Environ. Sci. Technol. 2000, 34, 1483-1488
Metal-Thiometalate Transport of Biologically Active Trace Elements in Sulfidic Environments. 2. Theoretical Evidence for Copper Thioarsenite Complexing J. A. TOSSELL* Department of Chemistry and Biochemistry, University of Maryland, College Park, Maryland 20742
To obtain an atomistic understanding of the structure and stability of the copper-thioarsenite complex characterized experimentally only by its molecular formula and formation constant in the previous paper (hereafter paper 1) we have carried out quantum mechanical calculations on several isomers of the CuH2AsOS2 complex, a number of related complexes and the corresponding uncomplexed ligands. Using a first principles molecular orbital theory approach which includes electron correlation (second-order MollerPlesset perturbation theory) with a polarized double-ζ valence orbital basis and effective core potentials we calculate the lowest energy isomer to be CuAsS(SH)(OH) (with bonds from Cu to S, SH and As and bonds from As to Cu, S and O), which has a calculated enthalpy of formation from Cu(OH2)2+(aq) and AsS(SH)(OH)-(aq) of about -100 kJ/mol, suggesting a high stability for this complex. Our method also reproduces the experimental trend in log K values for the formation of the simpler complexes CuCl2-, Cu(SH)2- ,and Cu(CN)2- and enables us to understand why the CuAsS(SH)(OH) complex is so stable. We find that the complexes of Cu+ with AsS(SH)2-, AsS(SH)(NH2)-, and AsS(SH)(CH3)- are also very stable, while the complexes with related As(V), P(III), and P(V) S-containing ligands are considerably less stable. The Cu+ complex formed from the purely oxidic ligand AsO(OH)2- is calculated to be unstable. The determining chemical characteristic of the strongly complexing ligands is the presence of two coordinating S atoms (one -S and one -SH) and an electron-rich As center. When the OH group on As is replaced by the strongly electron withdrawing F, the CuAs bond is broken, and the stability of the complex is greatly reduced. We have also calculated deprotonation enthalpies and estimated the acid dissociation constants for the weak acids As(OH)3, As(SH)(OH)2, As(SH)2(OH), and AsO(OH)3. We calculate AsO(OH)3, As(SH)(OH)2, and As(SH)2(OH) to all be much more acidic than As(OH)3. Thus, As(SH)2(OH) will be strongly deprotonated near neutral pH and its anion will complex very strongly with Cu+(aq), while As(OH)3 will exist mainly as the neutral molecule, and neither it nor its anion will complex significantly with Cu+.
* Corresponding author phone: (301)405-1868; fax: (301)3149121; e-mail:
[email protected]. 10.1021/es9901359 CCC: $19.00 Published on Web 03/08/2000
2000 American Chemical Society
Introduction To reproduce the observed Cu solubility of the sulfide mineral assemblage CuS(covellite)-Cu1.8S(digenite)-Cu3AsS4(enargite) in sulfidic solution it was necessary in paper 1 (Clarke and Helz, previous paper in this issue) to invoke the existence of a copper thioarsenite complex with molecular formula CuH2AsOS2, having a formation equilibrium constant of almost 1020. To support the existence of this Cu-thioarsenite complex and to assist in its characterization we have carried out a number of quantum mechanical calculations. We have previously employed quantum chemical methods to characterize various Zn and Cd complexes in aqueous solution (1) as well as complexes of Sb (2) and As (3) In the present work we have examined a number of different isomeric forms and different protonation states for the complex with structural formula CuAsS(SH)(OH) and have also examined the interaction with water of the most stable isomeric form of the neutral complex. We have also calculated enthalpies in solution for the formation of a number of different complexes of Cu+, including those with common ligands such as Cl-, as well as more exotic S, As, and P containing ligands such as AsS(SH)(OH)-. By examining a substantial number of Cu-ligand complexes we can test the accuracy of our approach and compare the stability of the thioarsenite complex with that of complexes formed both from the more common ligands and from other ligands structurally related to the thioarsenite. Enthalpies for the acid dissociation of As(OH)3, As(SH)(OH)2, As(SH)2(OH), and AsO(OH)3 have also been calculated to determine the relative acidity of these acids, so as to predict which acid anions would be prominent in solution in the pH range considered experimentally.
