Environ. Sci. Technol. 1999, 33, 4270-4277
Cadmium Complexation with Bisulfide F E I Y U E W A N G † A N D A N D R EÄ T E S S I E R * INRS-Eau, Universite´ du Que´bec, 2700 Rue Einstein, C.P. 7500, Sainte-Foy, Que´bec, Canada G1V 4C7
We have used in situ dialysis to measure the solubility of CdS(s) in sulfidic solutions. The solubility product for the reaction CdS(s) + H+ a Cd2+ + HS- at 25 °C and 1 atm was found to be 10-14.82(0.03 for a crystalline product and to be 10-14.15(0.06 and 10-14.40(0.06 for two precipitates, respectively. We show that the solubility of these three solids at various pH values (4.2-8.6) and sulfide concentrations (10-4.3-10-1.3 M) can be reproduced adequately by the following four bisulfide complexes: CdHS+, Cd(HS)2, Cd(HS)3-, and Cd(HS)42-; the log Kn values for the general equation Cd2+ + nHS- a Cd(HS)n2-n are 7.38 ( 0.68, 14.43 ( 0.01, 16.26 ( 0.58, and 18.43 ( 0.05, respectively. The species CdOHS-, which has been reported previously, does not explain our experimental results. Calculations using estimated concentrations of organic ligands (humic substances and organic thiols) known to be present in natural waters indicate that sulfide complexes largely dominate Cd speciation in natural waters at ∑S(-II) g 10-6 M.
Introduction When present in anoxic environments, reduced sulfur species are likely to control dissolved cadmium concentrations. Wherever waters become devoid of oxygen (e.g., in porewaters, groundwaters, or the hypolimnion of productive lakes or fjords), sulfide ions may be present. Cadmium, a class B or “soft sphere” metal, forms covalent bonds with Scontaining ligands such as bisulfide (1). Thus, in sulfidic environments, Cd is expected to form solid sulfide phases as well as sulfide complexes; the latter are expected to dominate Cd speciation in sulfidic environments even in the presence of large concentrations of organic matter (2, 3). When Cd concentrations predicted with the available thermodynamic data are compared with measured Cd concentrations in sulfidic waters, apparent supersaturations with respect to Cd-sulfide phases are generally observed (3-5). This failure of equilibrium models to adequately predict Cd solubility can be due to kinetic factors, to the presence of unidentified organic ligands, to inadequate characterization of the solid phase that controls Cd solubility, or simply to inadequate thermodynamic data. This paper examines the lattermost possibility. The first influential work on Cd complexation by sulfide is that of Ste-Marie et al. (6), who interpreted the solubility of Cd in the Cd-S(-II)-H2O system in terms of complexes with the bisulfide ligand, Cd(HS)n2-n (n ) 1-4), and the hydroxo complex CdOH+. Their study was performed at high ionic strength (I ) 1 M) for only one initial concentration of * Corresponding author phone: (418)654-2632; fax: (418)654-2600; e-mail:
[email protected]. † Present address: EVS Environment Consultants, 195 Pemberton Avenue, North Vancouver, BC, Canada V7P 2R4. 4270
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sulfides, and it involved an uncharacterized Cd-sulfide solid. Important experimental details concerning the experimental data, equilibration time and solid-solution separation were also lacking in this work. Zhang and Millero (7) have determined formation constants for Cd(HS)+ and Cd(HS)2 by titrating Cd solutions prepared in seawater with sulfide and measuring the free sulfide concentration by cathodic stripping square wave voltammetry. This procedure assumes that sulfide complexes do not