Interaction Between Group IIb Divalent Transition-Metal Cations and 3

Article pubs.acs.org/JPCA

Interaction Between Group IIb Divalent Transition-Metal Cations and 3‑Mercaptopropionic Acid: A Computational and Topological Perspective Sabyasachi Bagchi, Debasish Mandal, Deepanwita Ghosh, and Abhijit K. Das* Department of Spectroscopy, Indian Association for the Cultivation of Science, Jadavpur, Kolkata 700032, India S Supporting Information *

ABSTRACT: Density functional theory was applied to study the interaction of group IIb transition-metal cations (Zn2+, Cd2+, and Hg2+) with one and two fully or partially deprotonated 3-mercaptopropionic acid ligands. In this investigation, we determined the geometries of all possible complexes resulting from the coordination of the metal ions with the ligands at different binding sites selected on each ligand. The relative energies of the complexes, metal-ion affinities, free energies, and entropies were also determined. The natures of the bonds were critically analyzed by natural bond orbital (NBO) analysis and clarified further using the atoms-in-molecules (AIM) approach. The substantial influence of the solvent (water) polarization on the energetics, geometries, and bonding of the molecular complexes was also investigated by the conductor-like screening solvation model (COSMO). In an attempt to simulate the complexes in an aqueous environment, water molecules were added explicitly to complete the coordination sphere of the metal cations, and the corresponding metal-ion affinities were calculated to study the effect of microhydration.

1. INTRODUCTION Transition-metal ions have a wide range of applications in many areas of chemistry, including catalysis, organometallic reactions, and biochemistry.1 These metal ions exhibit unusual chemical behavior inside cells. The transition-metal ions can become toxic when their concentration exceeds the natural levels and act as potent disrupters of biological systems. The toxicity of heavy-metal ions owes to their ability to coordinate strongly and react with biologically active groups and to compete effectively for binding sites with the essential metal ions. The toxicities produced by transition metals generally involve neurotoxicity, hepatotoxcity, and nephrotoxicity. Cadmium is a potent carcinogen in rodents2,3 and has also been accepted as a category 1 human carcinogen.4 Being a wellestablished toxic ion, cadmium acts as a major environmental pollutant, causing severe damage to the living world.5−8 Cadmium is emancipated in aquatic environments during rock mineralization processes9 and widely used industrial practices such as electroplating and galvanizing; pigmentation in the paint industry; and the production of nickel−cadmium batteries, pesticides, alloys, chemical reagents, and nuclear rods. Its toxicity depends not only on the total dissolved concentration but also on the free Cd2+ concentration, which controls Cd−organism interactions.10−14 From this point of view, the study of the complexation of cadmium by natural organic ligands in oceans and estuaries is very important.15−20 Increasing and extensive use of mercury in human applications has resulted in an environmental contamination that has made the problem more important over the past few years because it involves the understanding of the fundamentals of interactions © 2013 American Chemical Society

of mercury ions with inorganic and low-molecular-weight organic ligands. Another paradigm of transition-metal toxicity is the release of toxic Zn2+ ions when ZnO nanoparticles dissolve under aqueous conditions to form hydrated Zn2+ cations.21 This dissolution is increased under acidic conditions and in the presence of biological components such as amino acids and peptides.21,22 Cytotoxicity of Zn2+ involves disruption of cellular zinc homeostasis, leading to lysosomal and mitochondrial damage and ultimately cell death.23 Upon metalation, organic molecules form complexes that affect the metal ion’s behavior on a very large scale, especially by changing their toxicity, bioavailability, and mobility. The harmful presence of certain heavy metals is thus controlled by complexation with organic molecules. Among organic compounds, molecular species that contain thiols such as glutathione and metallothioneins or species that are produced by biodegradation of sulfur-containing compounds (such as methane thiols or 3-mercaptopropionic acid) are present in large concentrations in aquatic or marine environments and can form stable complexes with metal ions.24,25 Incidentally, 3mercaptopropionic acid is a molecule that is abundant in the oceanic marine environment and is found to coordinate effectively with cadmium ions, resulting in partitioning of the resultant complexes in aqueous phase, thereby arresting the mobility of the metal ion, which reduces its toxicity. Received: October 3, 2012 Revised: January 18, 2013 Published: January 18, 2013 1601

dx.doi.org/10.1021/jp309792v | J. Phys. Chem. A 2013, 117, 1601−1613

The Journal of Physical Chemistry A

Article

Vairavamurthy et al.26 investigated the structure of cadmium-3mercaptopropionic acid complexes using electrospray mass spectrometry and extended X-ray absorption fine structure (EXAFS) spectroscopy. Belcastro et al.27 performed a detailed density functional theory investigation of the coordination modes, geometrical structures, and binding energies of the complexes formed by cadmium ions and one or two 3mercaptopropionic acid molecules. The objective of the present study is to investigate thoroughly the interaction of one or two fully or partially deprotonated 3-mercaptopropionic acid ligands with the group IIb transition-metal cations (Zn2+, Cd2+ and Hg2+). The present work aimed at exploring the structures of the complexes by finding different possible coordination modes of complexation and determining the stability of these complexes by examining their binding energies and metal-ion affinities (MIAs). We also performed detailed analysis of the bonds and interactions of the complexation processes by two different methods. Efforts were made to evaluate the feasibility of water-assisted complexation processes and delineate the influence of solvent on the energetics and geometries of the molecule through calculations based on the conductor-like screening solvation model (COSMO). All of these studies were performed using density functional theory. The importance of the work is mainly two-fold. First, these ions are biologically important because of their toxic and polluting effects and can compete among themselves for complexation with organic compounds containing thiols. Second, this work can provide insight into the mechanism of complex formation and binding of transition metals with hydrophilic thiols, which are an integral part of many biological systems, leading to the detoxification of the toxic metal ions.

the complete counterpoise (CP) method of Boys and Bernardi.42 In fact, the magnitudes of the BSSE correction terms were carefully analyzed to assay the degree of basis set completeness, and we found very small BSSE values, suggesting that the employed basis sets were effectively complete for the present investigation.43 Metal-ion affinity (MIA) was computed as the negative of the enthalpy variation for the metalation process, which can be represented by the following ligand−M2+ reaction nMPA + M2 + → (MPA)n −M2 +

(1)

MIA was explicitly calculated as MIA = − (Eel [(MPA)n −M2 +] − Eel(n MPA) − Eel(M2 +) + {Evib[(MPA)n −M2 +] − Evib(n MPA)})

