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Arsenic(III) carbonate complexing
Applied Geochemistry Applied Geochemistry 20 (2005) 1218–1225 www.elsevier.com/locate/apgeochem
Arsenic(III) carbonate complexing Carla S. Neuberger 1, George R. Helz
*
Department of Chemistry and Biochemistry, Water Resources Research Center, University of Maryland, College Park, MD 20742, USA Received 8 September 2004; accepted 6 January 2005 Editorial handling by R. Fuge Available online 19 March 2005
Abstract A recent hypothesis that As(III) forms strong complexes with (bi)carbonate is tested by measuring the solubility of As2O3 in concentrated (up to 0.72 m) bicarbonate solutions at near-neutral pH and 25.0 ± 1.8 °C. A small, but statistically significant solubility enhancement is observed in NaHCO3-containing solutions compared to NaCl solutions of essentially the same ionic strength. The effect is well-explained by formation of a single complex: AsðOHÞ03 þ HCO 3 ¼ AsðOHÞ2 CO3 þ H2 O, Kc = 0.22 ± 0.04 (corrected to infinite dilution). This complexÕs small stability constant value is consistent with a published, quantum-based computational prediction. The measurements also yield a value for the solubility of arsenolite that agrees reasonably with recently reported values: 1=2As2 O3 ðarsenoliteÞ þ 3=2 H2 O ¼ AsðOHÞ03 , Ks = 0.178 ± 0.004 (also corrected to infinite dilution). The new data show that As(III) carbonate complexes will be negligible at carbonate concentrations found in most natural waters, but could be significant, though minor, in extremely carbonate-rich waters, such as found in interior drainage basins. Ó 2005 Elsevier Ltd. All rights reserved.
1. Introduction Kim et al. (2000) and Lee and Nriagu (2003) have proposed a novel hypothesis that dissolved As(III) carbonato complexes (As(CO3)+, AsðCO3 Þ and 2, AsðOHÞ2 CO 3 ) might be ‘‘the most stable inorganic arsenic species in the aquatic environment.’’ Among the arguments offered to support this view are: (a) high concentrations of HCO 3 (0.02–0.6 M) promote leaching of As from aquifer materials (see also Anawar et al., 2004), (b) ion exchange chromatograms of solutions containing CO2-saturated arsenite display anom-
*
Corresponding author. E-mail address: [email protected] (G.R. Helz). 1 Current address: Bettis Atomic Power Laboratories, West Miflin, PA 15122-0079, USA.
alous features compared to solutions containing only arsenite, (c) extrapolation of stability constants for lanthanide carbonato complexes vs. ionic radius suggests that carbonato complexes of As3+ would have large stability constants and (d) formation of extremely stable As(III) carbonato complexes can explain As(III) leaching from sulfide minerals in anaerobic aquifers. Several of these arguments can be challenged. The last argument has become immaterial as evidence has mounted that processes other than leaching of sulfide minerals generate most of the dissolved As in anaerobic aquifers (Appelo et al., 2002; Harvey et al., 2002; Smith et al., 2003; Stu¨ben et al., 2003; McArthur et al., 2004; Islam et al., 2004; Szramek et al., 2004; Horneman et al., 2004). Kim et al. (2000) did not fully exclude the possibility that HCO 3 releases As from aquifer materials through its capacity to displace adsorbed As by
0883-2927/$ - see front matter Ó 2005 Elsevier Ltd. All rights reserved. doi:10.1016/j.apgeochem.2005.01.007
C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225
ion-exchange or through its capacity to buffer pH in the alkaline range. Kim et al. (2000) neglected hydrolysis of As3+; Tossell (2004) has used quantum-chemical calculations to show that the As(III) carbonato complexes predicted from the lanthanide correlation would be very unstable relative to AsðOHÞ03 . Nonetheless, the possible importance of As(III) carbonato complexes in natural waters remains to be actually quantified. Here, HCO 3 is shown indeed to enhance the solubility of As2O3 at near-neutral pH in a manner consistent with carbonate complexing, but the magnitude of the effect is small and observable only at high HCO 3 concentrations.
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2. Methodology Arsenic trioxide (J.T. Baker Analyzed, Primary Redox Standard, certified 99.95–100.05% As2O3) was recrystallized in water by cycling 4 times between room temperature and boiling. At the end of this treatment, the solids were left overnight in the aqueous phase. Then, most of the solution, with any suspended particles, was decanted, and the solids were filter-washed with water and air-dried. The product was a coarse powder. Two batches were prepared. X-ray diffraction evidence shows that one batch was a mixture of arsenolite and claudetite (upper panel of Fig. 1), whereas the other
Intensity
500 400 300 200 100 x10^3 Arsenolite - As2O3 Claudetite - As2O3 Claudet
10
20
30
40
50
60
70
80
90
2-Theta(˚) 12.0
Intensity
10.0 8.0 6.0 4.0 2.0 x10^3 ^ Arsenolite - As2O3
20
30
40
50 60 2-Theta(˚)
70
80
90
Fig. 1. X-ray powder diffraction patterns of two batches of As2O3 (CuKa) after recrystallization. The upper panel is for a batch that contains mostly arsenolite, but appears also to contain some claudetite (monoclinic As2O3), as shown by the doublet at 28° 2h. This batch was used in Runs 1 and 2. The lower panel is for a batch that contains only arsenolite (cubic As2O3). This batch was used in Runs 3 and 4.