contribute to the measured current and that sulfide losses due to the reaction of sulfide with waste mercury is negligible (8). The stability constants they determined are conditional stability constants that are valid in the seawater medium only. As discussed in their paper, these stability constants could be underestimates because the authors did not take into account in their calculation the formation of Cd-chloro complexes and because they assumed that the Cd-sulfide complexes are nonlabile. Al-Farawati and van den Berg (8) determined conditional formation constants for Cd(HS)+ and Cd(HS)2 in seawater by flow analysiscathodic stripping voltammetry (FA-CSV) using either the direct titration of sulfide by Cd or a ligand competition method (8). The most trustworthy formation constants for Cd-sulfide complexes are probably those derived from the recent solubility study of Daskalakis and Helz (9). Their work covered a wide range of pH values (2.3-9) and of sulfide concentrations (∑S(-II) ) 10-3.4-100.6 M) and involved strict procedures to minimize sulfide oxidation as well as the use of a wellcharacterized CdS(s) solid phase (greenockite). Cadmium solubility was explained by three sulfide complexes: Cd(HS)3-, Cd(HS)42-, and CdOHS-. Although the range of sulfide concentrations used in their experiments can be reached in marine systems, it is higher than the sulfide concentrations normally found in freshwater anoxic systems (usually e10-5 M). As a consequence, the Cd-sulfide complexes that fit their data at high sulfide concentrations are not necessarily those [e.g., Cd(HS)+ and Cd(HS)2] that would predominate at the low sulfide concentrations of freshwaters. The above considerations stress the need for a more complete set of thermodynamic data on the Cd-sulfide system if we are to understand Cd fixation and mobility in freshwater sulfidic environments. In this paper, we present measurements of the formation constants for Cd-sulfide complexes. Our experimental conditions covered ranges of bisulfide concentrations and pH reaching into the realm of values appropriate for natural anoxic waters, particularly for freshwater porewaters.
Materials and Methods Manipulations were performed in a glovebox (Anaerobic System model 1025, Forma Scientific Inc.) filled with ultrapure N2 (99.996%). All solutions were made with deoxygenated ultrapure Milli-Q water (>18MΩ‚cm) and were further purged with N2 for at least 30 min before transferring them to the glovebox. All reagents were of analytical grade unless otherwise specified. Nitric acid (Anachemia) was of Environmental Grade. Ionic strength was adjusted with sodium nitrate (Suprapur, Merck). Buffer solutions were prepared either with sodium acetate (Suprapur; Merck)-acetic acid (Environmental Grade, Anachemia), 2,4-lutidine (Sigma), or imidazole (Sigma). Bisulfide stock solutions were freshly prepared from Na2S (Aldrich) and standardized by iodometry. We used three kinds of CdS(s) in our study. High-purity crystalline CdS(s) (called herein CdS(s)I) was purchased from Aldrich (powder, 100 days). They reported a value of log Ksp) -14.36 ( 0.26 at 25 °C and I ) 0 (Table 2). However, recalculation of their Ksp with the 4272
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I (M)
pH
0.097 0.051 0.017 0.011 0.047 0.013 0.007
7.61 7.75 7.81 8.05 7.93 8.21 8.16
0.097 0.051 0.017 0.011 0.047 0.013 0.007
7.62 7.95 7.90 8.10 7.50 7.78 8.07
0.097 0.051 0.017 0.011 0.047 0.013 0.007
7.85 8.02 8.12 8.25 7.95 8.07 8.15
logΣ[EDTA]
logΣ[S(-II)]