(2)

where MPA is a particular ligand, n represents an integer (1 or 2 in this particular case), and M2+ represents a particular metal ion. Eel is the electronic energy obtained from self-consistentfield (SCF) calculations, and Evib is the zero-point energy at 0 K and includes the zero-point energy as well as temperature corrections from 0 to 298 K, as enthalpy was calculated at 298 K. A self-consistent reaction field (SCRF) method was employed to take into account the influence of water as a solvent and to analyze the general complexation mechanism, energetics, and thermochemistry of the complexes, particularly in the aqueous phase. The aqueous-phase calculations were carried out at the same level of theory as the gas-phase ones. The conductor-like screening solvation model (COSMO)44 was used to incorporate the solvent effect. All COSMO calculations in this study were performed using the default choice of the Gaussian 09 program with the recommended standard parameters. The dielectric constant of water used for the calculations with solvent was 78.35 D. Natural bond orbital (NBO) analysis was performed on the optimized structure of each complex to determine the NBO charges on various atoms involved in complexation and also to evaluate bond orders by determining the percentage contributions of the orbitals involved in bond formation in different metalated complexes. The NBO analysis developed by Reed et al.45 to study orbital interactions is a reliable tool for the rationalization of chemical bonds. The NBO population analysis and Wiberg bond index calculation were performed using the NBO 3.1 program46 implemented in Gaussian 09. Our effort to more deeply elucidate the nature and strength of the bonding interactions of different complexes obtained by metalation of free ligands led us to perform a subsequent topological analysis using Bader’s theory of Atoms in Molecules (AIM),47−49 which is based on the topological analysis of the electron density, ρ(r). The necessary indicator for the existence of a particular bond in a complex, regardless of its nature, is the presence of a (3, −1)-type bond critical point (BCP) (in which curvature of the density at the BCP along the line connecting two nuclei is positive and the two curvatures along mutually perpendicular lines and perpendicular to the bond path are negative). Other topological properties such as the gradient (∇ρBCP) and Laplacian (∇2ρBCP) of the electron density were calculated to characterize the nature and strength of the said interactions. The Laplacian of the density (∇2ρ) identifies the regions of space wherein the electronic charge is locally depleted or concentrated. The former situation is typically associated with interactions between closed-shell systems,47

2. COMPUTATIONAL DETAILS All electronic structure calculations were carried out using the Gaussian 09 28 suite of quantum chemistry programs. Equilibrium geometries of the free ligands and their metalated complexes were determined by full optimization without symmetry constraints followed by harmonic frequency calculations to assess the nature of the stationary points and to estimate the zero-point vibrational energy (ZPVE) corrections. The changes in enthalpy (ΔH), entropy (TΔS), and free energy (ΔG) were also evaluated by thermochemical analysis at 298 K. Optimization and vibrational analysis of all systems were performed using the Becke three-parameter hybrid (B3LYP) exchange-correlation functional29,30 in the framework of density functional theory (DFT) in conjunction with the BS basis set. DFT is widely recognized as an effective quantum chemistry method for studying molecular properties, with its reliability in predicting the geometries and energetics of metal cation−π and metal cation−heteroatom complexes compared to other quantum chemistry methods.29−31 The composition of the basis set BS is as follows: 6-311++G** for H, C, S, and O32,33 and LANL2DZ relativistic pseudopotentials34−36 for Zn, Cd, and Hg. To choose the basis set, we considered works on the determination of metal affinities for biological systems,37−40 as well as the fact that a relativistic pseudopotential is essential for elements such as mercury for which the nonrelativistic allelectron approach cannot be successfully applied.41 Because of the size of the basis sets used, the basis set superposition errors (BSSEs) were taken into account for all stable complexes using 1602



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Figure 1. B3LYP/BS-optimized structures of MPA and MPAH and the coordination modes in the interaction of M2+ (M = Zn, Cd, and Hg) ions with one MPA or MPAH ligand.

where the kinetic energy contribution is greater than the potential energy and there is a reduction of electronic charge along the bond path characterized by ∇2ρ > 0 and valid for the ionic bonds, hydrogen bonds, and van der Waals interactions. In contrast, the latter situation characterizes covalent bonds, where the electron density is concentrated in the internuclear region and is unambiguously related to ∇2ρ < 0. Additional information on the chemical bond can be obtained through energetic considerations related to the interplay of changes in the potential and kinetic energies. The Laplacian of ρ appears in the local expression of the virial theorem50 as ℏ2 2 ∇ ρ(r ) = 2G(r ) + V (r ) 4m

E (r ) =

1 V (r ) 2

(4)

V (r ) =

1 2 ∇ ρ(r ) − 2G(r ) 4

(5)

Finally, the criterion used here to determine the nature of bonds was the ratio −G(r)/V(r). For −G(r)/V(r) > 1, the interaction is noncovalent; for 0.5 < −G(r)/V(r) < 1, it is partly covalent.51,52 Weak interactions have positive values for both ∇2ρ(r) and H(r), medium interactions have positive ∇2ρ(r) but negative H(r), and strong interactions have negative values for both ∇2ρ(r) and H(r).53 All topological properties were calculated using XAIM 1.0.54

(3)

3. RESULTS AND DISCUSSION In accordance with the existing experimental findings26 and previous theoretical investigations27 regarding the formation mechanism of different complexes between cadmium ion and 3mercaptopropionic acid, in both the gaseous and liquid phases, we also considered completely deprotonated (MPA) and partially deprotonated (MPAH) ligands. The possible modes of coordination between the metal cations and ligands considered were mainly based on the classification of hard

The quantity G(r) appearing in this expression is a form of the kinetic energy density of the electrons that is positive everywhere, and V(r) is the potential energy density, which is negative everywhere. The electronic energy density H(r), defined as H(r) = G(r) + V(r), is an index of the amount of covalent nature in the chemical interactions. The sign of H(r) determines whether the accumulation of charge at a given point r is stabilizing [H(r) < 0] or destabilizing [H(r) > 0]. The bond energies, E(r), were calculated using the equations 1603



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Table 1. B3LYP/BS Total Energies (Zero-Point-Corrected and BSSE-Corrected) and Relative Energies (ΔE) for Various Conformers of M2+−MPA and M2+−MPAH (M = Zn, Cd, Hg) Complexesa complex Zn −MPA(S,O) Zn2+−MPA(O,O) Zn2+−MPAH(S,O)b Zn2+−MPAH(S,O)a Zn2+−MPAH(S) Zn2+−MPAH(O,O) Cd2+−MPA(S,O) Cd2+−MPA(O,O) Cd2+−MPAH(S,O)b Cd2+−MPAH(S,O)a Cd2+−MPAH(S) Cd2+−MPAH(O,O) Hg2+−MPA(S,O) Hg2+−MPA(O,O) Hg2+−MPAH(S,O)b Hg2+−MPAH(S,O)a Hg2+−MPAH(S) Hg2+−MPAH(O,O) 2+

a

E (au) −731.030242 −730.953775 −731.377303 −731.356068 −731.320751 −731.265787 −713.481172 −713.426076 −713.840986 −713.821432 −713.800721 −713.746957 −708.139339 −708.069637 −708.505121 −708.489091 −708.479276 −708.428752

(−731.065271) (−730.953775) (−731.482751) (−731.471436) (−731.444537) (−731.375804) (−713.525306) (−713.471257) (−713.950933) (−713.942318) (−713.929666) (−713.864551) (−708.175022) (−708.123691) (−708.602893) (−708.596735) (−708.590009) (−708.522619)

Eb (au)

ΔE (kcal mol−1)

ΔEb (kcal mol−1)

−713.479952 −713.425486 −713.840006 −713.820162 −713.803301 −713.746485

0.0 (0.0) 47.9 (45.4) 0.0 (0.0) 13.3 (7.1) 35.4 (23.9) 69.9 (67.1) 0.0 (0.0) 34.5 (33.9) 0.0 (0.0) 12.2 (5.4) 25.2 (13.3) 59.0 (54.2) 0.0 (0.0) 43.7 (32.2) 0.0 (0.0) 10.1 (3.8) 16.2 (8.1) 47.9 (50.3)

