) 1 for solutions saturated with Hg(l) and aS0 ) 1 for solutions saturated with rhombic sulfur. Geochemists sometimes describe sulfide equilibria by the fugacity of S2(g). At 298 K, aS0 and fS2 are related by log aS0 ) 0.5 log fS2 + 6.98. Cinnabar’s reflectivity at λ > 575 nm, which is responsible for its brilliant red color, decreases with increasing log aS0 (12). polysulfides bind strongly to certain other metals (14-17). Subsequent investigations (18-20) have produced Hg speciation models in conflict with that of Schwarzenbach and Widmer. In this paper, we investigate the solubility of cinnabar over a wider range of conditions than Schwarzenbach and Widmer in order to explore these issues. We also present the first systematic study of the role of S0 on Hg(II) speciation.
Chemical Background for This Study At 298 K and 1 atm, cinnabar is the only stable compound in the Hg-S system. Above 315 °C, cinnabar transforms to metacinnabar. Impure forms of the latter can be stabilized at ambient temperature by Fe, Zn, and Se (13). As is common with sulfide phases (ref 21, p 66), cinnabar is a nonstoichiometric solid in which Hg vacancies are coupled with interstitial S substitution (12). Extrapolation of measurements at 375-600 K (12) suggests that stable cinnabar can have compositions of Hg1.000S to Hg0.9673S at 298 K. Figure 1 illustrates two generally unrecognized consequences of non-stoichiometry. First, the activity of HgS must decline as the HgS content of cinnabar declines with increasing aS0. At log aS0 ) 0, i.e., rhombic sulfur saturation, aHgS is estimated to be 0.76 (see Appendix A, Supporting Information, for details on how this estimate is made). Concentrations of mononuclear Hg complexes in cinnabarsaturated solutions scale with aHgS, so they may vary by ∼25% due to variations in cinnabar composition. The estimated magnitude of this effect is small in relation to experimental uncertainties in this study, so we do not discuss it further. The second effect of non-stoichiometry is the reduction in the reflectance of cinnabar at optical wavelengths >575 nm (Figure 1). This reflectance gives cinnabar its vermillion color. For metastable cinnabar compositions of Hge0.9S, this reflectance virtually disappears (12). Such materials would appear black, especially if present as finely divided precipitates. Thus red color is an unreliable diagnostic of HgS polymorphism. We will return to this point.
S0013-936X(96)01001-2 CCC: $14.00
1997 American Chemical Society
Figure 1 shows that cinnabar of varying composition is stable over a 7 log unit range in aS0. This aS0 variation may, in principle, have a large influence on Hg solubility through its control of the stability of aqueous species. Two kinds of aS0 sensitive species must be considered, polymercurous cations (Hgn2+) and polysulfides (HxSnx-2, x ) 0-2). The first are destabilized by rising aS0 (due to concomitant decline in aHg0, Figure 1) while the second are stabilized. In the solid state, a series of homopolyatomic mercury cations, Hgn2+ are known, including the so-called mercurous ion, Hg22+ (22-24). In the presence of electron-pair donating ligands in water, these species tend to disproportionate to Hg0 plus Hg(II) species:
Hg22+ + nL- f HgIILn2-n + Hg0
(1)
Here, L- designates a general ligand, which in principle could be even a water molecule. For example, in calomel (Hg2Cl2)-saturated water, the Hg22+ ion is decomposed by H2O and Cl- during dissolution, and Hg(II) species predominate in the aqueous phase (25). Thiol and sulfide ligands very strongly drive reaction 1 to the right, not only by forming extremely stable complexes with HgII but also by oxidizing Hg0, releasing H2 (26). For these reasons, polyatomic Hg cations can be neglected in sulfidic solutions. On the other hand, the importance of mercury polysulfides is an open question. In sulfidic natural waters, evidence for the common occurrence of dissolved, zero-valent S is accumulating (27-30). Zero-valent sulfur dissolves readily in sulfidic solutions forming polysulfides:
(n - 1)S0 + HS- + (x - 1)H+ a HxSnx-2 x ) 0-2, n ) 2-6 or more (2) Isotope exchange measurements (31) and other evidence indicate that these reactions are labile and therefore amenable to thermodynamic treatment (15, 16). From mass action laws for reaction 2, concentrations of polysulfides can be shown to scale as aS0(n-1). Consequently, longer chain (higher n) polysulfides are significant only at high aS0, near sulfur saturation (15, 32, 33). The n g 4 polysulfides are effective bidentate chelating agents for Group IB monovalent cations (14-17). Ring strain curtails a bidentate role for shorter polysulfides. Exploratory results suggest that Zn(II) and Cd(II) fail to form detectable polysulfide complexes (34, 35), but the possibility of mercury(II) polysulfides needs to be investigated. Mercury(II) polysulfide complexes are known in solids (36). In this paper, we report cinnabar solubility under two log aS0 conditions. In the first series, no S0 was introduced and care was taken to avoid its production through sulfide oxidation. The intent in this series was to position log aS0 at undefined but low levels where larger polysulfides (e.g., S42and S52-) would be negligible. In the second series, excess rhombic sulfur was added so that the solutions were always saturated with S0, poising aS0 at unity. A short, preliminary report on this work was published previously (37).