CdS(s)I from Aldrich -2.94 -6.16 -2.94 -5.98 -2.94 -5.87 -2.94 -5.80 -3.64 -6.46 -3.64 -6.30 -3.64 -6.17
logΣ[Cd]
logKsp (I ) 0)
-4.77 -4.89 -4.90 -5.00 -5.19 -5.23 -5.29
-14.78 -14.78 -14.81 -14.85 -14.85 -14.86 -14.83
CdS(s)II Precipitated from Thioacetamide -2.94 -5.92 -4.55 -2.94 -5.83 -4.65 -2.94 -5.83 -4.59 -2.94 -5.76 -4.68 -3.64 -6.35 -4.83 -3.64 -6.13 -4.90 -3.64 -6.20 -4.87 CdS(s)III Precipitated from Na2S -2.94 -5.82 -4.45 -2.94 -5.80 -4.49 -2.94 -5.64 -4.50 -2.94 -5.73 -4.47 -3.64 -6.12 -4.78 -3.64 -6.05 -4.82 -3.64 -6.04 -4.82
-14.31 -14.38 -14.44 -14.48 -14.41 -14.37 -14.43 -14.09 -14.17 -14.15 -14.23 -14.08 -14.17 -14.21
computer program HYDRAQL from their experimental data using the equilibrium constants reported in their paper gives lower values of Ksp (Table 2). The discrepancy between the Ksp value reported by Daskalakis and Helz (9) and the one we calculated from their experimental data is probably a consequence of an error made in their calculation of the equilibrium constant (K1) corresponding to their eq 1:
CdL3- + 2H+ ) Cd2+ + H2L3-
K1 ) 10-0.12 (I ) 0.1) (3)
where I is ionic strength and L is for DTPA. They reported (9) a value of K1 ) 100.22 at infinite dilution when correction is made with the Davies equation. Using the Davies equation, we calculate that K1 ) 10-0.34 or 10-0.56 depending if we consider their K1 as a mixed constant (i.e., where H+ is expressed as activity) or not. Applying the corrected values of K1 to the calculation of their Ksp leads to log Ksp ) -14.92 or log Ksp ) -15.14. These values are relatively close to the value we obtain for CdS(s)I in the present study (Table 2). Accuracy of the value of Ksp obtained by the indirect solubility method that we used is strongly dependent on the accuracy of the stability constants for the complex of the strong ligand with H+, with the metal studied (Cd) and with the cations involved to adjust the ionic strength (e.g., Na) since these constants are all influential in the calculation of [Cd2+]. The 4th acidity constant of EDTA (or the 5th acidity constant of DTPA) and the stability constants of EDTA (or DTPA) with monovalent cations such as Na+ are not wellestablished because of experimental difficulties involved in their determination (20). In the present study, we adopted the set of constants for EDTA given in NIST (14; see Table A in Supporting Information) for the calculation of Ksp; using other consistent sets of EDTA constants would have given slightly different Ksp values. Cadmium-Bisulfide Complexes. The solubility of cadmium in bisulfide solutions was measured at various values of ∑S(-II) (10-4.3-10-1.3 M) and pH (4.1-8.6) using the three CdS(s) phases; the experimental data are listed in Table 3. To formulate the chemical model that can reproduce these
TABLE 2. Comparison of CdS(s) Solubility Productsa Obtained in This Study ((95% Confidence Interval) with Values Reported by Othersb CdS(s) crystalline CdS greenockite
preparation Aldrich, 0.2-5 µm synthesized from Cd and S calculated calculated calculated precipitated Cd(NO3)2 + thioacetamide Cd(NO3)2 + Na2S CdSO4 + Na2S CdCl2 + H2S
log Ksp -14.82 ( 0.03 -14.36 ( 0.26 -14.92 ( 0.19 -15.14 ( 0.19 -14.04 ( 0.28 -14.08 -15.93 -14.40 ( 0.06 -14.15 ( 0.06 -12.28e -12.3
reference this study 9 recalculated from ref 9c recalculated from ref 9d 16 17 18 this study this study 6 19
a The solubility products correspond to eq 2. b Unless indicated, values are at 25 °C and I ) 0. c Assuming that K in ref 9 is not a mixed constant 1 (see text). d Assuming that K1 in ref 9 is a mixed constant (see text). e Value at I ) 1 M.