C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225
batch consisted only of arsenolite (lower panel of Fig. 1). The reason for the difference in the batches is not known; both were prepared from the same bottle of As2O3 reagent. As suggested later, the arsenolite + claudetite batch may have also contained some amorphous As2O3, which would not have been apparent by X-ray diffractometry. Solubility experiments were conducted in 500 mL Pyrex Erlenmeyer side-arm flasks into which were placed 40–80 g of recrystallized As2O3 and 250 mL of electrolyte. The electrolyte in Runs 1and 3 was 0.710 m NaCl; in Run 2 it was 0.355 m NaCl + 0.357 m NaHCO3 and in Run 4 it was 0.715 m NaHCO3. These solutions were prepared with distilled, deionized water and J.T. Baker Analyzed reagent salts. During runs, the flasks were continuously stirred with magnetic stir bars. Even though oxidation of As(III) by O2 at near-neutral pH is believed to be very slow (Cherry et al., 1979; Eary and Schramke, 1990; Kim and Nriagu, 2000), flasks were constantly swept by a gas to eliminate air. For the first 30 min, the gas was bubbled through the solution, and subsequently it was blown over the surface. In Runs 1 and 3 (NaCl electrolyte) the gas was N2 and in Runs 2 and 4 (containing HCO 3 ) the gas was CO2. Sweeping the HCO -containing runs with CO2 had the additional 3 advantage of fixing the pH in the near-neutral range. All experiments were performed at room temperature, which was recorded frequently and never fell outside the range 23.8–26.8 °C. Periodically during an experiment, 3–5 mL aliquots were extracted from the flasks and centrifuged in a small, bench-top centrifuge for 10 min; this produced a supernate that was free of visible particles. A measured volume of this supernate was then sampled with an automatic pipette. We used centrifugation rather than filtration because we thought As(OH)3 might be adsorbed on filters. However, in some experiments, filtration (0.02 lm Whatman Anatop 25) was used on duplicate samples and compared to centrifugation; differences between the two methods of removing particles were not systematic and in most cases gave similar results. Arsenic (III) was determined by an iodometric backtitration method with amperometric endpoint detection employing an automatic titrator (Brinkmann 760 DMS Titrino). A titration jar was filled with 100 mL of deionized water, 4 mL of 0.1 M acetic acid/acetate buffer (pH 4.0) and the measured volume of centrifuged supernate. Measured aliquots (0.100 or 1.00 mL) of 0.05 or 0.005 M I2 were added to the jar until the color of the solution remained yellow, indicating accumulation of excess I2. Iodine reacts with As(III) quantitatively as follows: þ H2 O þ I2 þ AsIII ðOHÞ03 ! AsV O2 ðOHÞ 2 þ 2I þ 3H
ð1Þ
Subsequently, the excess I2 was titrated with phenylarsine oxide solution that had been standardized with potassium bi-iodate. The moles of As(III) in the sample are equal to the difference between the moles of titrant needed to titrate the sample and the moles needed to titrate a volume of the I2 reagent solution equal to the volume added to the sample. The reaction of phenylarsine oxide with I2 is analogous to reaction 1: 2H2 O þ I2 þ C6 H5 AsIII O ! C6 H5 AsV OðOHÞ2 þ 2I þ 2Hþ ð2Þ
The pH in the reactor was measured at the start of each sampling with a glass electrode calibrated at pH 7 and 10 with commercial buffers. The precision of pH measurements was approximately ±0.02 and the reproducibility of replicate As determinations was about ±0.01 m.
3. Results Fig. 2 shows results from an experiment to determine the time necessary to reach reversible equilibrium. A solution was prepared by saturating 0.351 m Na2CO3 with As2O3 under N2; the pH was about 11.5. A day later, the N2 gas was switched to CO2 gas (1 atm). This quickly converted each mole of CO2 3 to two moles of 2 HCO driving the 3 ðCO2 þ CO3 þ H2 O ! 2HCO3 Þ, pH down to about 7.7 and forcing As(III) to start precipitating. At the same time, another solution containing 0.715 m NaHCO3 under CO2 gas was allowed to begin dissolving As2O3. By the end of one day, the As(III)
0.4
0.3 As(III) (m)
1220
0.2
0.1 From supersaturation From undersaturation
0.0 0
20
40 Time (h)
60
80
Fig. 2. Reversibility test. Data represented by circles are for a run that contained 0.351 m Na2CO3, pre-equilibrated with As2O3 for 24 h at pH 11.5 under N2. At t = 0, CO2 gas was introduced, converting Na2CO3 to 0.715 m NaHCO3 and driving the pH down to 7.7. Data represented by triangles are for a run that contained 0.715 m NaHCO3 at pH 7.7 under CO2 from the beginning.
C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225 Table 1 Experimental data
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Table 1 (continued)
Total As (m)
pH
Reaction time (h)
Run 1 0.250 0.239 0.254 0.243 0.236 0.247 0.247 0.257 0.255 0.246 0.253 0.242 0.231 0.230 0.231 0.237 0.220 0.226 0.252 0.231 0.237 0.260 0.262 0.260 0.244 ± 0.011
6.97 6.97 6.97 6.94 6.94 6.94 6.96 6.96 6.96 6.96 6.96 6.96 7.02 7.02 7.02 7.02 7.02 7.02 6.90 6.90 6.90 6.90 6.90 6.90 Mean ± SD
49 49 49 74 74 74 97 97 97 97 97 97 167 167 167 167 167 167 215 215 215 215 215 215
Run 2 0.253 0.239 0.254 0.251 0.260 0.253 0.264 0.267 0.269 0.269 0.272 0.276 0.249 0.254 0.249 0.247 0.245 0.219 0.266 0.264 0.261 0.269 0.260 0.270 0.258 ± 0.013
7.09 7.09 7.09 7.11 7.11 7.11 7.11 7.11 7.11 7.11 7.11 7.11 7.12 7.12 7.12 7.12 7.12 7.12 7.15 7.15 7.15 7.15 7.15 7.15 Mean ± SD
49 49 49 74 74 74 97 97 97 97 97 97 167 167 167 167 167 167 215 215 215 215 215 215
Run 3 0.181 0.160 0.169
5.54 5.54 5.54
48 48 48
Total As (m)
pH
Reaction time (h)
0.186 0.188 0.173 0.139 0.152 0.152 0.147 0.147 0.143 0.154 0.152 0.151 0.160 ± 0.016
5.64 5.64 5.64 5.54 5.54 5.54 5.67 5.67 5.67 5.67 5.67 5.67 Mean ± SD
73 73 73 144 144 144 218 218 218 218 218 218
Run 4 0.201 0.201 0.190 0.199 0.204 0.211 0.188 0.176 0.184 0.221 0.211 0.240 0.232 0.231 0.220 0.207 ± 0.019
7.74 7.74 7.74 8.03 8.03 8.03 7.97 7.97 7.97 7.71 7.71 7.71 7.71 7.71 7.71 Mean ± SD
48 48 48 73 73 73 144 144 144 218 218 218 218 218 218
Run 1: 0.710 m NaCl, arsenolite + claudetite, pH adjusted with NaOH at start. Run 2: 0.355 m NaCl, 0.357 m NaHCO3, arsenolite + claudetite. Run 3: 0.710 m NaCl, arsenolite only. Run 4: 0.715 m NaHCO3, arsenolite only.
solubility of these two runs was nearly the same (Fig. 2) and after a second day the solubilities had reached a final plateau. For all runs, solubilities did not change systematically with time after 48 h of equilibration. Note that the run that began from undersaturation appears to overshoot the equilibrium point slightly at 24 h. This presumably occurs because a few particles possess excess free energy despite the efforts to prevent this by recrystallizing the starting material; after these particles dissolve, the system equilibrates with less soluble particles. Table 1 presents the solubility data that will be used to explore the stability of As(III) carbonato complexes. Two features of the data are noteworthy. First, the solubility differs for the different As2O3 batches. The same electrolyte (0.710 m NaCl) was used in both Run 1 (arsenolite + claudetite) and Run 3 (arsenolite only), but the average solubility of the first batch was 0.244 (r 0.011) m whereas the average solubility of the third was 0.160 (r 0.016) m. The solubility of the arsenolite + claudetite
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batch remained metastably high throughout the weeklong experiments and never drifted downward toward that of the arsenolite-only batch (Table 1). The second noteworthy feature in Table 1 is that, for a given As2O3 batch, the As(III) solubility in HCO 3containing solutions is slightly, but distinctly greater than in solutions containing no HCO 3 , even though the solutions are compared at essentially the same ionic strength. For example, in HCO 3 -containing Run 2, average As(III) is 0.258 (r 0.013) m vs. 0.244 (r 0.011) m in HCO 3 -free Run 1. These means are significantly different at the 99% level of confidence. In Run 4, where HCO 3 is twice that in Run 2, the contrast relative to HCO -free Run 3 is even larger: 0.207 3 (r 0.019) m for Run 4 vs. 0.160 (r 0.016) m for Run 3. Thus HCO 3 appears to enhance As(III) solubility under the conditions of these experiments.