Computational Methods We use mainly the techniques of Hartree-Fock theory and many-body perturbation theory (while employing coupled cluster theory and density functional theory for a few of the molecules studied). The theoretical foundations and capabilities of these techniques are discussed in ref 4(a,b). Recent studies (5) have established that accurate calculation of reaction energies requires consideration of the instantaneous correlations between electron motions, using the techniques of configuration interaction, many-body perturbation theory, coupled cluster theory, or density functional theory. This is particularly important for transition-metal compounds, for which even structural properties are sensitive to electron correlation (6). We have calculated energy-optimized structures for all molecules using the simplest version of MollerPlesset many-body perturbation theory (second-order or MP2, since the sum of the zeroth- and first-orders is simply the Hartree-Fock result (7(a)). The MP2 method is now routinely used for calculations on both main-group and transition-metal organometallic compounds (7(b)). It is very efficient and reliable, failing to give a better description than the Hartree-Fock only for some pathological open-shell systems (7(c)). However, recent studies have shown that it can give poor results for some transition metal systems, producing substantial overbonding, i.e., bonds which are too short and too strong, which can be corrected (7(d)) by going to higher levels of theory, such as coupled cluster approachs. For the complexes Cu(OH2)2+, Cu(SH)2- ,and CuAsS(SH)(OH) we have therefore also carried out geometry optimizations at the fourth-order Moller-Plesset level (MP4) and coupled cluster with double excitations (CCD) levels, considerably more demanding approaches which recover a VOL. 34, NO. 8, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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very large fraction of the energy arising from electron correlation (7(e)). The basis sets used for the Cu-complex calculations were of the relativistic effective core potential type, as implemented by Stevens, Basch, and Krauss (designated SBK, neglecting core electrons but properly describing their effect upon the valence electrons, with the incorporation of relativistic effects (8) or of the standard 6-31G type (4(a,b)). Additional basis functions of d type (polarization functions) are added to each of the non-H atoms to better account for the polarization of the electron density during bond formation. For the studies on the As acids we utilized 6-31+G(2d, p) basis sets, while including diffuse functions to describe the electron density of the anion as well as polarization functions on both the heavy atoms and H. We employed the programs GAMESS (9) and GAUSSIAN94 (10). The calculation of hydration effects upon chemical reactions is currently a very active area within quantum chemistry, with a number of different approaches being employed by various researchers (11, 12). To evaluate hydration enthalpies we use a multipart approach. For monatomic ions, such as Cl-, and for the common small polyatomics, such as OH- or SH-, we use “experimental” hydration energies from the tables of Rashin and Honig (13). Of course the hydration energy of a neutral compound cannot be uniquely divided into contributions from cation and anion, but if we chose the hydration energy of one particular species as a reference, we can define most all other hydration energies with respect to it, obtaining quite consistent values. For H+ we use the newest “experimental” value of the hydration energy (-1150 kJ/mol), taken from Tissandier et al. (14(a)), although the best theoretical value is probably that of -1118 kJ/mol obtained by Tawa et al. (14(b)). For the large polyatomic ions containing Cu we determine the hydration energy by using the reformulated Born model of Rashin and Honig (ref 13), approximating the effective Born radius using the electron-density based criterion described in ref 11 (a surface of electron density 0.001e/(bohr)3) and implemented in GAUSSIAN94. The radii so obtained are quite similar to those estimated using the approach of ref 13 for spherical ions but make allowance for the nonspherical nature of the other ions. For the arsenic acids and their anions we employ both the Born model described above and the isodensity polarized continuum (IPCM) method (11), as implemented in GAUSSIAN94, to calculate hydration energies for both the neutral molecules and their corresponding anions (the IPCM procedure is not implemented in GAUSSIAN94 for the effective core potential basis sets used for the Cu complexes). The IPCM avoids an estimate of the Born radius by employing an electron density criterion for defining a surface, not necessarily spherical, enclosing the species. It is useful to consider what accuracy we might expect to attain in evaluating the energetics of the Cu complexes. As described in ref 5, even an extremely accurate method such as the composite G2 approach of Pople and co-workers, which incorporates electron correlation at a very high level and effectively extrapolates to the infinite basis set limit, cannot obtain reaction energies to better average accuracy than about 6 kJ/mol, resulting in equilibrium constant errors at room temperature of about a factor of 10. Most of the compounds in the G2 test set are also smaller than those considered here, and G2 is a considerably more accurate theory than the polarized SBK MP2 approach we employ. Nonetheless, we have found that relative complex formation enthalpies, within this series of related compounds, can be evaluated with an accuracy near to that of G2, even at our lower level of theory. Invariably, many serious approximations need to be made for calculations on such systems to be feasible. We believe we have developed a computationally feasible model and method to describe the relative energetics 1484
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FIGURE 1. Geometry calculated at the polarized SBK MP2 level for the most stable isomer of CuAsS(SH)(OH). of such complexes. It is important that the method also be simple enough computationally that we can apply it to a wide range of cations and ligands, so as to identify stable species for further more detailed theoretical or experimental study. The accuracy of the gas-phase reaction energies depend only upon the accuracy of the quantum mechanical method, which can be continuously refined to give more reliable results. The more problematic terms in the reaction energies in solution involve our approximations for the hydration energies and our inclusion of only a small number of atoms in our models for the reactive species.