Experimental Section Cinnabar was synthesized at 400-450 °C in evacuated quartz tubes from triply distilled Hg(l) and 99.999% pure sulfur powder. Repeated boiling in concentrated K2CO3 solution removed any excess sulfur, and leaching for 2-3 days in 0.1 M L-cysteine removed any high-energy crystal edges and corners. X-ray diffraction and X-ray emission spectra confirmed cinnabar. The following buffers were prepared in 0.7 KCl: pH 1-3, HCl/KCl; pH 4-5, HOAc/NaOAc; pH 6-8, NaH2PO4/Na2HPO4; pH 9-11, Na2B4O7/NaOH; and pH 12-13, KOH/KCl. The pH of these solutions was measured with a pH electrode standardized at two points against commercial buffers. All
solutions were deoxygenated by purging with N2 for at least 30 min prior to preparation of solubility experiments. Solubility procedures were similar to those used previously in this laboratory (34, 35, 37). Experiments were conducted in 50-mL glass ampules charged with 0.3 g of cinnabar, 50 mL of the desired buffer, and the desired aliquot of 1 M bisulfide solution. Ampules for determining mercury polysulfide speciation were also charged with 0.3 g of 99.999% pure sulfur powder. These operations were performed in a N2-filled glovebox. After fusion sealing, samples were allowed to equilibrate up to several months in a thermostated chamber (25 °C) with periodic shaking by hand. After equilibration, ampules were opened in a glovebox, syringe-filtered through 0.02-µm Anotop 25 inorganic membrane filters (Whatman) and analyzed. Total divalent sulfur (∑S-II) was measured either by potentiometric titration with HgCl2 or by iodimetric titration. Precision was ∼8% for the first procedure and ∼4% for the latter, based on triplicate determinations. Total mercury was measured via cold vapor atomic absorption (Perkin-Elmer Model 2380 AA spectrophotometer) at 253.7 nm, based on an established method (38-41). Estimated precision was 6%. Prior to Hg determinations, sulfur species were oxidized with a 0.9 M BrCl solution. This step was followed by hydroxylamine reduction to eliminate the excess BrCl, which would otherwise consume SnCl2 in the Hg0 generation flask. For full details about methods, see ref 42.