TABLE 3. Solubility of CdS(s) in Bisulfide Solutions (T ) 25 °C, M) CdS(s)I
CdS(s)II
CdS(s)III
pH
I
logΣ[S(-II)]
logΣ[Cd]
6.41 7.30 7.95 8.34 6.25 7.43 7.45 8.31 4.13 4.14 4.15 4.15 4.16 4.17 4.18 4.18 6.35 7.21 7.31 7.40 7.40 7.43 7.44 7.45 7.45 7.45 7.45 7.48 7.55 8.38 8.40 8.40 8.42 8.43 8.43 8.55
0.063 0.063 0.063 0.063 0.063 0.063 0.021 0.063 0.011 0.011 0.011 0.012 0.012 0.011 0.013 0.016 0.063 0.021 0.063 0.014 0.021 0.012 0.012 0.012 0.012 0.018 0.019 0.016 0.022 0.021 0.019 0.063 0.037 0.016 0.063 0.022
-1.28 -1.28 -1.28 -1.28 -1.28 -1.28 -1.98 -1.28 -3.98 -3.60 -4.28 -3.29 -3.11 -4.10 -2.99 -2.29 -1.28 -1.98 -1.28 -2.59 -1.99 -3.11 -2.96 -3.28 -2.98 -2.08 -2.11 -2.28 -1.96 -1.99 -2.11 -1.28 -1.58 -2.28 -1.28 -1.96
-8.22 -7.70 -8.31 -8.60 -8.10 -7.53 -9.56 -8.07 -10.40 -10.21 -10.50 -9.92 -9.78 -10.51 -9.75 -9.02 -7.72 -9.30 -7.17 -9.75 -8.75 -10.45 -10.10 -10.58 -10.27 -9.59 -9.05 -9.38 -8.98 -9.52 -9.90 -7.79 -8.63 -10.12 -8.01 -9.30
experimental data, we considered the following complexation reactions (9):
Cd2+ + nHS- ) Cd(HS)n2-n 2+
Cd
-
Kn
2-r-q
+ qHS + rH2O ) Cd(OH)r(HS)q
(4)
+ rH
+
∑[Cd] ) [Cd ] + ∑[Cd(HS) ] + ∑[Cd(OH) (HS) ] + ∑[Cd(OH) ] ) f(K , K , K , *K , pH, ∑S(-II), γ) 2-n
2+
n
2-r-q
r
2-m
q
sp
m
n
rq
m
Since the ionic strength was always less than 0.1 M, the activity coefficients, γ, of the chemical species were calculated with the Davies equation. The hydrolysis constants, *Km (m ) 1-4), were taken from NIST (14), and the solubility products Ksp were those determined in this study for the three CdS(s) solids (Table 2). The values of the formation constants Kn and Krq were adjusted using a stepwise regression method (SPSS) that minimizes the sum of squares of deviations in ∑[Cd] between modeled and experimental data. In a first attempt, the optimization program was applied to all our data simultaneously. With this approach, the two complexes Cd(HS)3- (log K3 ) 16.26 ( 0.58) and Cd(HS)42(log K4 ) 18.43 ( 0.05), which tend to be dominant at high ∑S(-II) and high pH, were sufficient to explain the experimental data and other complexes that would be expected to form at lower ∑S(-II) and lower pH were statistically nonsignificant. Table 3 indicates that ∑[Cd] in our experiments varied by 3 orders of magnitude. In such a case, the least-squares calculation is dominated by the high ∑[Cd] measurements, which in this study corresponded to high ∑S(-II) concentrations and high pH values (Table 3), and thus to predominance of the complexes Cd(HS)3- and Cd(HS)42-. This large influence of high ∑[Cd] values could prevent extraction of complexes that do not exert a strong influence in the fitting when all the data are considered simultaneously. To minimize this problem, we considered in a second step a data subset (obtained with the solid CdS(s)III, see Figure 2) of similar low pH values (ranging from 4.13 to 4.18) but varying ∑S(-II). The regression analysis of these data subset yielded the two complexes CdHS+ (log K1 ) 7.38 ( 0.68) and Cd(HS)2 (log K2 ) 14.43 ( 0.01); adding other complexes did not improve the quality of the fit for this subset of data. That Cd speciation for these data set is dominated by these two complexes is consistent with the slope slightly lower than 1 obtained when log ∑[Cd] is plotted as a function of log [HS-] (Figure 2). Indeed, the equilibrium constant corresponding to eq 4 is
Kn ) {Cd(HS)n2-n}/{Cd2+}{HS-}n Krq (5)
(7)
(8)
By combining eqs 2 and 8 and taking the log, one obtains
(6)
log {Cd(HS)n2-n} ) (n - 1) log {HS} + pH + log Ksp + log Kn (9)
A multiple-regression analysis was used to identify the Cdbisulfide complexes that best explain our experimental data:
which indicates that a plot of log ∑[Cd] as a function of log [HS-] should yield slopes of 0 and 1 when Cd speciation is
2+
Cd
+ mH2O ) Cd(OH)m
2-m
+ mH
+
*Km