0.0 34.2 0.0 12.5 23.0 58.7

Values in parentheses are for the solution phase with water as the solvent. bReference 27.

ion involving the sulfhydryl group is the least favored mode of complexation. Also, in coordinating with the sulfhydryl group, the metal ions form six-membered rings, which are more stable than the four-membered rings formed by the coordination of the metal ions with the carboxylic group. The results of NBO analysis of the most stable M2+−MPA, 2+ M −MPAH, M2+−(MPA)2, and M2+−MPA−MPAH complexes are reported in Table 2. In the most stable M2+−MPA complexes, namely, the M2+−MPA(S,O) complexes, the results of the NBO analysis reflect a strong charge-transfer electrostatic interaction between the S atom and the M2+ ion. The bonding analysis shows a charge-transfer phenomenon from the ligand to the metal-ion center. A closer examination of the M2+−S bonds [76.78% Zn2+ and 23.22% S for the Zn2+−MPA(S,O) complex, 76.40% Cd2+ and 23.60% S for the Cd2+−MPA(S,O) complex, and 68.47% Hg2+ and 31.53% S for the Hg2+− MPA(S,O) complex] reveals that these bonds are mainly due to the contributions of the p orbital of the sulfur atom [91.84%, 91.34%, and 96.43% for the Zn2+−MPA(S,O), Cd2+−MPA(S,O), and Hg2+−MPA(S,O) complexes, respectively] and the s orbital of the metal ion (87.10% for Zn2+, 91.34% for Cd2+, and 94.89% for Hg2+) but comparatively smaller contributions from the p orbital of the metal ions (12.11% for Zn2+, 7.92% for Cd2+, and 3.76% for Hg2+) and the s orbital of the sulfur atom (7.98%, 5.34%, and 3.37% for the respective complexes). The natural charges on the O atom (−0.916|e|, −0.909|e|, and −0.853|e| for Zn 2+−O, Cd 2+ −O, and Hg 2+−O bonds, respectively) and on the metal ions (1.262|e|, 1.286|e|, and 1.114|e| for Zn2+, Cd2+, and Hg2+, respectively) indicate that the M2+−O bonds have ionic character. The weakening of the C− O bond with respect to that in the free ligand [lengths of 1.33, 1.32, and 1.32 Å for the Zn2+−MPA(S,O), Cd2+−MPA(S,O), and Hg2+−MPA(S,O) complexes, respectively, compared to 1.26 Å for MPA) reflects the same phenomenon (see Table S1, Supporting Information). The Wiberg bond index values of the M2+−S bonds (0.7644, 0.7223, and 0.7540 for zinc, cadmium, and mercury, respectively) are greater than those of the M2+−O bonds (0.3128, 0.2589, and 0.2713 for zinc, cadmium and mercury, respectively), which is also corroborated by the fact

and soft acids and bases taking into account their polarizing power.55 Figure 1 depicts the optimized structures of the possible coordination modes of the metal cations (Zn2+, Cd2+, Hg2+) with one 3-mercaptopropionic acid molecule. In the search for the global minimum, a number of initial-guess geometries were considered, but the global minimum for different conformations of the metal-ion complexes were successfully located by analyzing the initial-guess geometries and the geometries reported in other works. Some of the initially optimized structures are also shown in Figure F1 in the Supporting Information. As can be seen from Figure 1, the metal cations can interact and coordinate with both a completely deprotonated (MPA) and partially deprotonated (MPAH) 3-mercaptopropionic acid molecule at different coordination sites to form complexes. The total and relative energies of various conformers of the complexes of M2+−MPA and M2+−MPAH in both the gas and solution phases are summarized in Table 1. It is clear from the table that the bicoordinated complexes, for both MPA and MPAH, in which the metal cations coordinate with both the sulfur and oxygen atoms, are the most stable. For the M2+− MPA complexes, the M2+−MPA(S,O) species was found to be the most stable complex formed. The O,O-bicoordinated species were found to be 47.9, 34.5, and 43.7 kcal mol−1 less stable than the corresponding M2+−MPA(S,O) species for zinc, cadmium, and mercury, respectively, in the gas phase, whereas in the solution phase, they were energetically less stable than their corresponding most stable species by 45.4, 33.9, and 32.2 kcal mol−1, respectively. This result can be attributed to the fact that zinc, cadmium, and mercury are typical soft acids, as they are large atoms with many filled d orbitals that interact and bind strongly with thiols, which are large, easily polarizable, and considered to be soft bases. However, we could not locate any S-monocoordinated species for MPA for any of the metal cations, in agreement with the work of Belcastro et al.,27 who reported that the structure is highly unstable. This indicates that, although the primary functional group in 3-mercaptopropionic acid that binds the metal ion is the sulfhydryl group rather than the carboxylic group, monocoordination with metal 1604



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Table 2. Results of NBO Analysis of the Most Stable M2+−MPA, M2+−MPAH, M2+−(MPA)2, and M2+−MPA−MPAH Complexes complex MPA

MPAH

Zn2+−MPA(S,O)

Zn2+−MPAH(S,O)b

Cd2+−MPA(S,O)

Cd2+−MPAH(S,O)b

Hg −MPA(S,O) 2+

Hg2+−MPAH(S,O)b

Zn2+−(MPA)2(O,S,O,S)

atom

natural charge

S O O1 S O OH Zn2+ S O O1 Zn2+ S O OH Cd2+ S O O1 Cd2+ S O OH Hg2+ S O O1 Hg2+ S O OH Zn2+ S S1 O1 O2 O3 O4

−0.813 −0.831 −0.821 −0.681 −0.645 −0.710 1.262 −0.405 −0.916 −0.582 1.304 −0.300 −0.774 −0.591 1.286 −0.414 −0.909 −0.592 1.328 −0.296 −0.762 −0.607 1.114 −0.305 −0.853 −0.583 1.142 −0.125 −0.720 −0.623 1.178 −0.553 −0.553 −0.834 −0.745 −0.834 −0.745

complex Cd2+−(MPA)2(O,S,O,S)

Hg2+−(MPA)2(O,S,O,S)

Zn2+−(MPA)−MPAH(O,S,O,S)

Cd2+−(MPA)−MPAH(O,S,O,S)