4. Discussion Because the enhancement of As(III) solubility by HCO 3 is weak, it was observable only at high ionic strength. Experimental results obtained at high ionic strength can be modeled various ways. Here, a simple ion association approach is adopted, and the extra solubility observed in HCO 3 -containing solutions is attributed to formation of a single carbonato complex. The choice of approach has been influenced by the previous work of Kim et al. (2000), Lee and Nriagu (2003) and Tossell (2004), all of whom describe the role of HCO 3 as that of a complexing ligand. As described below, the extra As(III) solubility observed in the presence of HCO 3 can be explained well by invoking the single complex, AsðOHÞ2 CO 3 . Tossell (2004) concluded that this complex would have modest stability in water, similar to the stability of As(III) oxyhydroxide dimers. He also found that the other complexes proposed by Kim et al. (2000), As(CO3)+ and AsðCO3 Þ 2 , would be unstable with respect to hydrolysis. Neuberger (2004) examined the possibility of fitting the data in Table 1 with either of the other carbonato complexes individually. Both produced statistically slightly poorer fits to the data. Additionally, As(CO3)+ produced fits in which the deviations between predicted and observed solubilities displayed undesirable trends relative to pH and ½HCO 3 . To derive a stability constant for AsðOHÞ2 CO 3 , total dissolved As(III) in the experiments is considered to consist of five species (brackets denote molal concentrations):
½X þ ½NaX0 ¼ ½X ð1 þ K NaX ½Naþ Þ
ð4Þ
Similarly, RNa ¼ ½Naþ þ R½NaX0 ¼ ½Naþ ð1 þ RðK NaX ½X ÞÞ
ð5Þ
Johnson and Pytkowicz (1979) estimate that stoichiometric formation constants for NaCl0 and NaHCO03 ion-pairs in seawater (ionic strength 0.7 m) are KNaX = 0.3. Lacking measured formation constants for NaAsOðOHÞ02 and NaAsðOHÞ2 CO03 , it is assumed that KNaX = 0.3 describes the stability of these ion-pairs as well. This assumption allows equating R(KNaX [X]) in Eq. (5) to 0.3R[X], which owing to the charge balance constraint must equal 0.3[Na+]. Substituting this last term and solving Eq. (5) at RNa = 0.71 m gives [Na+] = 0.601 m in the solutions. Using this value in Eq. (4) gives [X] + [NaX0] = 1.18[X]; thus Eq. (3) can be rewritten, RAsðIIIÞ ¼ ½AsðOHÞ03 þ 1:18½AsOðOHÞ 2 þ 1:18½AsðOHÞ2 CO 3
ð6Þ
Equilibrium concentrations in Eq. (6) are subject to the following mass action laws: 1=2As2 O3 ðsÞ þ 3=2H2 O $ AsðOHÞ03 ; K s ¼
c0 ½AsðOHÞ03 a0:5 s ð7Þ
þ AsðOHÞ03 $ AsOðOHÞ 2 þH ;
Ka ¼
pH c1 ½AsOðOHÞ 2 ½10
c0 ½AsðOHÞ03
AsðOHÞ03 þ HCO 3 $ AsðOHÞ2 CO3 þ H2 O; ½AsðOHÞ2 CO3 Kc ¼ c0 ½AsðOHÞ03 ½HCO 3
ð8Þ
ð9Þ
where c0 and c1 are activity coefficients for neutral and univalent dissolved species and as is the activity of the As2O3 component in the solid phase. Note that in Eq. (9) it has been assumed that the activity coefficients of HCO 3 and AsðOHÞ2 CO3 cancel; any systematic error introduced by this assumption will be embedded quantitatively in the Kc value obtained below. Substituting Eqs. (7)–(9) into Eq. (6) yields: a0:5 K s 1:18K a c0 1þ þ 1:18K c c0 ½HCO3 : RAsðIIIÞ ¼ c0 c1 10pH
RAsðIIIÞ ¼ ½AsðOHÞ03 þ ½AsOðOHÞ 2 þ ½NaAsOðOHÞ02 þ ½AsðOHÞ2 CO 3 þ ½NaAsðOHÞ2 CO03
The concentration of any Na ion-pair, [NaX0,] in the system can be related to the concentration of free anion, [X], by the following equation in which KNaX is the stoichiometric formation constant of the ion-pair in the ionic medium and [Na+] is the molal concentration of free Na+:
ð3Þ
ð10Þ
C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225
þpH 1 ½Naþ ½Cl K a K s a0:5 Þ s c1 ð10 0:5 1 þ K c K s as
0.0 -0.1 -0.2 0.0 0.1 0.2 0.3 0.4 0.5 0.6 HCO3 (m) 0.2 0.1 0.0 -0.1 -0.2 6.0
6.5
7.0 pH
7.5
8.0
ð11Þ
Note that in writing Eq. (11), it has been assumed that alkalinity, but not HCO 3 , is conserved; in principle, HCO can be non-conserved because of HCO 3 3þ AsðOHÞ3 ¼ H2 CO3 þ AsOðOHÞ 2. Substituting Eqs. (7)–(9) into 11 and rearranging yields: ½HCO 3 ¼
0.1
5.5
½Naþ ¼ ½Cl þ ½HCO 3 þ ½AsOðOHÞ2
þ ½AsðOHÞ2 CO 3
0.2
Observed - Calculated (m)
In order to use this equation to fit the data in Table 1, it is assumed that c0 = 1.09 at I = 0.7 based on measurements by Garrett et al. (1940) of the decrease in As2O3 solubility as a function of HCl concentration. In Runs 3 and 4 (arsenolite only), as is defined to be one, and a value for as is fit in Runs 1 and 2. This attributes the higher solubility of the arsenolite + claudetite batch to a higher activity of the As2O3 component. A Ka value of 109.17 was used (Nordstrom and Archer, 2003), and c1 was set to 0.68 based on the Davies equation. The final fit is not sensitive to uncertainties in c1 and Ka because these quantities affect only the computed concentrations of AsOðOHÞ 2 and its ion-pair, which are minor-to-negligible in Eq. (6) for the solutions considered here. The HCO 3 concentration in Eq. (10) is evaluated from a charge balance equation (neglecting H+ and OH):
1223
ð12Þ
Based on Eqs. (4) and (5), [Na+] and [Cl] can be replaced by RNa+/1.18 and RCl/1.18. Using these values and Eqs. 10 and 12, the entire set of data in Table 1 is fit by varying the values of Ks, Kc and as until good predictions of observed solubilities were obtained. Fig. 3 shows deviations between the calculated and observed solubilities as a function of ½HCO 3 and pH. Fitting was carried out by the non-linear least squares method, using SCIENTIST (MicroMath Scientific Software). The results are: Ks = 0.178 ± 0.004, Kc = 0.22 ± 0.04 and as (in the arsenolite + claudetite batch) = 2.12 ± 0.08. Quoted uncertainties are 1U and reflect only the statistics of fitting the data, not systematic errors such as might arise from model-related assumptions. A sensitivity analysis shows that changing either c1 or KNaX by ±25% causes the derived constants to change by less than the statistical uncertainty. Thus evaluation of the constants is insensitive to assumptions about these two parameters. The value of Ks for arsenolite is slightly higher than a value of 0.15 ± 0.02 determined for arsenolite by Pokrovski et al. (1996), but lower than the value of 0.20 selected by Nordstrom and Archer (2003) from a survey of literature. There are no previous experimental values of Kc to compare with the one obtained here, but TossellÕs
Fig. 3. Comparison of observed dissolved As(III) to that calculated from the 5-species model represented by Eq. (3) in text. The absence of trends vs. HCO 3 (upper panel) and pH (lower panel) indicates that the model accounts well for the variation of solubility in the data set. Data obtained only afterP48 h equilibration are plotted.