Results Sixty-eight solubility measurements were collected over the pH range, 1-12, and the total sulfide range, 10-3-10-1 M. Complete data are available in the Supporting Information and in ref 42. The variation of solubility with pH is illustrated in Figure 2, which shows data selected to fall within two restricted ∑S-II windows. The upper panel displays results obtained in the absence of S0, whereas the lower panel displays results from S0 saturated solutions. For comparison, solubilities calculated from equilibrium constants of Schwarzenbach and Widmer (11) at the midpoints of each ∑S2- window are shown. Inspection reveals that solubilities calculated from the Schwarzenbach and Widmer constants, which were obtained from ∑S-II ) 20 mM solutions, reproduce the new data reasonably well. It also appears that S0 modestly enhances observed solubilities as compared to the predictions, especially at neutral to alkaline pH where S0 readily reacts with HS- to form polysulfide ions. Various speciation models were tested with a nonlinear least squares routine (SCIENTIST, MicroMath Inc.) by fitting the new experimental data. The tests of a good fit were that log [∑Hgcalc/∑Hgobs] (a) did not vary systematically with the independent parameters, pH and ∑S-II (b) had a mean value near zero, and (c) had the smallest standard deviation achievable. Stoichiometric equilibrium constants derived from several models are presented in Table 1. The constants are appropriate for 0.7 M KCl at 298 K; uncertainties specified are (1σ. The H+ ion was equated to 10-pH, where pH is measured on the NBS scale. We used a value of 10-6.68 for the first ionization constant of H2S in 0.7 M NaCl (43, 44). Analysis of S0 Unsaturated Data. As a first step, we simply refined the Schwarzenbach and Widmer model to minimize squares of deviations with respect to the new data. This yielded model 2 in Table 1. Figure 3A shows the deviations in log [∑Hgcalc/∑Hgobs] before and after adjusting the constants. The new constants agree well with the original values (model 1), although a slightly higher constant in the acidic region and slightly lower constants in the alkaline region are obtained. The absence of a systematic offset suggests that Schwarzenbach and Widmer’s “black HgS” may in fact have been cinnabar, presumably a S-rich, non-stoichiometric variant having a diminished reflectivity for red light. However, this conclusion is tentative. Free energy data (12) suggest
VOL. 31, NO. 7, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2149
TABLE 1. Alternate Models Used To Explain the Solubility Results log10 equilibrium constants (298 K, I ) 0.7) reaction (1) HgS(cinn) + H2S(aq) a Hg(SH)20 (2) HgS(cinn) + SH- a HgS(SH)(3) HgS(cinn) + 2SH- a HgS22- + H2S (4) HgS(cinn) + 2H2S(aq) a Hg(H2S)(SH)20 (5) HgS(cinn) + SH- + H2S(aq) a Hg(SH)3(6) HgS(cinn) + (n - 1)S(rhom) + SH- a Hg(Sn)(SH)standard deviation of log [Hgcalc/Hgobs]
model 1a
model 2b
-5.76 -5.17 -6.73
-5.41 ( 0.12 -5.41 ( 0.29 -7.19 ( 0.24
0.68
0.57
model 3c
-6.93 -3.43 -2.27 0.86
model 4d -5.36 ( 0.10 -5.34 ( 0.30 -7.14 ( 0.16 -3.97 ( 0.17 0.57
a Schwarzenbach and Widmer (11) K values corrected to be consistent with K ) 10-6.68 for H S and corrected from I ) 1.0 to I ) 0.7 with Davies a 2 eq. b Same model as 1, but with equilibrium constants calculated by least squares fit to low aS0 data (n ) 43). c Barnes et al. model (18) optimized to fit low aS0 data in this paper; HgS(HS)22-, proposed in ref 18, did not improve this fit. d Combined fit to both low aS0 and S saturated data (n ) 68).
FIGURE 2. Mercury concentrations as a function of pH in two subsets of the data. In each diagram, the solubilities calculated from K values of Schwarzenbach and Widmer (11) are shown for comparison. The upper diagram depicts S0-absent results, whereas the lower diagram depicts S0-saturated results. Uncertainties of purely analytical origin (∼15%) are smaller than the symbols. The scatter arises in part from the 3-fold range of ∑S-II in each diagram. It also reflects systematic variability arising from sluggish relaxation of excess solid-state free energy that is produced by such things as small particle size, strain, defects, stacking faults, departures from equilibrium stoichiometry, etc. that metacinnabar would be only 0.4 log units more soluble than cinnabar, a difference possibly concealed by uncertainties in Table 1. Figure 3B presents a test of the speciation model of Barnes et al. (18). The residual trend indicates that the higher order dependence on aH2S implicit in this model is inconsistent with data in the ∑S-II < 0.1 M range. Our analysis does not prove that reactions 4 and 5 in Table 1 are unimportant at total sulfide concentrations above this range, but such sulfide concentrations would rarely be of concern in environmental chemistry. It is tempting to criticize model 3 by noting that mercury appears to have 3-coordination in the complexes in reactions