Hg2+−(MPA)−MPAH(O,S,O,S)

that the M2+−S bonds are stronger electrostatic bonds than the M2+−O bonds. The same trends of bond index values were observed in the solution phase, but the corresponding values were smaller in magnitude, indicating that the bonds were somewhat weaker in water, which is a polar solvent. The data compiled in Table S2 (Supporting Information) reflect the fact that the action of the solvent (water in the present case; COSMO, B3LYP/BS level) strongly modifies the NBO charge distribution over the atoms. Such a significant effect of the solvent on the charge distribution naturally implies that the solvent should influence the internuclear distances as well (see Table S1, Supporting Information). The calculated AIM parameters are presented in Table 3. For the M2+−MPA(S,O) complexes, the exploration of the critical points for the interaction of MPA with M2+ revealed the existence of two (3, −1) critical points that connect the cation with the sulfur and oxygen atoms. The Laplacians of the charge density at the (3, −1) critical points have positive values, indicating a depletion of electron energy, a characteristic of closed-shell interactions. It was found that the ∇2ρ values for the (3 −1) critical points corresponding to the M2+−S bonds

atom

natural charge

Cd2+ S S1 O1 O2 O3 O4 Hg2+ S S1 O1 O2 O3 O4 Zn2+ S S1 O1 OH O3 O4 Cd2+ S S1 O1 OH O3 O4 Hg2+ S S1 O1 OH O3 O4

1.297 −0.592 −0.592 −0.849 −0.749 −0.849 −0.749 1.161 −0.536 −0.536 −0.824 −0.759 −0.824 −0.759 1.132 −0.522 −0.487 −0.580 −0.675 −0.868 −0.694 1.236 −0.553 −0.528 −0.587 −0.675 −0.882 −0.700 1.091 −0.469 −0.463 −0.613 −0.715 −0.862 −0.694

for all three metal cations were greater than those for the critical points corresponding to the M2+−O bonds. Moreover, the values of the total energy density, H(r), for these two bond critical points (BCPs) were negative for all three metal cations, indicating these bonds involve medium interactions. Also, for all three metal cations, the criterion 0.5 < −G(r)/V(r) < 1.0 is satisfied for the BCPs corresponding to both types of bonds (see Table 3), indicating that, even though these bonds are electrostatic bonds, they are also faintly covalent in nature, with the M2+−S bonds being stronger than the M2+−O bonds. These results once again strongly support the bonding characteristics explained earlier by the NBO analysis and the theory of hard and soft acids and bases. In the case of partially deprotonated 3-mercaptopropionic acid (MPAH), the metalation process leads to the formation of M2+−MPAH complexes. As can be seen in Table 1, the M2+− MPAH(S,O)b species emerged as the most stable for all three +

metal cations. The M2 −MPAH(S,O)a species were energetically less stable (by 13.3, 12.2, and 10.1 kcal mol−1 for Zn, Cd, and Hg, respectively). This is due to the repulsion of the metal cations by the hydrogen atom attached to the oxygen atom to 1605



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Table 3. B3LYP/BS-Calculated Values of Electron Density (ρ), Laplacian of the Electron Density (∇2ρ), Kinetic Energy Density [G(r)], Potential Energy Density [V(r)], and Electronic Energy Density [H(r)] in au and the Ratio of G(r) and V(r) at the Bond Critical Point (BCP) for Metal Ion−Ligand Bonds of the Most Stable M2+−MPA, M2+−MPAH, M2+−(MPA)2, and M2+−MPA− MPAH Complexes complex Zn −MPA(S,O) 2+

Cd2+−MPA(S,O) Hg2+−MPA(S,O) Zn2+−MPAH(S,O)b Cd2+−MPAH(S,O)b Hg2+−MPAH(S,O)b Zn2+−(MPA)2(O,S,O,S)

Cd2+−(MPA)2(O,S,O,S)

Hg2+−(MPA)2(O,S,O,S)

Zn2+−(MPA)−MPAH(O,S,O,S)

Cd2+−(MPA)−MPAH(O,S,O,S)

Hg2+−(MPA)−MPAH(O,S,O,S)

bond

ρ

2 ∇ ρ

G(r)

V(r)

H(r)

−G(r)/V(r)

Zn −S Zn2+−O Cd2+−S Cd2+−O Hg2+−S Hg2+−O Zn2+−S Zn2+−O Cd2+−S Cd2+−O Hg2+−S Hg2+−O Zn2+−S Zn2+−S1 Zn2+−O1 Zn2+−O3 Cd2+−S Cd2+−S1 Cd2+−O1 Cd2+−O3 Hg2+−S Hg2+−S1 Hg2+−O1 Hg2+−O3 Zn2+−S Zn2+−S1 Zn2+−O1 Zn2+−O3 Cd2+−S Cd2+−S1 Cd2+−O1 Cd2+−O3 Hg2+−S Hg2+−S1 Hg2+−O1 Hg2+−O3

0.0892 0.0739 0.0769 0.0658 0.0692 0.0655 0.0645 0.0741 0.0512 0.0656 0.0386 0.0649 0.0514 0.0514 0.0573 0.0573 0.0487 0.0487 0.0499 0.0499 0.0539 0.0539 0.0399 0.0399 0.0605 0.0563 0.0250 0.0692 0.0569 0.0526 0.0183 0.0607 0.0632 0.0581 0.0113 0.0540

0.5171 0.1430 0.3715 0.1507 0.2815 0.1194 0.3241 0.0741 0.2543 0.1448 0.1624 0.1095 0.0905 0.0905 0.2710 0.2710 0.1249 0.1249 0.2497 0.2497 0.1119 0.1119 0.1683 0.1683 0.1132 0.1008 0.0835 0.3637 0.1400 0.1306 0.0773 0.3018 0.1198 0.1167 0.0380 0.2293

0.1387 0.0597 0.1085 0.0545 0.0819 0.0469 0.0879 0.0585 0.0671 0.0531 0.0419 0.0442 0.0389 0.0389 0.0748 0.0748 0.0396 0.0396 0.0656 0.0656 0.0392 0.0392 0.0438 0.0438 0.0476 0.0433 0.0249 0.0979 0.0471 0.0428 0.0174 0.0831 0.0455 0.0424 0.0107 0.0633

−0.1481 −0.0837 −0.1241 −0.0714 −0.0934 −0.0639 −0.0948 −0.0984 −0.0706 −0.0700 −0.0432 −0.0610 −0.0553 −0.0553 −0.0818 −0.0818 −0.0480 −0.0480 −0.0688 −0.0688 −0.0505 −0.0505 −0.0454 −0.0454 −0.0669 −0.0613 −0.0289 −0.1048 −0.0592 −0.0529 −0.0193 −0.0907 −0.0611 −0.0557 −0.0119 −0.0692

−0.0094 −0.0239 −0.0156 −0.0169 −0.0115 −0.0170 −0.0069 −0.0399 −0.0035 −0.0169 −0.0013 −0.0168 −0.0164 −0.0164 −0.0070 −0.0070 −0.0084 −0.0084 −0.0032 −0.0032 −0.0113 −0.0113 −0.0016 −0.0016 −0.0193 −0.0180 −0.0040 −0.0069 −0.0121 −0.0101 −0.0019 −0.0076 −0.0156 −0.0133 0.0012 −0.0059

0.9364 0.7137 0.8743 0.7636 0.8769 0.7340 0.9272 0.5945 0.9504 0.7586 0.9699 0.7246 0.7034 0.7034 0.9144 0.9144 0.8250 0.8250 0.9535 0.9535 0.7762 0.7762 0.9648 0.9648 0.7115 0.7063 0.8616 0.9342 0.7956 0.8090 0.9016 0.9162 0.7447 0.7612 0.8937 0.9147