(2004) prediction that AsðOHÞ2 CO 3 would prove to be similar in stability to the dimer, As2 O2 ðOHÞ 3 , is supported by the present result. Writing the formation reaction for the dimer in a format analogous to that of reaction 9, the following equilibrium constant is calculated from the data of Garrett et al. (1940): 0:16 AsðOHÞ03 þ AsOðOHÞ 2 $ As2 O2 ðOHÞ3 þ H2 O; K ¼ 10
ð13Þ This value is reasonably close to the 100.66 obtained here for Kc. The Gibbs free energy of formation of As2O3 in the arsenolite + claudetite batch, based on the value of as, was 1.8 kJ/mol less negative than the corresponding quantity in the arsenolite-only batch. It is tempting to interpret this as the difference in the Gibbs free energy of formation of claudetite vs. arsenolite. However, this idea contradicts the current view that claudetite is the stable polymorph of As2O3 at 25 °C and thus must have a more negative Gibbs free energy of formation than arsenolite (Pokrovski et al., 1996; Nordstrom and Archer, 2003). The fact that the solubility of the arsenolite + claudetite batch remained perched metastably during the weeklong experiment indicates that relaxation of excess Gibbs free energy within As2O3 solids is slow. The source of excess free energy is not known. It is possible that the arsenolite + claudetite batch con-
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C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225
tained X-ray amorphous As2O3 that formed during the preliminary recrystallization process. This phase may have controlled the solubility even though crystalline material was present. The results in this paper uphold qualitatively the contention of Kim et al. (2000) and Lee and Nriagu (2003) that As(III) can be complexed by carbonate. However, based on the value of Kc, the HCO 3 concentration would need to exceed about 0.04 m in or0 der for ½AsðOHÞ2 CO 3 to exceed 1% of ½AsðOHÞ3 at near-neutral pH. For comparison, the average HCO 3 concentration is 0.002 m in the ocean and about half that in rivers. Therefore, contrary to the above authors, we conclude that As(III) carbonate complexing will be negligible in most natural waters. To emphasize this point, a speciation calculation for conditions representative of many groundwaters is presented in Fig. 4. Possibly As(III) carbonate species are significant in surface and ground waters of evaporative basins. For example, ½HCO 3 in excess of 0.1 m is known in evaporative lakes in western U.S.A. (Livingston, 1963; Eugster and Hardie, 1973). Welch and Lico (1998) describe mildly alkaline, pH 7–9 groundwaters in Nevada with up to 0.05 m HCO 3 . These also contain exceptionally high As(III) concentrations. In such unusual environments, As(III) carbonate complexes might prove to be significant but minor.
-
H2CO3
1E-3
HCO3
23 CO
[molal concentration]
1E-5
0
As(OH)3
1E-7
-
) OH O( s A
1E-9
2
-
As(OH) 2CO3 1E-11
1E-13
5
6
7
8
9
10
pH
Fig. 4. Calculated species distribution in a solution containing 103 m total carbonate and 107 m total As(III). These concentrations are representative of those commonly found in groundwaters. Activity coefficients have been set to unity in this calculation.
Acknowledgements This work supported by Grant EAR 0229387 from the National Science Foundation. The authors thank Prof. Jack Tossell for originally suggesting this problem to us and for informative subsequent discussions. The manuscript benefited from constructive criticism by J. Webster and an anonymous reviewer.