Results Our discussion is based mainly upon results obtained using the polarized SBK basis and the MP2 method for the gasphase calculations. Results of additional calculations using other higher level methods for some of the Cu complexes are presented in the Supplementary Information and are briefly mentioned here to support the MP2 results. We have identified three different isomers of overall composition CuS(SH)As(OH). In Figure 1 we show the most stable isomer, which has the Cu bonded to As, S, and the SH, with bond distances of 2.34, 2.09, and 2.19 Å, respectively. The As is bonded to Cu, S, and O, while the distance from As to the -SH group is now 3.74 Å, clearly a nonbonded value. The two other isomers we have found lie higher in energy by about 100 kJ/mol and so can be excluded from further consideration. The calculated Cu-SH bond distance in this complex is very similar to the value of 2.15 Å we calculate for Cu(SH)2-1 at the polarized SBK MP2 level. For the Cu(SC10H13)2- ion a Cu-SR distance of 2.137 Å is found experimentally (15). The Cu-S distance calculated at the Hartree-Fock level (i.e. without the MP2 correlation correction) is 2.30 Å. For gasphase CuCl our polarized SBK Hartree-Fock and MP2 bond distances are 2.168 and 2.065 Å, respectively, while very large basis set calculations utilizing the multiconfiguration Hartree-Fock method (another way of incorporating electron correlation, (16)) give a distance of 2.059 Å and the experimental bond distances is 2.052 Å. Thus it appears that our polarized SBK MP2 distances reproduce experimental Cu-S and Cu-Cl distances well, at least in simple molecules. For the most stable isomer of CuAsS(SH)(OH) we find a Cu-As bond at the MP2 but not at the Hartree-Fock level of theory. This bond is also present in the equilibrium geometries obtained using MP4, CCD, QCISD (quadratic configuration interaction with single and double excitations), and CASSCF (complete active space SCF) methods, as shown
compare the calculated Cu+ -O distances with experiment since Cu(OH2)n+ is unstable in aqueous solution (although Cu1+ is of course stable in the present of complexing ligands). Since the interaction energy with CuAsS(SH)(OH) for the second water molecule is smaller than the self-interaction energy of water, the best simple formulation of this species in solution is as CuAsS(SH)(OH)‚‚‚H2O, a species in which the Cu+ is essentially four coordinate. As a simple approximation to the complexation reactions of Cu+(aq) in solution we have considered reactions such as
FIGURE 2. Geometry calculated at the polarized SBK MP2 level for CuAsS(SH)(OH)(H2O). in the Supplementary Information. However, even at the Hartree-Fock level the CuAsS(SH)(OH) complex is one of the most stable of all the complexes studied, although it shows only the expected Cu-S and Cu-SH bonds. The bond between nominally Cu(I) and nominally As(III) is certainly unexpected and unusual, and we have been unable to find any compounds in which such a Cu(I)-As(III) bond has been characterized. However, several compounds of trivalent As(III) have been characterized (17) in which one or more normal substituents such as -Cl are replaced by electronrich transition-metal organometallic fragments such as -Cr(CO)5. The most stable isomer of CuAsS(SH)(OH) is also calculated to be stable with respect to both protonation and deprotonation. At the polarized SBK MP2 level (using the experimental hydration energy of H+ and the Born hydration energy of CuAsS2(OH)-, based on the Born radius obtained from GAUSSIAN94) we calculate a energy change of +42 kJ/mol for the deprotonation reaction:
CuAsS(SH)(OH) f CuAsS2(OH)- + H+ Ignoring entropy effects this enthalpy would suggest a pK at room temperature of about +7 for CuAsS(SH)(OH). As discussed more fully below, entropic effects are generally found to increase solution pKa’s by about 5 units. Thus at the near neutral pH values considered in paper 1 the neutral complex would certainly predominate. For the protonation of CuAsS(SH)(OH) in aqueous solution we calculate the energy to be about +140 kJ/mol, since the energy change for binding a proton to CuAsS(SH)(OH) is considerably less than the experimental hydration enthalpy of H+. Thus, the neutral CuAsS(SH)(OH) complex is expected