2150
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 31, NO. 7, 1997
FIGURE 3. Deviation plots showing fits to the low aS0 data. Horizontal lines are guides for the eye. (A) Open symbols show deviations produced when solubilities are calculated from K values of Schwarzenbach and Widmer (model 1, Table 1); closed symbols show deviations when these K values are optimized by least squares fitting (model 2, Table 1). (B) Open symbols show deviations produced by K values of Barnes et al. (18); closed symbols show deviations after optimization (model 3, Table 1); optimization does not remove the trend produced by this model. (C) Deviations produced by Zhang and Millero K values. 4 and 5. Mercury generally possesses 2- or 4-coordination when surrounded by S ligands, although 3-fold coordination is known (45). However, solubility data obtained under the constraints that aHgS ≈ 1 and aH2O ≈ 1 cannot establish the number of formula units of HgS and H2O in a complex and therefore do not establish the coordination number. For example, without reducing the quality of the fits, HgS22- in all the models of Table 1 could be replaced by Hg(OH)2(SH)22-. The latter differs from HgS22- by only two H2O molecules, but suggests a 4-coordinate rather than 2-coordinate complex. As shown in Figure 3C, the stability constants of Zhang and Millero (20) for Hg(HS)+ and Hg(HS)20 vastly underpredict cinnabar solubilities. Zhang and Millero used cathodic
FIGURE 4. Species distribution vs pH. In solutions saturated with rhombic sulfur (left-hand plot), the polysulfide complex Hg(Sn)(SH)predominates in the near-neutral pH range important in most natural waters. Conditions: I ) 0.7, 298 K. Chloride ligands are unable to compete with sulfide and polysulfide ligands in solutions containing >10-10 M sulfide. stripping voltammetry to monitor free sulfide as they titrated ∼10-6 M NaHS into seawater containing ∼10-6 M total dissolved Hg. Equilibration time after aliquot additions was 1 min. Two kinds of systematic error could have intervened. First, as the authors acknowledge, the voltammetric method might have detected some Hg-bound sulfide, causing overestimation of free S-II and corresponding underestimation of stability constants. Second, because the test conditions were far from equilibrium, the bound sulfide might have become kinetically trapped in a metastable species, consistent with the very low solubilities predicted from the resulting stability constants (Figure 3C). A number of other speciation models were postulated and tested in the course of this work, including the fairly complex one described in our preliminary report (37). In the end, none produced improvement upon model 2 when tested against only the low aS0 data. Analysis of S0 Saturated Data. According to Figure 2, sulfur saturated solutions display higher solubility than can be accounted for by the reactions in models 1 and 2, indicating the need to consider mercury-polysulfide complexes. To investigate this, we used the equilibrium constants in ref 15 to calculate the concentration of free polysulfides at saturation with rhombic sulfur in each of the high aS0 samples. This can be done from the known equilibrium constants for reactions of the form of reaction 2 (15), making use of the observed pH and ∑S-II values and letting aS0 ) 1. The calculation requires no knowledge of mercury polysulfide complexes because ∑Hg is negligible as compared to the major dissolved sulfur species in the solutions under consideration. The calculated concentrations of free polysulfides were then subtracted from ∑S-II to obtain the concentrations of total free sulfide, which was partitioned between H2S and HS- according to pH. This procedure corrects for the consumption of free sulfide during the formation of polysulfides. Concentrations of mercury polysulfide complexes were modeled with reactions of the form
(n - 1)S(rhom) + xH+ + yHS- + HgS(cinn) a mercury polysulfide (3) Various reactions of this nature were proposed and tested by fitting their equilibrium constants to the experimental data; aS0 was assigned a unit value for the high aS0 samples. A single mercury polysulfide complex, Hg(Sn)(SH)-, produced an excellent fit (model 4, Table 1). This species can be viewed as a sulfurized derivative of HgS(SH)-; i.e.