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C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225 Lee, J.S., Nriagu, J.O., 2003. Arsenic carbonate complexes in aqueous systems. In: Cai, Y., Braids, O.C. (Eds.), Biogeochemistry of Environmentally Important Trace Elements. American Chemical Society Symposium Series 835, pp. 33– 40. Livingston, D.A., 1963. Data of Geochemistry. G. Chemical Composition of Rivers and Lakes. US Geol. Surv. Prof. Pap., 440-G. US Government Printing Office, Washington, DC. McArthur, J.M., Banerjee, D.M., Hudson-Edwards, K.A., Mishra, R., Purohit, R., Ravenscroft, P., Cronin, A., Howarth, R.J., Chatterjee, A., Talukder, T., Lowry, D., Houghton, S., Chadha, D.K., 2004. Natural organic matter in sedimentary basins and its relation to arsenic in anoxic ground water: the example of West Bengal and its worldwide implications. Appl. Geochem. 19, 1255–1293. Neuberger, C., 2004. Arsenic (III) speciation in sulfidic- and carbonate-containing waters; thioarsenites as ligands for Ag(I), Pb(II) and Hg(II). Ph.D. Dissertation, Univ. Maryland, College Park. Nordstrom, K.D., Archer, D.G., 2003. Arsenic thermodynamic data and environmental geochemistry. In: Welch, A.H., Stollenwerk, K.G. (Eds.), Arsenic in Ground Water. Kluwer Academic Publishers, pp. 1–25. Pokrovski, G., Gout, R., Schott, J., Zotov, A., Harrichoury, J.-C., 1996. Thermodynamic properties and stoi-
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chiometry of NAs(III) hydroxide complexes at hydrothermal conditions. Geochim. Cosmochim. Acta 60, 737–749. Smith, J.V.S., Jankowski, J., Sammut, J., 2003. Vertical distribution of As(III) and As(V) in a coastal sandy aquifer: factors controlling the concentration and speciation of arsenic in the Stuarts Point groundwater system, northern New South Wales, Australia. Appl. Geochem. 18, 1479– 1496. Stu¨ben, D., Berner, Z., Chandrasekharam, D., Karmakar, J., 2003. Arsenic enrichment in groundwater of West Bengal, India: geochemical evidence for mobilization of As under reducing conditions. Appl. Geochem. 18, 1417– 1434. Szramek, K., Walter, L.M., McCall, P., 2004. Arsenic mobility in groundwater/surface water systems in carbonate-rich Pleistocene glacial drift aquifers (Michigan). Appl. Geochem. 19, 1137–1155. Tossell, J.A., 2004. Calculation of the interaction of bicarbonate ion with arsenites in aqueous solution and with the surfaces of Al hydroxide minerals. In: Advances in Arsenic Research. American Chemical Society of Symposium Series, in press. Welch, A.H., Lico, M.S., 1998. Factors controlling As and U in shallow ground water, southern Carson Desert, Nevada. Appl. Geochem. 13, 521–539.
Arsenic(III) carbonate complexing Carla S. Neuberger 1, George R. Helz
*
Department of Chemistry and Biochemistry, Water Resources Research Center, University of Maryland, College Park, MD 20742, USA Received 8 September 2004; accepted 6 January 2005 Editorial handling by R. Fuge Available online 19 March 2005
Abstract A recent hypothesis that As(III) forms strong complexes with (bi)carbonate is tested by measuring the solubility of As2O3 in concentrated (up to 0.72 m) bicarbonate solutions at near-neutral pH and 25.0 ± 1.8 °C. A small, but statistically significant solubility enhancement is observed in NaHCO3-containing solutions compared to NaCl solutions of essentially the same ionic strength. The effect is well-explained by formation of a single complex: AsðOHÞ03 þ HCO 3 ¼ AsðOHÞ2 CO3 þ H2 O, Kc = 0.22 ± 0.04 (corrected to infinite dilution). This complexÕs small stability constant value is consistent with a published, quantum-based computational prediction. The measurements also yield a value for the solubility of arsenolite that agrees reasonably with recently reported values: 1=2As2 O3 ðarsenoliteÞ þ 3=2 H2 O ¼ AsðOHÞ03 , Ks = 0.178 ± 0.004 (also corrected to infinite dilution). The new data show that As(III) carbonate complexes will be negligible at carbonate concentrations found in most natural waters, but could be significant, though minor, in extremely carbonate-rich waters, such as found in interior drainage basins. Ó 2005 Elsevier Ltd. All rights reserved.
1. Introduction Kim et al. (2000) and Lee and Nriagu (2003) have proposed a novel hypothesis that dissolved As(III) carbonato complexes (As(CO3)+, AsðCO3 Þ and 2, AsðOHÞ2 CO 3 ) might be ‘‘the most stable inorganic arsenic species in the aquatic environment.’’ Among the arguments offered to support this view are: (a) high concentrations of HCO 3 (0.02–0.6 M) promote leaching of As from aquifer materials (see also Anawar et al., 2004), (b) ion exchange chromatograms of solutions containing CO2-saturated arsenite display anom-
*
Corresponding author. E-mail address: [email protected] (G.R. Helz). 1 Current address: Bettis Atomic Power Laboratories, West Miflin, PA 15122-0079, USA.
alous features compared to solutions containing only arsenite, (c) extrapolation of stability constants for lanthanide carbonato complexes vs. ionic radius suggests that carbonato complexes of As3+ would have large stability constants and (d) formation of extremely stable As(III) carbonato complexes can explain As(III) leaching from sulfide minerals in anaerobic aquifers. Several of these arguments can be challenged. The last argument has become immaterial as evidence has mounted that processes other than leaching of sulfide minerals generate most of the dissolved As in anaerobic aquifers (Appelo et al., 2002; Harvey et al., 2002; Smith et al., 2003; Stu¨ben et al., 2003; McArthur et al., 2004; Islam et al., 2004; Szramek et al., 2004; Horneman et al., 2004). Kim et al. (2000) did not fully exclude the possibility that HCO 3 releases As from aquifer materials through its capacity to displace adsorbed As by
0883-2927/$ - see front matter Ó 2005 Elsevier Ltd. All rights reserved. doi:10.1016/j.apgeochem.2005.01.007
C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225
ion-exchange or through its capacity to buffer pH in the alkaline range. Kim et al. (2000) neglected hydrolysis of As3+; Tossell (2004) has used quantum-chemical calculations to show that the As(III) carbonato complexes predicted from the lanthanide correlation would be very unstable relative to AsðOHÞ03 . Nonetheless, the possible importance of As(III) carbonato complexes in natural waters remains to be actually quantified. Here, HCO 3 is shown indeed to enhance the solubility of As2O3 at near-neutral pH in a manner consistent with carbonate complexing, but the magnitude of the effect is small and observable only at high HCO 3 concentrations.
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2. Methodology Arsenic trioxide (J.T. Baker Analyzed, Primary Redox Standard, certified 99.95–100.05% As2O3) was recrystallized in water by cycling 4 times between room temperature and boiling. At the end of this treatment, the solids were left overnight in the aqueous phase. Then, most of the solution, with any suspended particles, was decanted, and the solids were filter-washed with water and air-dried. The product was a coarse powder. Two batches were prepared. X-ray diffraction evidence shows that one batch was a mixture of arsenolite and claudetite (upper panel of Fig. 1), whereas the other
Intensity
500 400 300 200 100 x10^3 Arsenolite - As2O3 Claudetite - As2O3 Claudet
10
20
30
40
50
60
70
80
90
2-Theta(˚) 12.0
Intensity
10.0 8.0 6.0 4.0 2.0 x10^3 ^ Arsenolite - As2O3
20
30
40
50 60 2-Theta(˚)
70
80
90
Fig. 1. X-ray powder diffraction patterns of two batches of As2O3 (CuKa) after recrystallization. The upper panel is for a batch that contains mostly arsenolite, but appears also to contain some claudetite (monoclinic As2O3), as shown by the doublet at 28° 2h. This batch was used in Runs 1 and 2. The lower panel is for a batch that contains only arsenolite (cubic As2O3). This batch was used in Runs 3 and 4.