2+

which the metal cations are coordinated. The M2+−MPAH(S) class of complexes comes next in the order of stability. This is once again consistent with the observation that monocoordination with metal ion involving the sulfhydryl group is the least favored mode of complexation, even though the sulfhydryl group is the main functional group for metal-ion coordination. The most unstable species are the M2+−MPAH(O,O) complexes, in which the metal ions are bicoordinated with the two oxygen atoms. Because of the hard nature of the oxygen atoms, the binding of the metal cations with the least favored carboxylic group makes these species unstable. In addition, the formation of six-membered rings in the M2+− MPAH(S,O)b complexes accounts for their stabilities in comparison to the presence of four-membered rings in the M2+−MPAH(O,O) complexes. As an obvious succession to the preceding discussion, we attempted to determine the impact of the solvent (water) on complex formation. It was observed that the relative stabilities of the complexes followed the same pattern, the relative energies being slightly lower for different

sets of complexes, for all of the metal ions, as compared to the values obtained from the gas-phase calculations. In the course of elucidating the nature of bonding in these complexes, we now turn our attention to the results of the NBO analysis reported in Table 2. From the electronic charge distribution of the most stable M2+−MPAH(S,O)b complexes, it is evident that more charge was transferred from the sulfur atom to the metal cations than to the oxygen atom. The charge densities on the oxygen atom were −0.774|e|, −0.762|e|, and −0.720|e| for the Zn2+−MPAH(S,O)b, Cd2+−MPAH(S,O)b, and Hg2+−MPAH(S,O)b complexes, respectively, which are more negative than that of the free ligand by −0.129|e|, −0.117| e|, and −0.075|e|, respectively. This observation corroborates the ionic nature of the M2+−O bonds in these complexes. An obvious effect of this trend is the increase of the C−O bond lengths in the complexes with respect to that in the free ligand (MPAH), as reported in Table S1 of the Supporting Information. In these complexes, the analysis of the M2+−S bonds [73.30% Zn2+ and 26.70% S for the Zn2+−MPAH(S,O)b complex, 72.16% Cd2+ and 27.84% S for the Cd2+−MPAH1606



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Table 4. Variations in Enthalpy (ΔH = MIA), Free Energy (ΔG), and Entropy (TΔS) for the Formation Process of M2+−MPA and M2+−MPAH (M = Zn, Cd, Hg) Complexes Computed at the B3LYP/BS Levela complex Zn −MPA(S,O) Zn2+−MPA(O,O) Zn2+−MPAH(S,O)b Zn2+−MPAH(S,O)a Zn2+−MPAH(S) Zn2+−MPAH(O,O) Cd2+−MPA(S,O) Cd2+−MPA(O,O) Cd2+−MPAH(S,O)b Cd2+−MPAH(S,O)a Cd2+−MPAH(S) Cd2+−MPAH(O,O) Hg2+−MPA(S,O) Hg2+−MPA(O,O) Hg2+−MPAH(S,O)b Hg2+−MPAH(S,O)a Hg2+−MPAH(S) Hg2+−MPAH(O,O) 2+

a

MIA (ΔH298 K) (kcal mol−1) 626.6 579.1 431.1 417.5 396.2 361.6 579.3 544.7 391.6 379.1 366.2 332.4 592.1 591.1 408.1 398.1 391.9 360.1

(158.2) (113.2) (137.0) (129.7) (113.7) (70.3) (116.4) (82.7) (100.2) (94.5) (86.8) (46.0) (138.6) (106.5) (123.8) (119.9) (115.8) (73.4)

MIA (ΔH298 K)b (kcal mol−1)

free energy (ΔG298 K) (kcal mol−1) 616.6 571.1 421.4 408.3 389.9 354.4 569.9 536.5 382.3 370.1 359.5 325.4 582.8 579.8 399.2 389.6 385.1 353.5

579.3 545.1 393.9 381.4 370.4 334.7

(148.7) (105.2) (127.3) (120.3) (106.5) (62.4) (107.4) (74.4) (90.9) (85.7) (80.0) (37.9) (129.9) (98.8) (114.7) (112.2) (108.8) (66.1)

entropy (TΔS298 K) (kcal mol−1) 10.0 (9.5) 8.0 (8.0) 9.7 (9.74) 9.2 (9.4) 6.3 (7.1) 7.2 (7.9) 9.4 (9.0) 8.2 (8.3) 9.3 (9.3) 9.0 (8.8) 6.7 (6.8) 7.0 (8.1) 9.3 (8.7) 11.3 (7.7) 8.9 (9.1) 8.5 (7.7) 6.8 (7.0) 6.6 (7.3)


(S,O)b complex, and 61.27% Hg2+ and 38.73% S for the Hg2+− MPAH(S,O)b complex) reveals that these bonds are mainly due to charge transfer from the sulfur atom to the metal cations. A detailed inspection also indicates that the p orbital of the sulfur atom [94.61%, 96.69%, and 97.91% for the Zn2+− MPAH(S,O)b, Cd2+−MPAH(S,O)b, and Hg2+−MPAH(S,O)b complexes, respectively] and the s orbital of the metal ion (93.71% for Zn2+, 96.67% for Cd2+, and 98.03% for Hg2+) contribute more to the formation of the M2+−S bonds, whereas the p orbital of the metal ions (5.74% for Zn2+, 2.86% for Cd2+, and 0.98% for Hg2+) and the s orbital of the sulfur atom (5.11%, 3.08%, and 1.85% for the respective complexes) contribute less. A careful assessment of the Wiberg bond index values provides another window toward examining the M2+−S and M2+−O bonds. For the M2+−S bonds, the bond index values are 0.8626, 0.8401, and 0.9350 for zinc, cadmium, and mercury, respectively, whereas for the M2+−O bonds, the corresponding values are 0.1603, 0.1029, and 0.0745. This indicates that the M2+−S bonds are much stronger than the M2+−O bonds. Similar trends were observed in the solution phase where water was used as the solvent, the only change being the lowering of the bond index values for the two types of bonds. This behavior can be attributed to the phenomenon of the exertion of the polar solvent’s reaction field, which leads to weakening of the bonds. As a second strategy to define and describe the bonding interactions of the most stable M2+−MPAH(S,O)b complexes, we focused on the results of the AIM analysis reported in Table 3. We were able to identify two (3, −1) bond critical points belonging to two separate bond paths connecting the sulfur and the oxygen centers of MPAH with the metal cations. At the bond critical points, the values of the Laplacian of the charge density were positive, which is an obvious consequence of closed-shell interactions indicating the depletion of energy. The M2+−S bonds were stronger than the M2+−O bonds in all of the M2+−MPAH(S,O)b complexes, as the ∇2ρ values at the (3, −1) critical points corresponding to M2+−S bonds were greater than those at the (3, −1) critical points corresponding to M2+− O bonds. The total energy densities, H(r), for these two types