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FIGURE 2. Solubility of CdS(s)III in bisulfide solutions: plots of log ∑[Cd] as function of log [HS-] for three data subsets of nearly constant pH. dominated by the complexes CdHS+ and Cd(HS)2, respectively. The regression analysis of two other data subsets (pH 7.21-7.48 and pH 8.31-8.55) yielded the complexes Cd(HS)3and Cd(HS)42-, in agreement with the slopes between 2 and 3 shown in Figure 2. Figure 3A shows that the model involving the four bisulfide complexes of Cd explains well our measured ∑[Cd] whatever the CdS(s) form used in the solubility experiment. The stoichiometry and formation constants obtained in this study for Cd complexes in bisulfide solutions are compared in Table 4 with those reported in previous studies. The complexes Cd(HS)n2-n have been consistently reported in studies of Cd-bisulfide solutions (6-9). The values that we obtained for the formation constants of these complexes agree reasonably well with those [Cd(HS)n2-n; n ) 1-4] reported by Ste-Marie et al. (6). This agreement is surprising and probably fictive, given that their stability constants were determined at I ) 1 M and that the model they used to interpret their solubility data involved an unrealistically high stability constant for the formation of CdOH+ (1017.76) as compared to the value commonly accepted (103.9; ref 14). Jacobs and Emerson (21) expressed doubts about the importance of the complex CdOH+ assumed by Ste-Marie et al. (6); they reinterpreted their experimental data and reported that the species CdOHSH, CdHS2-, and CdS22- can account for the solubility of CdS(s) reported by Ste-Marie et al. (6) in the alkaline region. All three species are however inconsistent with our experimental data as they are also with those of Daskalakis and Helz (9). The values reported by Zhang and Millero (7) are conditional stability constants for CdHS+ and Cd(HS)2 in seawater; they are lower than ours (Table 4), but this result was expected mainly because they did not take into account side reactions of Cd in seawater (7, 8). The conditional constants obtained by voltammetric techniques in seawater vary depending on the method used (direct titration of sulfide by Cd; ligand competition) even when the constants have been corrected for side reactions of Cd (see Table 4; refs 7 and 8). The set of complex stoichiometries and stability constants that we obtain agrees reasonably well with those of the bisulfide complexes proposed by Daskalakis and Helz (9) if we combine the recalculated value of their Ksp for greenockite and their stability constants for Cd-bisulfide complexes given in their Table 3, which were expressed as a function of CdS(s) (Table 4). The experimental conditions used by Daskalakis and Helz (9) did not allow them to determine stability constants for the CdHS+ and Cd(HS)2 species, and they reported upper limits to the constants of these species. 4274
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FIGURE 3. Comparison between total dissolved Cd concentrations measured (∑[Cd]measured) in bisulfide solutions in equilibrium with CdS(s) and those calculated with models (∑[Cd]calculated). (Panel A) Our Cd experimental data and our stability constants and Ksp values given in Tables 2 and 4. (Panel B) Our experimental data and Daskalakis and Helz (9) model with the stability constants that we recalculated from ref 9 [i.e., species Cd(HS)n shown in Table 4 and species CdOHS-] and our Ksp values given in Table 2 for our various CdS(s) solids. (Panel C) Daskalakis and Helz (9) experimental data and our stability constants given in Table 4 and a log Ksp value of -14.92. These complexes, particularly CdHS+, are expected to predominate at low pH and low ∑S(-II) values. It is probable that their experimental conditions did not encompass low enough ∑S(-II) concentrations to capture these species. The agreement for the stability constant of Cd(HS)42- between the two studies is excellent. Their K value for the formation of Cd(HS)3- is larger than ours and the upper limit for Cd(HS)2 is lower; since they were unable to resolve Cd(HS)2, it is possible that they attributed some solubility due to the latter species instead to Cd(HS)3-. The main difference between our results and those of Daskalakis and Helz (9) arises from the inclusion of the mixed ligand complex CdOHS- in their model. This complex is inconsistent with our measurements, under our experimental conditions; introducing this species into our model did not improve any of the fits of our solubility data. Moreover, Figure 3B shows that using the set of constants that we recalculated from the data of Daskalakis and Helz (9) does not reproduce well our solubility data. To obtain log ∑[Cd]calculated shown in Figure 3B, we included the stability constant for the species CdOHS- and the upper limits for K values of the complexes Cd(HS)+ and Cd(HS)2; neglecting these latter two species or