C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225
batch consisted only of arsenolite (lower panel of Fig. 1). The reason for the difference in the batches is not known; both were prepared from the same bottle of As2O3 reagent. As suggested later, the arsenolite + claudetite batch may have also contained some amorphous As2O3, which would not have been apparent by X-ray diffractometry. Solubility experiments were conducted in 500 mL Pyrex Erlenmeyer side-arm flasks into which were placed 40–80 g of recrystallized As2O3 and 250 mL of electrolyte. The electrolyte in Runs 1and 3 was 0.710 m NaCl; in Run 2 it was 0.355 m NaCl + 0.357 m NaHCO3 and in Run 4 it was 0.715 m NaHCO3. These solutions were prepared with distilled, deionized water and J.T. Baker Analyzed reagent salts. During runs, the flasks were continuously stirred with magnetic stir bars. Even though oxidation of As(III) by O2 at near-neutral pH is believed to be very slow (Cherry et al., 1979; Eary and Schramke, 1990; Kim and Nriagu, 2000), flasks were constantly swept by a gas to eliminate air. For the first 30 min, the gas was bubbled through the solution, and subsequently it was blown over the surface. In Runs 1 and 3 (NaCl electrolyte) the gas was N2 and in Runs 2 and 4 (containing HCO 3 ) the gas was CO2. Sweeping the HCO -containing runs with CO2 had the additional 3 advantage of fixing the pH in the near-neutral range. All experiments were performed at room temperature, which was recorded frequently and never fell outside the range 23.8–26.8 °C. Periodically during an experiment, 3–5 mL aliquots were extracted from the flasks and centrifuged in a small, bench-top centrifuge for 10 min; this produced a supernate that was free of visible particles. A measured volume of this supernate was then sampled with an automatic pipette. We used centrifugation rather than filtration because we thought As(OH)3 might be adsorbed on filters. However, in some experiments, filtration (0.02 lm Whatman Anatop 25) was used on duplicate samples and compared to centrifugation; differences between the two methods of removing particles were not systematic and in most cases gave similar results. Arsenic (III) was determined by an iodometric backtitration method with amperometric endpoint detection employing an automatic titrator (Brinkmann 760 DMS Titrino). A titration jar was filled with 100 mL of deionized water, 4 mL of 0.1 M acetic acid/acetate buffer (pH 4.0) and the measured volume of centrifuged supernate. Measured aliquots (0.100 or 1.00 mL) of 0.05 or 0.005 M I2 were added to the jar until the color of the solution remained yellow, indicating accumulation of excess I2. Iodine reacts with As(III) quantitatively as follows: þ H2 O þ I2 þ AsIII ðOHÞ03 ! AsV O2 ðOHÞ 2 þ 2I þ 3H
ð1Þ
Subsequently, the excess I2 was titrated with phenylarsine oxide solution that had been standardized with potassium bi-iodate. The moles of As(III) in the sample are equal to the difference between the moles of titrant needed to titrate the sample and the moles needed to titrate a volume of the I2 reagent solution equal to the volume added to the sample. The reaction of phenylarsine oxide with I2 is analogous to reaction 1: 2H2 O þ I2 þ C6 H5 AsIII O ! C6 H5 AsV OðOHÞ2 þ 2I þ 2Hþ ð2Þ
The pH in the reactor was measured at the start of each sampling with a glass electrode calibrated at pH 7 and 10 with commercial buffers. The precision of pH measurements was approximately ±0.02 and the reproducibility of replicate As determinations was about ±0.01 m.
3. Results Fig. 2 shows results from an experiment to determine the time necessary to reach reversible equilibrium. A solution was prepared by saturating 0.351 m Na2CO3 with As2O3 under N2; the pH was about 11.5. A day later, the N2 gas was switched to CO2 gas (1 atm). This quickly converted each mole of CO2 3 to two moles of 2 HCO driving the 3 ðCO2 þ CO3 þ H2 O ! 2HCO3 Þ, pH down to about 7.7 and forcing As(III) to start precipitating. At the same time, another solution containing 0.715 m NaHCO3 under CO2 gas was allowed to begin dissolving As2O3. By the end of one day, the As(III)
0.4
0.3 As(III) (m)
1220
0.2
0.1 From supersaturation From undersaturation
0.0 0
20
40 Time (h)
60
80
Fig. 2. Reversibility test. Data represented by circles are for a run that contained 0.351 m Na2CO3, pre-equilibrated with As2O3 for 24 h at pH 11.5 under N2. At t = 0, CO2 gas was introduced, converting Na2CO3 to 0.715 m NaHCO3 and driving the pH down to 7.7. Data represented by triangles are for a run that contained 0.715 m NaHCO3 at pH 7.7 under CO2 from the beginning.
C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225 Table 1 Experimental data
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Table 1 (continued)
Total As (m)
pH
Reaction time (h)
Run 1 0.250 0.239 0.254 0.243 0.236 0.247 0.247 0.257 0.255 0.246 0.253 0.242 0.231 0.230 0.231 0.237 0.220 0.226 0.252 0.231 0.237 0.260 0.262 0.260 0.244 ± 0.011
6.97 6.97 6.97 6.94 6.94 6.94 6.96 6.96 6.96 6.96 6.96 6.96 7.02 7.02 7.02 7.02 7.02 7.02 6.90 6.90 6.90 6.90 6.90 6.90 Mean ± SD
49 49 49 74 74 74 97 97 97 97 97 97 167 167 167 167 167 167 215 215 215 215 215 215
Run 2 0.253 0.239 0.254 0.251 0.260 0.253 0.264 0.267 0.269 0.269 0.272 0.276 0.249 0.254 0.249 0.247 0.245 0.219 0.266 0.264 0.261 0.269 0.260 0.270 0.258 ± 0.013
7.09 7.09 7.09 7.11 7.11 7.11 7.11 7.11 7.11 7.11 7.11 7.11 7.12 7.12 7.12 7.12 7.12 7.12 7.15 7.15 7.15 7.15 7.15 7.15 Mean ± SD
49 49 49 74 74 74 97 97 97 97 97 97 167 167 167 167 167 167 215 215 215 215 215 215
Run 3 0.181 0.160 0.169
5.54 5.54 5.54
48 48 48
Total As (m)
pH
Reaction time (h)
0.186 0.188 0.173 0.139 0.152 0.152 0.147 0.147 0.143 0.154 0.152 0.151 0.160 ± 0.016
5.64 5.64 5.64 5.54 5.54 5.54 5.67 5.67 5.67 5.67 5.67 5.67 Mean ± SD
73 73 73 144 144 144 218 218 218 218 218 218
Run 4 0.201 0.201 0.190 0.199 0.204 0.211 0.188 0.176 0.184 0.221 0.211 0.240 0.232 0.231 0.220 0.207 ± 0.019
7.74 7.74 7.74 8.03 8.03 8.03 7.97 7.97 7.97 7.71 7.71 7.71 7.71 7.71 7.71 Mean ± SD
48 48 48 73 73 73 144 144 144 218 218 218 218 218 218
Run 1: 0.710 m NaCl, arsenolite + claudetite, pH adjusted with NaOH at start. Run 2: 0.355 m NaCl, 0.357 m NaHCO3, arsenolite + claudetite. Run 3: 0.710 m NaCl, arsenolite only. Run 4: 0.715 m NaHCO3, arsenolite only.