to predominate. We have also considered the interaction of water molecules with Cu+ and CuAsS(SH)(OH), at the MP2 level. By calculating the energy for coordinating four H2O molecules to free Cu+ to form Cu(OH2)4+1 and then adding the Born energy of this species obtained using the Rashin and Honig prescription we obtain a total hydration energy for Cu1+ (at the polarized SBK MP2 level) of -583 kJ/mol, compared to the experimental value (13) of -601 kJ/mol. The interaction energy of a water molecule with neutral CuAsS(SH)(OH) is only about -61 kJ/mol, little more than the hydration energy of H2O in water (approximately the magnitude of the enthalpy of vaporization, 41.5 kJ/mol), and the interaction energy of a second water molecule is only -37.5 kJ/mol. The O of the first water coordinates to the Cu of the CuAsS(SH)(OH) complex with a bond distance of 2.05 Å and causes elongation of the Cu-SH bond by about 0.08 Å and essentially no change in the Cu-As or Cu-S distances, as seen in Figure 2. Addition of a second molecule of water causes changes of no more than 0.01 Å in the bond distances shown in Figure 2. For comparison, the calculated Cu-O distance in Cu(OH2)41+ at the same level of theory is about 2.12 Å. It is impossible to
Cu(OH2)21+ + Cl- f Cu(OH2)Cl + H2O
(1)
Cu(OH2)21+ + 2Cl- f CuCl21- + 2H2O
(2)
and
Cu(OH2)21+ + AsS(SH)(OH)1- f CuAsS(SH)(OH) + 2H2O (3) in which the coordination number of Cu+ is kept at 2 or 3. Of course, both the Cu and As bearing species involved are probably better formulated with additional water molecules as mentioned above, but this simple approach should allow for a proper comparison between the different ligands and it greatly reduces the computational time required (we did establish for the case of CuAsS(SH)(OH) formation that changing eq 3 by adding two additional waters to both reactant and product Cu species changed the overall energetics by less than 5 kJ/mol). Energetic results are given in Table 1 for several reactions of type (1), (2), and (3). We give gas-phase reaction energies obtained at the polarized SBK MP2 level, the hydration energy calculated as the Born hydration energies of the ions (based upon the radii evaluated by GAUSSIAN94, hydration energies of neutrals ignored), the resulting enthalpy change in solution in kJ/mol and the quantity (∆Hin solution)/(2.303RT), which would be equal to log K in the absence of any entropic contributions. We have also established that the gas-phase reaction energy changes by less than 10 kJ/mol for the AsS(SH)(OH)- complex when we optimize all structures at the MP4 level (rather than the MP2). For the formation of Cu(SH)2- (eq type 2) and CuAsS(SH)(OH) (eq type 3) we also calculated the vibrational spectra for all reactants and products and determined the changes in vibrational zeropoint energy and the vibrational, rotational, and translational (VRT) contributions to the enthalpy and free energy. The VRT contribution to ∆G for Cu(SH)2- is only about -13 kJ/ mol, while that for CuAsS(SH)(OH) is about -48 kJ/mol, since in eq type 3 2 mol of reactants are converted to 3 mol of product. This primarily entropic gas-phase effect would tend to increase log K for type 2 reactions by about 2 units and for type 3 reactions by about 8 units, if hydration terms do not compensate for it. This effect has not been included in the numbers shown in Table 1. The Cl-, SH-, and CN- ligands are included primarily because their complexes are simple and reasonably well characterized (although even for these species a range of experimental equilibrium constants have been reported). The other As- and P-containing ligands are included because of our interest in evaluating the complexing abilities of other large, complicated ligands and because they eludicate the influence of chemical substitutions on stability and help us to better understand the high stability of the thioarsenite complex. The increasing stability of the complexes from Clto SH- to CN- is described well by the calculations, as is shown by the comparison with experimental log formation constants given in the last column of Table 1. The calculated VOL. 34, NO. 8, 2000 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Calculated Gas-Phase Energies (from Polarized SBK MP2 Calculations), Hydration Enthalpies, Reaction Enthalpies in Solution, - ∆H/(2.303RT), and Experimental log K Values for the Formation of Various Cu+ Complexes (Energies in kJ/mol) ∆Egas-phase