TABLE 4. Comparison of Stoichiometry and Log K Values ((95% Confidence Interval) Obtained in This Study for Cd Complexes with Bisulfide Ligand with Those Obtained in Other Studies log K (I ) 0; T ) 25 °C) reaction Cd2+
+
HS-
)
CdHS+
Cd2+ + 2HS- ) Cd(HS)2
this study 7.38 ( 0.68
14.43 ( 0.01
Cd2+ + 3HS- ) Cd(HS)3-
16.26 ( 0.58
Cd2+ + 4HS- ) Cd(HS)42-
18.43 ( 0.05
other sources 7.55 ( e7.66 e8.22
0.16a
6.3 ( 0.2c 7.84d 9.13 ( 0.02e 7.63-8.43 f 14.61 ( 0.16a e13.36 e13.92 12.7 ( 0.2c 14.24d 14.99-15.69f 16.49 ( 0.20a 16.44 ( 0.36 17.00 18.85 ( 0.20a 17.89 ( 0.01 18.45
refs
6 9 recalculated from ref 9b 7 8 8 8 6 9 recalculated from ref 9b 7 8 8 6 9 recalculated from ref 9b 6 9 recalculated from ref 9b
a Values at I ) 1.0 M. b Recalculated from ref 9 using log K c sp ) -14.92. Conditional constants for seawater uncorrected for side reactions. Conditional constant from ref 7 corrected by ref 8 for side reactions of Cd. e Conditional constant obtained by titration of sulfide with Cd and corrected for side reactions of Cd. f Conditional constants obtained by the ligand competition method and corrected for side reactions of Cd. d
using the original set of stability constant given by Daskalakis and Helz (9) did not improve the fitting of our experimental data. The set of complexes and the corresponding equilibrium constants derived from Daskalakis and Helz (9) solubility data strongly overestimate some of our experimental data (Figure 3B). The large overestimations (upper left part of Figure 3B) are for two of our data subsets of similar pH (pH ∼8.4 and ∼7.4, respectively) but varying ∑S(-II) (from 10-3.3 to 10-1.6 M). Under these experimental conditions, the Daskalakis and Helz (9) model predicts that CdOHS- should dominate dissolved Cd speciation (i.e., in most cases, CdOHSshould represent >98% of total dissolved Cd). Equilibrium with CdS(s) indicates that the concentration of CdOHSshould be independent of ∑S(-II) at a given pH, which shows in Figure 3B as two flat series of points at predicted log ∑[Cd] of -8.6 and -7.6, respectively. Thus, these overestimations suggest that the species CdOHS- does not exist; a similar conclusion has been reached recently in a study of CdS(s) solubility at high pH (22). Figure 3C shows that our model explains reasonably well the solubility data of Daskalakis and Helz (9), with the exception of five runs for which ∑[Cd] is largely underpredicted. These latter solubility data, for an unknown reason, are incompatible with ours; their only evident particularity is that they correspond to the lowest ∑S(-II) concentrations (10-3.43-10-2.24 M) and ionic strength used by Daskalakis and Helz (9) in their experiments. Relevance to Natural Waters. We examine below the relative importance of the bisulfide ligand in complexing Cd as compared to other ligands such as humic substances [the sum of humic (HA) and fulvic (FA) acids] and organic thiols that can also be present in sulfidic natural waters. To compare the relative importance of the various Cd complexes, we use the computer code WHAM (Windermere Humic Aqueous Model; ref 23). We assume the following concentrations of dissolved components as input to WHAM: pH ) 6.5, [Cd] ) 1 nM, [Ca] ) 150 µM, [Mg] ) 50 µM, [Na] ) 100 µM, [Fe] ) 250 µM, [inorganic carbon] ) 100 µM, [cysteine] ) 10 µM, and [humic substances] ) 24 mg L-1. The concentration of humic substances is based on 12 mg L-1 dissolved organic