solubility of these two runs was nearly the same (Fig. 2) and after a second day the solubilities had reached a final plateau. For all runs, solubilities did not change systematically with time after 48 h of equilibration. Note that the run that began from undersaturation appears to overshoot the equilibrium point slightly at 24 h. This presumably occurs because a few particles possess excess free energy despite the efforts to prevent this by recrystallizing the starting material; after these particles dissolve, the system equilibrates with less soluble particles. Table 1 presents the solubility data that will be used to explore the stability of As(III) carbonato complexes. Two features of the data are noteworthy. First, the solubility differs for the different As2O3 batches. The same electrolyte (0.710 m NaCl) was used in both Run 1 (arsenolite + claudetite) and Run 3 (arsenolite only), but the average solubility of the first batch was 0.244 (r 0.011) m whereas the average solubility of the third was 0.160 (r 0.016) m. The solubility of the arsenolite + claudetite
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C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225
batch remained metastably high throughout the weeklong experiments and never drifted downward toward that of the arsenolite-only batch (Table 1). The second noteworthy feature in Table 1 is that, for a given As2O3 batch, the As(III) solubility in HCO 3containing solutions is slightly, but distinctly greater than in solutions containing no HCO 3 , even though the solutions are compared at essentially the same ionic strength. For example, in HCO 3 -containing Run 2, average As(III) is 0.258 (r 0.013) m vs. 0.244 (r 0.011) m in HCO 3 -free Run 1. These means are significantly different at the 99% level of confidence. In Run 4, where HCO 3 is twice that in Run 2, the contrast relative to HCO -free Run 3 is even larger: 0.207 3 (r 0.019) m for Run 4 vs. 0.160 (r 0.016) m for Run 3. Thus HCO 3 appears to enhance As(III) solubility under the conditions of these experiments.
4. Discussion Because the enhancement of As(III) solubility by HCO 3 is weak, it was observable only at high ionic strength. Experimental results obtained at high ionic strength can be modeled various ways. Here, a simple ion association approach is adopted, and the extra solubility observed in HCO 3 -containing solutions is attributed to formation of a single carbonato complex. The choice of approach has been influenced by the previous work of Kim et al. (2000), Lee and Nriagu (2003) and Tossell (2004), all of whom describe the role of HCO 3 as that of a complexing ligand. As described below, the extra As(III) solubility observed in the presence of HCO 3 can be explained well by invoking the single complex, AsðOHÞ2 CO 3 . Tossell (2004) concluded that this complex would have modest stability in water, similar to the stability of As(III) oxyhydroxide dimers. He also found that the other complexes proposed by Kim et al. (2000), As(CO3)+ and AsðCO3 Þ 2 , would be unstable with respect to hydrolysis. Neuberger (2004) examined the possibility of fitting the data in Table 1 with either of the other carbonato complexes individually. Both produced statistically slightly poorer fits to the data. Additionally, As(CO3)+ produced fits in which the deviations between predicted and observed solubilities displayed undesirable trends relative to pH and ½HCO 3 . To derive a stability constant for AsðOHÞ2 CO 3 , total dissolved As(III) in the experiments is considered to consist of five species (brackets denote molal concentrations):
½X þ ½NaX0 ¼ ½X ð1 þ K NaX ½Naþ Þ
ð4Þ
Similarly, RNa ¼ ½Naþ þ R½NaX0 ¼ ½Naþ ð1 þ RðK NaX ½X ÞÞ
ð5Þ
Johnson and Pytkowicz (1979) estimate that stoichiometric formation constants for NaCl0 and NaHCO03 ion-pairs in seawater (ionic strength 0.7 m) are KNaX = 0.3. Lacking measured formation constants for NaAsOðOHÞ02 and NaAsðOHÞ2 CO03 , it is assumed that KNaX = 0.3 describes the stability of these ion-pairs as well. This assumption allows equating R(KNaX [X]) in Eq. (5) to 0.3R[X], which owing to the charge balance constraint must equal 0.3[Na+]. Substituting this last term and solving Eq. (5) at RNa = 0.71 m gives [Na+] = 0.601 m in the solutions. Using this value in Eq. (4) gives [X] + [NaX0] = 1.18[X]; thus Eq. (3) can be rewritten, RAsðIIIÞ ¼ ½AsðOHÞ03 þ 1:18½AsOðOHÞ 2 þ 1:18½AsðOHÞ2 CO 3
ð6Þ
Equilibrium concentrations in Eq. (6) are subject to the following mass action laws: 1=2As2 O3 ðsÞ þ 3=2H2 O $ AsðOHÞ03 ; K s ¼
c0 ½AsðOHÞ03 a0:5 s ð7Þ
þ AsðOHÞ03 $ AsOðOHÞ 2 þH ;
Ka ¼
pH c1 ½AsOðOHÞ 2 ½10
c0 ½AsðOHÞ03
AsðOHÞ03 þ HCO 3 $ AsðOHÞ2 CO3 þ H2 O; ½AsðOHÞ2 CO3 Kc ¼ c0 ½AsðOHÞ03 ½HCO 3
ð8Þ
ð9Þ
where c0 and c1 are activity coefficients for neutral and univalent dissolved species and as is the activity of the As2O3 component in the solid phase. Note that in Eq. (9) it has been assumed that the activity coefficients of HCO 3 and AsðOHÞ2 CO3 cancel; any systematic error introduced by this assumption will be embedded quantitatively in the Kc value obtained below. Substituting Eqs. (7)–(9) into Eq. (6) yields: a0:5 K s 1:18K a c0 1þ þ 1:18K c c0 ½HCO3 : RAsðIIIÞ ¼ c0 c1 10pH
RAsðIIIÞ ¼ ½AsðOHÞ03 þ ½AsOðOHÞ 2 þ ½NaAsOðOHÞ02 þ ½AsðOHÞ2 CO 3 þ ½NaAsðOHÞ2 CO03
The concentration of any Na ion-pair, [NaX0,] in the system can be related to the concentration of free anion, [X], by the following equation in which KNaX is the stoichiometric formation constant of the ion-pair in the ionic medium and [Na+] is the molal concentration of free Na+:
ð3Þ
ð10Þ
C.S. Neuberger, G.R. Helz / Applied Geochemistry 20 (2005) 1218–1225
þpH 1 ½Naþ ½Cl K a K s a0:5 Þ s c1 ð10 0:5 1 þ K c K s as
0.0 -0.1 -0.2 0.0 0.1 0.2 0.3 0.4 0.5 0.6 HCO3 (m) 0.2 0.1 0.0 -0.1 -0.2 6.0
6.5
7.0 pH
7.5
8.0
ð11Þ
Note that in writing Eq. (11), it has been assumed that alkalinity, but not HCO 3 , is conserved; in principle, HCO can be non-conserved because of HCO 3 3þ AsðOHÞ3 ¼ H2 CO3 þ AsOðOHÞ 2. Substituting Eqs. (7)–(9) into 11 and rearranging yields: ½HCO 3 ¼
0.1
5.5
½Naþ ¼ ½Cl þ ½HCO 3 þ ½AsOðOHÞ2
þ ½AsðOHÞ2 CO 3
0.2