of bond critical points (BCPs) were negative for all three metal cations, which indicates that these bonds are due to medium interactions. The values of −G(r)/V(r) at both these bond critical points lie in between 0.5 and 1.0, which indicates that these bonds are also weakly covalent in nature. These results are clearly in good agreement with the NBO results. The metal-ion affinities (MIAs) (ΔH298 K), free energies (ΔG298 K), and entropies (TΔS298 K) of the M2+−MPA and M2+−MPAH complexes in both gaseous and solution phases are displayed in Table 4. The results clearly support our discussions in the previous sections that the M2+−MPA(S,O) and M2+−MPAH(S,O)b complexes for all three metal cations are the most stable complexes. Interestingly, for the fully deprotonated 3- mercaptopropionic acid, monothio complexes were not found to exist in zwitterionic form [M2+−MPA(S)]; rather, the charge-neutralized cyclic species were found to be most stable. Thus, for both MPA and MPAH ligands, the formation of a cyclic structure upon metalation involving the sulfhydryl group is the most preferred mode of complexation compared to metalation involving the carboxylic group, which is in accordance with the theory of hard and soft acids and bases. However, it was observed that, for all complexes, the metal-ion affinities and free energy values decreased considerably in the solution phase. The insertion of the reaction field of the polar solvent (water) into the system was found to influence these quantities most strongly. Also, judging from the MIA values, it is evident that both fully deprotonated (MPA) and partially deprotonated (MPAH) ligands bonded with the metal ions in the order Zn2+ > Hg2+ > Cd2+. To simulate a more realistic aqueous system, we carried out microhydration of the metal cations, which completed the coordination spheres of the ions. In the most stable M2+− MPA(S,O) complexes two, three, and four water molecules were added successively to the metal cations. It was observed that, for all three metal cations, when two water molecules were added, one water molecule remained in the inner sphere whereas the other was pushed out from the coordination sphere and the system was stabilized through strong hydrogen bonds. When three water molecules were added, two remained in the 1607



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Figure 2. B3LYP/BS-optimized coordination modes in the interaction of M2+ (M = Zn, Cd, and Hg) ions with two MPA ligands or one MPA and one MPAH ligand.

inner sphere, whereas when four water molecules were added, three water molecules remained in the inner sphere. The MIA values also increased with increasing number of water molecules for all three metal cations (see Table S2, Supporting Information). However, the MIA values were also greater than the corresponding MIA values for the same complexes in the gas phase. This is due to the strong hydrogen bonds that stabilizes the hydrated complexes. The anomaly between the results obtained by the conductor-like screening solvation model (COSMO) and those obtained by microhydration can be attributed to the fact that the solvation energies calculated by the COSMO model were not very accurate, because the nonnegligible thermal effects, entropic effects, and nonelectrostatic effects between the solute and the solvent were not considered here. For the interactions of the metal cations with two ligand molecules, two possible categories of interactions were considered. We first considered the interaction of the metal cations with two fully deprotonated (MPA) ligands. The

possible modes of coordination of metal cations with two MPA ligands are depicted in Figure 2. As already established, the coordination of a metal ion with the sulfhydryl group giving rise to a cyclic structure of the resultant complex for one MPA ligand is the most favorable mode of coordination. It was expected that, for two ligand molecules, the particular mode in which the metal-ion coordinates with two ligands involving the sulfhydryl group would also be the most favorable. From Table 5, it is evident that the results were in accord with our expectations. For all three metal cations, M 2 + − (MPA)2(O,S,O,S) species were most stable, having a distorted tetrahedral structure. The O−M2+−S angle of the complexes with one MPA cycle is 98.47° for Zn, 92.66° for Cd, and 89.17° for Hg. The S−M2+−S angles had values of 121.44°, 128.14°, and 138.86° for Zn, Cd, and Hg, respectively, whereas the corresponding values for the O−M2+−O angles were 108.02°, 113.94°, and 118.25°. In this regard, the tetrahedral geometry of the complexes arising from the interaction of Cd2+ ions with other ligands, reported in earlier works,27 agree to a good extent 1608



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Table 5. B3LYP/BS Total Energies (Zero-Point-Corrected and BSSE-Corrected) and Relative Energies (ΔE) for Various Conformers of M2+−(MPA)2 and M2+−MPA−MPAH (M = Zn, Cd, Hg) Complexesa complex Zn −MPA2(O,S,O,S) Zn2+−MPA2(S,S,O) Zn2+−MPA2(S,S) Zn2+−MPA−MPAH(O,S,O,S) Zn2+−MPA−MPAH(S,S,O) Zn2+−MPA−MPAH(S,S) Cd2+−MPA2(O,S,O,S) Cd2+−MPA2(S,S,O) Cd2+−MPA2(S,S) Cd2+−MPA−MPAH(O,S,O,S) Cd2+−MPA−MPAH(S,S,O) Cd2+−MPA−MPAH(S,S) Hg2+−MPA2(O,S,O,S) Hg2+−MPA2(S,S,O) Hg2+−MPA2(S,S) Hg2+−MPA−MPAH(O,S,O,S) Hg2+−MPA−MPAH(S,S,O) Hg2+−MPA−MPAH(S,S) 2+

a

E (au) −1396.608986 −1396.595834 −1396.559006 −1397.211691 −1397.200391 −1397.146744 −1379.059775 −1379.049553 −1379.022518 −1379.661466 −1379.651035 −1379.609453 −1373.707565 −1373.702095 −1373.695591 −1374.320125 −1374.314212 −1374.282421

(−1396.879282) (−1396.856673) (−1396.822535) (−1397.312096) (−1397.303791) (−1397.269051) (−1379.324784) (−1379.30916) (−1379.288662) (−1379.761442) (−1379.756596) (−1379.7354) (−1373.970832) (−1373.964779) (−1373.955367) (−1374.413017) (−1374.411699) (−1374.401796)

Eb (au)

ΔE (kcal mol−1)

ΔEb (kcal mol−1)

−1379.055931 −1379.046346 −1379.020363 −1379.679118 −1379.672113 −1379.610182

0.0 (0.0) 8.2 (14.2) 31.3 (35.6) 0.0 (0.0) 7.1 (5.2) 40.7 (21.8) 0.0 (0.0) 6.4 (9.8) 23.3 (22.7) 0.0 (0.0) 6.5 (3.0) 32.6 (13.3) 0.0 (0.0) 3.4 (3.7) 7.5 (9.7) 0.0 (0.0) 3.7 (0.8) 23.6 (7.0)

0.0 6.0 22.3 0.0 4.4 43.3


with our results. The M2+−(MPA)2(S,S,O) species, in which the metal ions were bicoordinated with one ligand through the sulfhydryl group and monocoordinated with the other one through the sulfur atom, were less stable (by 8.2, 6.4, and 3.4 kcal mol−1 for Zn, Cd, and Hg, respectively) than the most stable species. The M2+−(MPA)2(S,S) species were found to belong to the energetically least stable category of complexes. In the M2+−(MPA)2(O,S,O,S) species, two six-membered rings involving the metal cations are present, whereas in the M2+− (MPA)2(S,S,O) complexes, there is one six-membered ring, and in the M2+−(MPA)2(S,S) complexes, there is no ring-like structure. This indicates that the mechanical strain in the M2+− (MPA)2(O,S,O,S) complexes was least, followed by the M2+− (MPA)2(S,S,O) complexes, whereas it was a maximum in the M2+−(MPA)2(S,S) species, which also accounts for the order of stability of these complexes. The results in the solution phase also exhibited the same trends. The phenomenon of bond formation in these complexes is a consequence of substantial modulation of the electronic charge density distribution on the atoms of the ligand molecule and the metal ion that take part in coordination and metalation. Close inspection of the NBO results for the M 2+ − (MPA)2(O,S,O,S) complexes, as reported in Table 2, shows their electronic behavior. Analysis of the bonding between the metal cations and the ligands reveals that the bonds are solely due to charge transfer from the ligand atoms (S and O atoms) to the metal center and that the interactions are electrostatic in nature. However, we observe an intriguing feature for all these complexes. The charges transferred to the metal cations from the sulfur atoms were less than those transferred from the sulfur atoms in the complexes in which a single MPA ligand molecule was coordinated with the metal cations [M2+−MPA(S,O) complexes]. The Wiberg bond index values for the M2+−S and M2+−S1 bonds (0.4660, 0.4788, and 0.4076 for zinc, cadmium, and mercury, respectively) are greater than those for the M2+− O1 and M2+−O3 bonds (0.3128, 0.2589, and 0.2713 for zinc, cadmium, and mercury, respectively). These values indicate that the bonds of the metal cations with sulfur atoms and with the oxygen atoms were of almost the same strength, but