HgS(SH)- + (n - 1)S0 a Hg(Sn)SH)-
(4)
We conjecture that at high pH, a sulfurized analog of HgS22also forms: Hg(Sn)22-. Our high aS0 data extend only to pH 9.4, and in this range, Hg(Sn)22- produced insignificant improvement in the fit. Nevertheless, salts of Hg(S6)22- and Hg(S4)22- have been synthesized, and their structures have been determined (36); their two polysulfide chains form bidentate rings that tetrahedrally coordinate Hg2+ ions. Although the additional solubility caused by S0 can be represented by a single reaction, several mercury polysulfides, differing in n, may in fact exist (as is true for free polysulfides). Figure 4 shows the distribution of species calculated from model 4 in Table 1. A key point is that over a broad, nearneutral pH range, Hg(Sn)(SH)- is the predominant Hg(II) complex in solutions saturated with S0. The value of n in Hg(Sn)(SH)- cannot be determined from our type of experiment for the following reason. General experience suggests that n lies in the range 4-6. According to reaction 4, such complexes would decline as the 3rd to 5th power of aS0. Because Hg(Sn)(SH)- exceeds HgS(SH)- by less than an order of magnitude at S0 saturation (Figure 4), more than a mere 2-fold decline in aS0 would submerge Hg(Sn)(SH)with respect to HgS(SH)-, rendering it unobservable by solubility methods. Controlling aS0 accurately within such a narrow range exceeds current capabilities. Implications Regarding Natural Waters. An important finding in this paper is that S0 promotes Hg solubility. A
VOL. 31, NO. 7, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2151
substantial body of information demonstrates that dissolved S0 is common (a) in water columns of anoxic marine basins (28, 29, 46-48), (b) in and above anoxic wetlands (27, 4951), and (c) in anoxic marine sediments (52-54). In many of these examples, dissolved S0 peaks near the oxic/sulfidic interface due to biotic and abiotic sulfide oxidation. While existing studies quantify the amount of S0, most have failed to define its activity, aS0, which can be calculated from measurements of ∑S-II, pH, and ∑S0. Thompson and Helz (16) used this method to estimate aS0 in the water column of the Black Sea, concluding that it ranged from 0.4 near the top of the sulfide zone to 0.06 at depth. Using this approach with the data of Luther et al. (49) yields values from 3.7 to 7.0 (indicating supersaturation) in the Great Sippewissett Salt Marsh and from 0.76 to 1.05 in a subtidal core from Orleans, MA. Similarly, the data from Great Marsh, Lewes, DE (27), suggest aS0 values between 1.0 and 1.6 (15). While the evidence remains sparse, aS0 values in natural anoxic environments appear to fall commonly in the range 0.1-10. This range is critical with respect to Hg behavior. At the lower end, Hg(Sn)(SH)- will be negligible relative to HgS(SH)-, whereas in the middle and upper part of this range Hg(Sn)(SH)- will predominate at near-neutral pH. Evidence that aS0 exceeds unity in salt marshes where sulfide oxidation is rapid has important implications regarding Hg mobility and bioavailability. Under conditions where Hg(Sn)(SH)- is the predominant species, the solubility of cinnabar scales as aS0(n-1) (reaction 4). The value of n is probably g4; a smaller n would deprive these mercury polysulfides of the extra stability available from the chelate effect, because the polysulfide chain would be too short to form a ring that incorporates the central Hg ion. Thus the solubility of HgS is probably proportional to aS0g3. A 10-fold supersaturation with respect to rhombic S0 (i.e., aS0 ) 10) therefore produces at least 1000-fold increase in HgS solubility. Hence, high Hg mobility should be expected in environments, such as salt marshes, where rapid S0 production leads to S0 supersaturation. Transformation of HgS(SH)- to Hg(Sn)(SH)- with increasing aS0 produces a less well solvated form of dissolved Hg. On a molar basis, molecular sulfur’s (S8) intrinsic solubility (55) is less than the solubility of the four-ring polycyclic aromatic hydrocarbons. This comparison suggests that enclosure of an Hg2+ ion by a ring of sulfur atoms might increase mercury’s rate of uptake by living cells.
Acknowledgments Funding for this study was provided by NSF Grants EAR9206542 and EAR-9405432 and by the Philadelphia Academy of Natural Sciences (ANS). We would like to give special thanks to C. Gilmour and G. Riedel (ANS, St. Leonard, MD) for their advice on optimizing our cold vapor Hg apparatus, to B. LaSorsa (Battelle Marine Sciences Laboratory) for the recipe for BrCl, and to J. C. Cornwell (University of Maryland) for references on zero-valent sulfur in nature.
Supporting Information Available Appendix A giving the effect of cinnabar’s non-stoichiometry on its solubility and Appendix B giving the solubility data (3 pp) will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the Supporting Information from this paper or microfiche (105 × 148 mm, 24× reduction, negatives) may be obtained from Microforms Office, American Chemical Society, 1155 16th St. NW, Washington, DC 20036. Full bibliographic citation (journal, title of article, names of authors, inclusive pagination, volume number, and issue number) and prepayment, check or money order for $12.00 for photocopy ($14.00 foreign) or $12.00 for microfiche ($13.00 foreign), are required. Canadian residents should add 7% GST. Supporting Information is also available via the World Wide Web at URL http://www.chemcenter.org.