∆Hhydration
ClSHCN-
-555.8 -598.6 -583.6
+562 +559 +551
2Cl2SH2CN-
-766.6 -831.2 -871.1
+748 +746 +746
CS2OHAs(III)O(OH)21As(III)S(OH)21As(III)S(SH)(OH)1As(III)S(SH)21As(III)S(SH)F1As(III)S(SH)NH21As(III)S(SH)CH31As(V)S2(OH)21P(III)S(SH)(OH)1P(V)S2(OH)21-
-418.0 -364.7 -393.8 -487.0 -466.5 -403.3 -493.0 -460.0 -406.2 -439.0 -420.1
+396 +417 +391 +378 +373 +386 +383 +373 +373 +386 +394
ligand
a
-∆Hin solution* (2.303RT)-1
expa log K
-1.0 +7.0 +5.7
13.0
Eq Type (2) -18.6 -85.2 -125.1 (-121 exp.18c)
3.3 15.0 22.0
≈5 17.2 16-24
Eq Type (3) -22.0 +52.3 -2.8 -109.0 -93.5 -17.3 -110.0 -87.0 -33.2 -53.0 -26.1
3.9 -9.2 0.5 19.1 16.4 3.0 19.3 15.3 5.8 9.3 4.6
∆Hin solution Eq Type (1)
+6.2 -39.6 -32.6
19.8
Reference 18(a,b).
enthalpy change for the Cu(CN)2-1 complex also agrees well with the experimental value (18(c)), given in parentheses. Our calculations also indicate that AsS(SH)(OH)- is indeed the best ligand for Cu+ in the AsO(OH)2- - AsS(SH)2- series (better than the AsS(SH)21- end member). Replacing the -OH group in AsS(SH)(OH)- by -NH2 gives a marginally more stable complex, while replacement of -OH by -F reduces the complex stability substantially. Calculated structures for some of these complexes are shown in Figure 3. The withdrawal of electron density from As in the AsS(SH)F complex causes a breaking of the Cu-As bond. The analogue thiophosphite, PS(SH)(OH)1-, and the analogous As(V) and P(V) ligands AsS2(OH)21- and PS2(OH)21- all give weaker complexes. The reasonably good comparison of calculated and experimental stabilities for the simple ligands supports the rather simple approach used in this study. Specifically, the gas-phase reaction energetics are evaluated at only a moderately accurate level (polarized SBK MP2), the coordination number of the Cu+ is kept small and (almost) fixed, the hydration energy of the ions is evaluated within the Born model, and that of all the neutral molecules is ignored, zeropoint vibrational energy effects are ignored, and entropic effects are ignored. The entropy is in fact expected to increase substantially when ions combine in solution to form a complex, leading to an increase in the formation constant. For example, for the formation of Ag(CN)2-1 the tabulation of Marcus (19) gives an entropic contribution to log K at room T of +5.9. For the formation of CuCl(aq) the same tabulation gives an entropic contribution to log K of about +6.2. We might therefore expect our purely enthalpic calculated log K values to be less positive than experiment by about 6 units. There is certainly some tendency in our results (Table 1) to underestimate the log K values compared to experiment, but the neglect of entropic contributions is only one of our many approximations so adding a standard entropic correction would not improve agreement with experiment. Nonetheless, our results indicate that such quantum mechanical calculations may be a useful and fairly simple way to identify strongly complexing ligands, since the relative log K values seem to be more accurate than the absolute values. 1486
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FIGURE 3. Geometries calculated at the polarized SBK MP2 level for CuAsS(SH)(NH2) and CuAsS(SH)F. To establish that AsS(SH)(OH)- indeed exists in solution so that it can complex with Cu+ we must establish that As-
TABLE 2. Calculated Energetics for Acid Dissociation Reactions of Various As(III) and As(V) Hydroxides and Bisulfides (Energies in kJ/mol)a ∆Egas-phase
reaction
6-31G* SCF (1)
6-31G* MP2 (2)
6-31+G(2d,p) MP2//6-31+ G(2d,p)SCF (3)
As(OH)3 f AsO(OH)21- + H+ As(SH)(OH)2 f AsS(OH)21- + H+ As(SH)2(OH) f AsS(SH)(OH)1- + H+ AsO(OH)3 f AsO2(OH)21- + H+
1536.1 1379.4 1395.7 1402.8
1468.6 1367.3 1389.6 1421.7
1436.6 1349.7 1366.0 1385.2
a
∆Hhydration anion -neutral IPCM (4)
Born (5)
∆Hsolution 6-31+G(2d,p) ∆ZPE MP2 and IPCM, (6-31G* SCF) plus ZPE ∆Hsolution* pKa (6) (3) -1150 -(4) -(6) (2.303RT)-1 exp
-220.8 -213 -187.7 -195.6 -173 -206.6 -221
-31.0 -21.8 -27.3 -30.5
+34.8 -9.8 -6.9 -1.9
6.1 -1.7 -1.2 -0.3
9.2 2.2
Proton hydration energy assumed to be -1150 kJ/mol.