Observed - Calculated (m)
In order to use this equation to fit the data in Table 1, it is assumed that c0 = 1.09 at I = 0.7 based on measurements by Garrett et al. (1940) of the decrease in As2O3 solubility as a function of HCl concentration. In Runs 3 and 4 (arsenolite only), as is defined to be one, and a value for as is fit in Runs 1 and 2. This attributes the higher solubility of the arsenolite + claudetite batch to a higher activity of the As2O3 component. A Ka value of 109.17 was used (Nordstrom and Archer, 2003), and c1 was set to 0.68 based on the Davies equation. The final fit is not sensitive to uncertainties in c1 and Ka because these quantities affect only the computed concentrations of AsOðOHÞ 2 and its ion-pair, which are minor-to-negligible in Eq. (6) for the solutions considered here. The HCO 3 concentration in Eq. (10) is evaluated from a charge balance equation (neglecting H+ and OH):
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ð12Þ
Based on Eqs. (4) and (5), [Na+] and [Cl] can be replaced by RNa+/1.18 and RCl/1.18. Using these values and Eqs. 10 and 12, the entire set of data in Table 1 is fit by varying the values of Ks, Kc and as until good predictions of observed solubilities were obtained. Fig. 3 shows deviations between the calculated and observed solubilities as a function of ½HCO 3 and pH. Fitting was carried out by the non-linear least squares method, using SCIENTIST (MicroMath Scientific Software). The results are: Ks = 0.178 ± 0.004, Kc = 0.22 ± 0.04 and as (in the arsenolite + claudetite batch) = 2.12 ± 0.08. Quoted uncertainties are 1U and reflect only the statistics of fitting the data, not systematic errors such as might arise from model-related assumptions. A sensitivity analysis shows that changing either c1 or KNaX by ±25% causes the derived constants to change by less than the statistical uncertainty. Thus evaluation of the constants is insensitive to assumptions about these two parameters. The value of Ks for arsenolite is slightly higher than a value of 0.15 ± 0.02 determined for arsenolite by Pokrovski et al. (1996), but lower than the value of 0.20 selected by Nordstrom and Archer (2003) from a survey of literature. There are no previous experimental values of Kc to compare with the one obtained here, but TossellÕs
Fig. 3. Comparison of observed dissolved As(III) to that calculated from the 5-species model represented by Eq. (3) in text. The absence of trends vs. HCO 3 (upper panel) and pH (lower panel) indicates that the model accounts well for the variation of solubility in the data set. Data obtained only afterP48 h equilibration are plotted.
(2004) prediction that AsðOHÞ2 CO 3 would prove to be similar in stability to the dimer, As2 O2 ðOHÞ 3 , is supported by the present result. Writing the formation reaction for the dimer in a format analogous to that of reaction 9, the following equilibrium constant is calculated from the data of Garrett et al. (1940): 0:16 AsðOHÞ03 þ AsOðOHÞ 2 $ As2 O2 ðOHÞ3 þ H2 O; K ¼ 10
ð13Þ This value is reasonably close to the 100.66 obtained here for Kc. The Gibbs free energy of formation of As2O3 in the arsenolite + claudetite batch, based on the value of as, was 1.8 kJ/mol less negative than the corresponding quantity in the arsenolite-only batch. It is tempting to interpret this as the difference in the Gibbs free energy of formation of claudetite vs. arsenolite. However, this idea contradicts the current view that claudetite is the stable polymorph of As2O3 at 25 °C and thus must have a more negative Gibbs free energy of formation than arsenolite (Pokrovski et al., 1996; Nordstrom and Archer, 2003). The fact that the solubility of the arsenolite + claudetite batch remained perched metastably during the weeklong experiment indicates that relaxation of excess Gibbs free energy within As2O3 solids is slow. The source of excess free energy is not known. It is possible that the arsenolite + claudetite batch con-
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tained X-ray amorphous As2O3 that formed during the preliminary recrystallization process. This phase may have controlled the solubility even though crystalline material was present. The results in this paper uphold qualitatively the contention of Kim et al. (2000) and Lee and Nriagu (2003) that As(III) can be complexed by carbonate. However, based on the value of Kc, the HCO 3 concentration would need to exceed about 0.04 m in or0 der for ½AsðOHÞ2 CO 3 to exceed 1% of ½AsðOHÞ3 at near-neutral pH. For comparison, the average HCO 3 concentration is 0.002 m in the ocean and about half that in rivers. Therefore, contrary to the above authors, we conclude that As(III) carbonate complexing will be negligible in most natural waters. To emphasize this point, a speciation calculation for conditions representative of many groundwaters is presented in Fig. 4. Possibly As(III) carbonate species are significant in surface and ground waters of evaporative basins. For example, ½HCO 3 in excess of 0.1 m is known in evaporative lakes in western U.S.A. (Livingston, 1963; Eugster and Hardie, 1973). Welch and Lico (1998) describe mildly alkaline, pH 7–9 groundwaters in Nevada with up to 0.05 m HCO 3 . These also contain exceptionally high As(III) concentrations. In such unusual environments, As(III) carbonate complexes might prove to be significant but minor.
-
H2CO3
1E-3
HCO3
23 CO
[molal concentration]
1E-5
0
As(OH)3
1E-7
-
) OH O( s A
1E-9
2
-
As(OH) 2CO3 1E-11
1E-13
5
6
7
8
9
10
pH
Fig. 4. Calculated species distribution in a solution containing 103 m total carbonate and 107 m total As(III). These concentrations are representative of those commonly found in groundwaters. Activity coefficients have been set to unity in this calculation.
Acknowledgements This work supported by Grant EAR 0229387 from the National Science Foundation. The authors thank Prof. Jack Tossell for originally suggesting this problem to us and for informative subsequent discussions. The manuscript benefited from constructive criticism by J. Webster and an anonymous reviewer.
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