2152
9
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 31, NO. 7, 1997
Users should select Electronic Publications and then Environmental Science and Technology under Electronic Editions. Detailed instructions for using this service, along with a description of the file formats, are available at this site. To down-load the Supporting Information, enter the journal subscription number from your mailing label. For additional information on electronic access, send electronic mail to
[email protected] or phone (202) 872-6333.
Literature Cited (1) Lindqvist, O.; Johansson, K.; Aastrup, M.; Andersson, A.; Bringmark, L.; Hovsenius, G.; Håkanson, L.; Iverfeldt, Å.; Meili, M.; Timm, B. Water Air Soil Pollut. 1991, 55, 1-251. (2) Weiner, J. G.; Stokes, P. M. Environ. Toxicol. Chem. 1990, 9, 821-823. (3) Driscoll, C. R.; Yan, C.; Schofield, C. L.; Munson, R. Holsapple, J. Environ. Sci. Technol. 1994, 28, 136A-143A. (4) Grieb, T. M.; Driscoll, C. R.; Gloss, S. P.; Schofield, C. L.; Bowie, G. L.; Porcella, D. B. Environ. Toxicol. Chem. 1990, 9, 919-930. (5) Bloom, N. S. Can. J. Fish. Aquat. Sci. 1992, 49, 1010-1017. (6) Gilmour, C. C.; Henry, E. A. Environ. Pollut. 1991, 71, 131-169. (7) Gilmour, C. C.; Henry, A. A.; Mitchell, R. Environ. Sci. Technol. 1992, 26, 2281-2287. (8) Watras, C. J.; Bloom, N. S.; Claas, S. A.; Morrison, K. A.; Gilmour, C. C.; Craig, S. R. Water Soil Air Pollut. 1995, 80, 735-745. (9) Matilainen, T. Water Air Soil Pollut. 1995, 80, 757-764. (10) Mason, R. P.; Reinfelder, J. R.; Morel, F. M. M. Environ. Sci. Technol. 1996, 30, 1835-1845. (11) Schwarzenbach, G; Widmer, M. Helv. Chim. Acta 1963, 46, 26132628. (12) Potter, R. W.; Barnes, H. L. Am. Mineral. 1978, 63, 1143-1152. (13) Dickson, F. W.; Tunell, G. Am. Mineral. 1959, 44, 471-487. (14) Cloke, P. L. Geochim. Cosmochim. Acta 1963, 27, 1229-1319. (15) Shea, D.; Helz, G. R. Geochim. Cosmochim. Acta 1988, 52, 18151825. (16) Thompson, R. A.; Helz, G. R. Geochim. Cosmochim. Acta 1994, 58, 2971-2983. (17) Berndt, M. E.; Bluttram, T.; Earley, D., III; Seyfried, W. E., Jr. Geochim. Cosmochim. Acta 1994, 58, 587-594. (18) Barnes, H. L.; Romberger, S. B.; Stemprok, M. Econ. Geol. 1967, 62, 957-982. (19) Shikina, N. D.; Zotov, A. V.; Khodakovskii, I. L. Geochem. Int. 1981, 109-117. (20) Zhang, J.-Z.; Millero, F. J. Anal. Chim. Acta 1994, 284, 497-504. (21) Vaughan, D. J.; Craig, J. R. Mineral Chemistry of Metal Sulfides; Cambridge University Press: Cambridge, l976; 493 pp. (22) Cutforth, B. D.; Davies, C. G.; Dean, P. A. W.; Gillespie, R. J.; Ireland, P. R.; Ummat, P. K. Inorg. Chem. 1973, 12, 1343(23) Schultz, A. J.; Williams, J. M.; Mire, N. D.; MacDiarmid, A. G.; Heeger, A. J. Inorg. Chem. 1978, 17, 646-649. (24) Cutforth, B. D.; Gillespie, R. J.; Ireland, P.; Sawyer, J. F.; Ummat, P. K. Inorg. Chem. 1983, 22, 1344-1347. (25) Marcus, Y. J. Phys. Chem. Ref. Data 1980, 9, 1307-1323. (26) Yamamoto, M. Chemosphere 1995, 31, 2791-2798. (27) Boulegue, J.; Lord, C. J.; Church, T. M. Geochim. Cosmochim. Acta 1982, 46, 453-464. (28) Hastings, D.; Emerson, S. Limnol. Oceanogr. 