(SH)2(OH) is a strong enough acid that it will be appreciably dissociated near neutral pH. AsO(OH)3 and As(OH)3 have experimental pKa’s of 2.2 and 9.2, respectively, but the pKa’s of the As sulfide acids are not known. If we can establish that As(SH)(OH)2 and As(SH)2(OH) have about the same deprotonation energetics as AsO(OH)3 we would expect them to be deprotonated near neutral pH. For these smaller maingroup species we have used somewhat more accurate quantum mechanical methods, evaluating the gas-phase geometries with an all-electron 6-31+G(2d, p) basis set (containing diffuse functions and multiple polarization functions) at the Hartree-Fock level, incorporating correlation at the MP2 level using the same basis sets, calculating the zero-point vibrational energies and the VRT contributions to enthalpy and free energy at the Hartree-Fock level, and evaluating hydration effects for both the neutral molecule and the anion using the IPCM method. Our results are given in Table 2. We reproduce the observed order of acidity for As(OH)3 and AsO(OH)3 and predict that the S containing acids will be even stronger than the As(V) acid AsO(OH)3. Of course the higher acidity of the -SH acids compared to the -OH acids is expected on qualitative grounds, (e.g. H2S is a much stronger acid than H2O), and we can now confirm this expectation. Note that for acid dissociation the entropy change is negative (20), the reverse of that for the complex formation reactions, with many acid dissociation reactions averaging about -92 J/(mol*K) for ∆S, giving an unfavorable entropic contribution to the equilibrium and so increasing the pKa by about 4.8 units. Increasing our values of ∆G/(2.303RT) (the purely enthalpic contribution to pKa) by 4.8 units clearly brings them into better agreement with experiment for As(OH)3 and AsO(OH)3. More important, the correct qualitative order of relative acidities is given by the ∆Hin solution (and even by the gas-phase energies). Given our calculated pKa values (and the experimental value for As(OH)3) we would expect that As(OH)3 would not be deprotonated at the pH values employed in paper 1 but that As(SH)(OH)2 and As(SH)2(OH) (and probably As(SH)3 as well) would be deprotonated, so that AsS(OH)2-, AsS(SH)(OH)-, and AsS(SH)2-would be potential ligands for Cu+. The calculations reported here are fairly simple by contemporary quantum mechanical standards, but they are accurate enough to provide important data which confirm the stability of the CuAsS(SH)(OH) complex hypothesized to fit the experimental solubility data in paper 1 while at the same time allowing a deeper, atomistic understanding of the structure of this complex. They are also economical enough that they can be used to study the relative stabilities of a large number of different complexes. We confirm that CuAsS(SH)(OH) is indeed a very stable species, with coordination of S, SH, and As to the Cu. It is also stable toward both protonation and deprotonation. Calculated formation constants, estimated entirely from the calculated enthalpies in solution, are consistent with experimental data for simple
complexes such as CuCl2- and correctly predict the high stability of the CuAsS(SH)(OH) complex. Several other As and P sulfide-containing anions are also found to complex strongly to Cu, and trends in stability can be explained on the basis of their electronic structure. Our calculations also confirm that As(SH)2(OH) is a much stronger acid than As(OH)3, although this might admittedly have been predicted on qualitative grounds.
Acknowledgments This work was supported by DOE grant DE-FG02-94ER14467.
Supporting Information Available Higher-level quantum mechanical methods give essentially the same results as those presented here, with CuAsS(SH)(OH) having an unusual Cu(I)-As(III) bond and a high stability. This material is available free of charge via the Internet at http://pubs.acs.org.
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Received for review February 8, 1999. Revised manuscript received January 11, 2000. Accepted January 17, 2000. ES9901359