carbon and the assumption that 50% of dissolved organic matter is carbon (24) and that all organic matter is humic substances. For the speciation calculation, we also assume that the ratio of FA:HA is 9 (23). The concentrations chosen for the inorganic components and the dissolved organic carbon have been reported for anoxic porewaters of an oligotrophic lake (3). The concentration of cysteine is based on maximum values of low molecular weight hydrophilic thiols reported in sulfidic porewaters (25, 26); cysteine is chosen as a model compound because it is among the most abundant thiols reported in natural waters (27) and because relevant acidity and stability constants are available for this compound (ref 28; Table A in the Supporting Information). The thermodynamic database of WHAM was modified to include our equilibrium constants for the Cd-bisulfide complexes (Table 4) and the acidity and stability constants of cysteine (Table A in the Supporting Information). Figure 4 shows that Cd-bisulfide complexes should largely dominate Cd speciation in sulfidic waters at ∑S(-II) g 10-6 M; humic substances and cysteine complexes cannot compete effectively with bisulfide for Cd. For example, unrealistically high concentrations of organic ligands (∼5 × 10-3 M cysteine and ∼200 mg L-1 humic substances) would be required to overcome bisulfide complexation of Cd at ∑S(-II) ) 10-5 M. Concentrations of ∑S(-II) g 10-6 M are commonly found in lacustrine (3, 29) and marine porewaters (30, 31), in the hypolimnion of lakes (32-34), and in marine anoxic basins (35, 36). The speciation of Cd in the Cd-sulfide system is shown in Figure 5 as a function of pH at two constant values of ∑S(-II) concentrations chosen to represent freshwater (10-5 M) and marine (10-3 M) environments, respectively. Figure 5A shows that our model predicts that Cd(HS)2 is the dominant species for ∑S(-II) ) 10-5 M in the pH range (pH 6-8) normally found in sulfidic freshwaters; it would predict that, in the same pH range, CdHS+ would predominate for ∑S(-II) between 10-7 and 10-6 M and Cd2+ for ∑S(-II) below 10-7 M. Figure 5B illustrates that Cd(HS)2 is the most abundant species at ∑S(-II) ) 10-3 M, according to our model; it would VOL. 33, NO. 23, 1999 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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Until there is good agreement between constants and stoichiometries obtained by solubility measurements with those obtained from other methods, they should be treated with caution.
Acknowledgments Financial support from the Natural Sciences and Engineering Research Council of Canada, the Que´bec Fonds pour la Formation de Chercheurs et l’Aide a` la Recherche, and the U.S. Environmental Protection Agency is acknowledged. We thank K. D. Daskalakis, L. Hare, G. R. Helz, J. R. Kramer, and an anonymous reviewer for their critical comments on the manuscript. F.W. was supported by a postdoctoral fellowship from the Institut National de la Recherche Scientifique.
Supporting Information Available FIGURE 4. Distribution of various forms of dissolved Cd as a function of ∑S(-II) for assumed concentrations of 24 mg L-1 humic substances and 10-5 M cysteine. Cd-CYS and Cd-humics are abbreviations for complexes of Cd with cysteine and humic substances, respectively.
A table listing the equilibrium constants and a figure showing the X-ray diffraction patterns (2 pages). This material is available free of charge via the Internet at http://pubs.acs.org.
Literature Cited
FIGURE 5. Distribution of Cd species as a function of pH predicted by our model at ∑S(-II) of (A) 10-5 and (B) 10-3 M. predict that Cd(HS)3- would become the dominant species at ∑S(-II) between 10-2 and 10-1 M and that it is only above this ∑S(-II) concentration that Cd(HS)42- would be the most abundant species. Such high sulfide concentrations are rarely found in natural waters. In conclusion, we have shown by solubility measurements that the solubility of CdS(s) in bisulfide solutions can be described by four complexes of the form Cd(HS)n2-n (n ) 1-4). It would be desirable that complementary work use different techniques over a wide range of solution conditions embracing those of natural waters to unravel more fully the Cd-bisulfide system. Indeed, complexes that best reproduce solubility data are not necessarily the actual ones present; such complexes only reproduce measured solubilities in chemical systems restricted to the experimental conditions. 4276
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Received for review March 12, 1999. Revised manuscript received August 26, 1999. Accepted August 30, 1999. ES990283Z
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