1988, 33, 391-396. (29) Luther, G. W., III. In Black Sea Oceanography; Izda, E., Murray, J. W., Eds.; Kluwer Academic Publishers: Dordrecht, The Netherlands, 1991; pp 187-204. (30) Luther, G. W., III; Ferdelman, T.; Tsamakis, E. Estuaries 1988, 11, 281-285. (31) Fossing, H.; Jørgensen, B. B. Biogeochemistry 1990, 9, 223-245. (32) Schwarzenbach, G.; Fischer, A. Helv. Chim. Acta 1960, 43, 13651390. (33) Giggenbach, W. F. Inorg. Chem. 1974, 13, 1724-1730. (34) Daskalakis, K.; Helz G. R. Geochim. Cosmochim. Acta 1993, 57, 4923-4931. (35) Daskalakis, K.; Helz, G. R. Environ. Sci. Technol. 1992, 26, 24622468. (36) Mu ¨ ller, A.; Schimanski, J.; Schimanski, U.; Bo¨gge, H. Z. Naturforsch. 1985, 40B, 1277-1288. (37) Paquette, K.; Helz, G. R. Water Air Soil Pollut. 1995, 80, 10531056. (38) Gill, G. A.; Fitzgerald, W. F. Mar. Chem. 1987, 20, 227-243. (39) Hatch, W. R.; Ott, W. L. Anal. Chem. 1968, 40, 2085-2087. (40) Bloom, N.; Fitzgerald, W. R. Anal. Chim. Acta 1988, 208, 151161. (41) Ingle, J. D., Jr.; Crouch, S. R. Spectrochemical Analysis; Prentice Hall: Englewood Cliff, NJ, 1988.
(42) Paquette, K. E. Solubility of Cinnabar (Red HgS) and Implications for Mercury Speciation in Sulfidic Waters. Ph.D. Dissertation, University of Maryland, College Park, MD 1994. (43) Goldhaber, M. B.; Kaplan, I. R. Mar. Chem. 1975, 3, 83-104. (44) Williamson, M. A.; Rimstidt, J. D. Geochim. Cosmochim. Acta 1992, 56, 3867-3880. (45) Bowmaker, G. A.; Dance, I. G.; Dobson, B. C.; Rogers, D. A. Aust. J. Chem. 1984, 37, 1607-1618. (46) Jacobs, L.; Emerson, S. Earth Planet. Sci. Lett. 1982, 60, 237252. (47) Luther, G. W., III; Church, T. M.; Powell, D. Deep-Sea Res. 1991, 38 (Suppl.), S1121-S1137. (48) Jørgensen, B. B.; Fossing, H.; Wirsen, C. O.; Jannasch, H. W. Deep-Sea Res. 1991, 38 (Suppl.), S1083-S1103. (49) Luther, G. W., III; Giblin, A. E.; Varsolona, R. Limnol. Oceanogr. 1985, 30, 727-736. (50) Cutter, G. A.; Velinsky, D. J. Mar. Chem. 1988, 23, 311-327.
(51) Witcomb, J. H.; Delaune, R. D.; Patrick, W. H., Jr. Mar. Chem. 1989, 26, 205-214. (52) Troelsen, H.; Jørgensen, B. B. Estuarine, Coastal Shelf Sci. 1982, 15, 255-266. (53) Schimmelmann, A.; Kastner, M. Geochim. Cosmochim. Acta 1993, 57, 67-78. (54) Gagnon, C.; Mucci, A.; Pelletier, E. Geochim. Cosmochim. Acta 1995, 59, 2663-2675. (55) Boulegue, J. Phosphorus Sulfur 1978, 5, 127-128.
Received for review December 3, 1996. Revised manuscript received February 4, 1997. Accepted March 4, 1997.X ES961001N X
Abstract published in Advance ACS Abstracts, May 15, 1997.
VOL. 31, NO. 7, 1997 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
9
2153