interestingly, the metal cation−sulfur bonds were weaker than those in the complexes in which a single ligand was coordinated with the metal cations [M2+−MPA(S,O) complexes]. The increase in the M2+−S and M2+−S1 bond lengths (see Table S1, Supporting Information) also reflects the weakening of these bonds. This can be attributed to the crowding effect due to the presence of too many atoms around the metal cations. Similar behavior was observed for these complexes in the solution phase as well. At this stage, it is pertinent to substantiate the bonding interactions of the most stable M2+−(MPA)2(O,S,O,S) complexes by the AIM analysis summarized in Table 3. For all three metal cations, four (3, −1) bond critical points, belonging to four separate bond paths that connect the sulfur and oxygen centers of the two MPA ligands with the metal cations, were identified. The Laplacians of the charge density (∇2ρ) at all of these (3, −1) critical points were found to have positive values, which once again indicates that all of the considered interactions were closed-shell interactions, indicating energy depletion. The ∇2ρ values at the (3, −1) critical points corresponding to M2+−S and M2+−S1 bonds were slightly less than those at the (3, −1) critical points corresponding to M2+−O1 and M2+−O3 bonds, which indicates that all of these bonds were of nearly equal strengths. Interestingly, the ∇2ρ values at the bond critical points corresponding to the M2+−S and M2+−S1 bonds were much smaller than those at the critical points corresponding to the M2+−S bonds in the M2+−MPA(S,O) complexes, which supports the fact that the metal cation−sulfur bonds in the present case were weaker than the same bonds in the complexes in which a single ligand was coordinated with the metal cations [M2+−MPA(S,O) complexes]. These findings are corroborated by the NBO analysis presented earlier. The negative values of the total energy density, H(r), at all of the bond critical points for all of the metal cations indicate that these interactions were medium interactions. The −G(r)/V(r) values for all the considered bonds also indicate that they had minimal covalent nature, although all these bonds were 1609



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Table 6. Variations in Enthalpy (ΔH = MIA), Free Energy (ΔG), and Entropy (TΔS) for the Formation Process of M2+− (MPA)2 and M2+−MPA−MPAH (M = Zn, Cd, Hg) Complexes Computed at the B3LYP/BS Levela complex

MIA (ΔH298 K) (kcal mol−1)

Zn2+−MPA2(O,S,O,S) Zn2+−MPA2(S,S,O) Zn2+−MPA2(S,S) Zn2+−MPA−MPAH(O,S,O,S) Zn2+−MPA−MPAH(S,S,O) Zn2+−MPA−MPAH(S,S) Cd2+−MPA2(O,S,O,S) Cd2+−MPA2(S,S,O) Cd2+−MPA2(S,S) Cd2+−MPA−MPAH(O,S,O,S) Cd2+−MPA−MPAH(S,S,O) Cd2+−MPA−MPAH(S,S) Hg2+−MPA2(O,S,O,S) Hg2+−MPA2(S,S,O) Hg2+−MPA2(S,S) Hg2+−MPA−MPAH(O,S,O,S) Hg2+−MPA−MPAH(S,S,O) Hg2+−MPA−MPAH(S,S)

734.6 (217.9) 727.3 (205.0) 705.1 (184.7) 700.5 (207.7) 693.8 (202.7) 660.6 (181.7) 687.8 (167.8) 681.8 (158.7) 665.1(146.3) 652.5 (159.4) 646.3 (156.4) 620.2 (143.5) 693.9 (188.2) 691.4 (184.8) 687.3 (179.2) 666.1 (183.3) 662.1 (182.2) 642.3 (176.2)

a

MIA (ΔH298 K)b (kcal mol−1)

free energy (ΔG298 K) (kcal mol−1) 712.5 707.6 687.6 681.3 672.9 644.3 666.4 662.7 648.4 634.1 626.1 604.0 674.9 670.4 670.4 646.1 643.3 625.8

688.2 681.8 665.1 668.8 664.3 624.2

(196.5) (185.8) (167.7) (186.7) (184.1) (164.5) (147.7) (140.4) (129.3) (138.8) (138.4) (126.8) (168.6) (166.2) (161.9) (163.3) (163.7) (159.3)

entropy (TΔS298 K) (kcal mol−1) 22.1 19.7 17.5 19.2 20.9 16.3 21.4 19.1 16.7 18.4 20.2 16.2 19.0 21.0 16.9 20.0 18.8 16.5

(21.4) (19.2) (17.0) (21.0) (18.6) (17.2) (20.1) (18.3) (17.0) (20.6) (18.0) (16.7) (19.6) (18.6) (17.3) (20.0) (18.5) (16.9)


complexes involving the metal ions plays an important role in the order of stability of the complexes. We now consider the results of NBO analysis to gain insight into the chemical bonds in the most stable M2+−MPA− MPAH(O,S,O,S) complexes. Bond formation upon metalation is a direct consequence of charge transfer from the ligand atoms (S, S1, O1, and O3) to the metal cations, as seen from Table 2, and the interactions are electrostatic interactions. The Wiberg bond index values for the M2+−S bonds are 0.5620, 0.5094, and 0.6067, and those for M2+−S1 bonds are 0.5031, 0.4445, and 0.4778 for zinc, cadmium and mercury cations, respectively. This indicates that the bonds of the metal cations with the sulfur atoms are of nearly equal strengths. Interestingly, these bonds are weaker than the M2+−S bonds in the complexes where a single ligand molecule is coordinated with the metal cations, that is, in both the M2+−MPA(S,O) complexes and M2+−MPAH(S,O)b complexes. The increase in the M2+−S and M2+−S1 bonds in the M2+−MPA−MPAH(O,S,O,S) complexes compared to those in the M2+−MPA(S,O) and M2+− MPAH(S,O)b complexes clearly dictates this phenomenon, which can again be attributed to the crowding effect due to the presence of a number of atoms around the metal cations. The bond indices for M2+−O1 bonds are 0.0951, 0.0616, and 0.0505 for Zn2+, Cd2+ and Hg2+, respectively, whereas for M2+−O3 bonds the corresponding values are 0.2538, 0.1892, and 0.1863. Consequently, we can infer that the M2+−O1 bonds are weaker than the M2+−O3 bonds. The reduction in the C−O1 bond lengths in the complexes compared to the free ligand supports this fact. To justify our findings, the natures of bonds in the most stable M2+−MPA−MPAH(O,S,O,S) complexes were further substantiated by the AIM analysis presented in Table 3. The exploration of critical points for the interaction of ligands with M2+ ions reveals the existence of four (3, −1) critical points. The Laplacian of the charge density (∇2ρ) corresponding to each (3, −1) critical point for all three metal cations is positive, which indicates that the considered interactions are of closedshell type with depletion of energy. The ∇2ρ values at the two

basically charge-transfer bonds with pure electrostatic interactions. The other possible category of interaction of the metal cations with two ligands is one in which the metal cations coordinate with one fully deprotonated ligand (MPA) and one partially deprotonated ligand (MPAH). The possible coordination modes are presented in Figure 2, and the energetics of these complexes are collected in Table 5. From this table, it is evident that the tetracoordinated complexes for all of the metal cations [M2+−MPA−MPAH(O,S,O,S)] are more stable than the tricoordinated and bicoordinated ones. These tetracoordinated complexes have distorted tetrahedral geometries. The S− M2+−O3 angles have values of 105°, 98.72°, and 95.29° for zinc, cadmium, and mercury, respectively. The S−M2+−S1 angles are 129.73°, 137.41°, and 147.94° for zinc, cadmium, and mercury, respectively, whereas the O1−M2+−O3 angles are 98.30°, 102.71°, and 104.71° for the same metal ions. The S1−M2+− O1 angles have values of 87.74°, 81.68°, and 74.34° for Zn2+, Cd2+, and Hg2+, respectively. The M2+−S and M2+−O bond lengths are quite similar to those in other cyclic complexes, as is evident from Table S1 (Supporting Information). The tricoordinated species [M2+−MPA−MPAH(S,S,O)] were also found to be stable and to lie 7.1, 6.5, and 3.7 kcal mol−1 above the corresponding global minima for zinc, cadmium, and mercury cations, respectively. These small relative energy differences, which show their stability, are reported in Table 5. The bicoordinated complexes M2+−MPA−MPAH(S,S) were found to be the least stable species, and their relative energy differences were larger (40.7, 32.6, and 23.6 kcal mol−1 for Zn2+, Cd2+, and Hg2+, respectively) than their corresponding global minima. The behavior of the complexes in solution was similar, but the relative energy values were lower in magnitude than those in the gas phase. It is noteworthy that, for this category of complexes, the relative stability order clearly shows the cyclic sulfhydryl group to be the most favorable functional group for metalation. This is once again in concurrence with the theory of hard and soft acids and bases as stated earlier. Also, the number of members in the ringlike structures in the 1610



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critical points corresponding to the M2+−S and M2+−S1 bonds are nearly equal, but these values are smaller than the corresponding values for M2+−S bonds for the M2+−MPA(S,O) and M2+−MPAH(S,O)b complexes. Also, the ∇2ρ values at the (3, −1) critical points for the M2+−O1 bonds are smaller than those for the M2+−O3 bonds, which shows that the M2+− O3 bonds are stronger than the M2+−O1 bonds. These facts support the results obtained from the NBO analysis. The total energy densities, H(r), at all bond critical values corresponding to all bonds for the metal cations are negative, as is characteristic of medium interactions. The trends in the −G(r)/V(r) values at all of the (3, −1) critical points corresponding to all of the bonds reflect the fact that, even though the bonds are mainly electrostatic in nature, a weak covalent nature is also exhibited. The metal-ion affinities (MIAs) (ΔH298 K), free energies (ΔG298 K), and entropies (TΔS298 K) of the M2+−(MPA)2 and M2+−MPA−MPAH complexes in both gaseous and solution phases, reported in Table 6, also support the order of stability of all these complexes as discussed earlier. Judging from the values of the mentioned parameters, it can be stated that the formation of cyclic structures involving the sulfhydryl group upon metalation is preferred to the formation of complexes by metalation involving the carboxylic group. This inference is consistent with the theory of hard and soft acids and bases. The decrease in the MIA and free energy values of the complexes in the solution phase can be attributed to the substantial influence of the solvent (water) polarization on the energetics of the complexes. The MIA values reveal that two ligand molecules bind with the metal ions in the same order Zn2+ > Hg2+ > Cd2+, as in the case where one ligand molecule is coordinated with the metal cations. Proceeding as earlier, we attempted to simulate a system closer to the one in aqueous solution and added water molecules explicitly to complete the coordination spheres of the metal cations. The bare Cd2+ ion exhibits an octahedral geometry when it binds to water molecules,56 and the tetrahedral conformer is 60.9 kcal mol−1 less stable than the octahedral one. This energy difference decreases with increasing size of the ligand molecule, and for the [Cd− (H2O)5−CH3] complex, the tetrahedral form is 37.7 kcal mol−1 more stable than the octahedral form. Also, an EXAFS57 study revealed that a cadmium ion in aqueous solution can coordinate with a maximum of four thiols. Our attempt to microhydrate the most stable M 2+−(MPA)2 and M2+−MPA−MPAH complexes (for all three metal cations) with two water molecules (so that octahedral geometries could be obtained) produced certain structures (see Figure F2, Supporting Information) that revealed that the complexes are stabilized through hydrogen bonds and the water molecules are shifted out of the coordination spheres of the metal cations. Similar results were also obtained by Belcastro et al.27 in the case of the Cd2+ ion, from which we can confirm the preference for tetracoordination of Zn2+, Cd2+, and Hg2+ ions in both fully and partially deprotonated 3-mercaptopropionic acid. The metalion affinity (MIA) values for these microhydrated complexes are greater than those obtained in gas-phase calculations, which can be attributed to the formation of hydrogen bonds that stabilize the hydrated complexes. As discussed earlier, the apparent anomalies between the results obtained by the COSMO model and microhydration are due to the limitations of the COSMO model.

4. CONCLUSIONS In this article, the electronic structures of complexes formed by metalation of one and two partially or deprotonated 3mercaptopropionic acid molecules with zinc, cadmium, and mercury cations were investigated in the gas phase and in solution with water as the (polar) solvent. Among the complexes formed by the coordination of metal cations and one ligand molecule, cyclic structures involving bicoordination with the sulfhydryl group were the most stable species compared to the bicoordinated complexes involving the carboxylic group. Interestingly, monocoordinated complexes involving the sulfhydryl group were also not stable. In the case of complexes formed by the interaction of metal ions with two ligand molecules, the tetracoordinated complexes having distorted tetrahedral geometries and involving the sulfhydryl groups were the most stable complexes compared to the tricoordinated or bicoordinated linear structures. These findings are well harmonized with the theory of hard and soft acids and bases. The metal-ion affinities (MIAs), free energies, and entropies also support these facts. The MIA values also reveal that, irrespective of the coordination of one or two ligand molecules with the metal cations, both the MPA and MPAH molecules bind with the metal cations in the order Zn2+ > Hg2+ > Cd2+. The nature of coordination of these complexes remained essentially the same when they were considered in aqueous environment, that is, when a polarizing field was applied to the free cations. Microhydration of the metal cations by explicit addition of water molecules showed clearly that tetrahedral coordination is preferred to octahedral coordination. Particular emphasis was also paid to the critical analysis of the interactions present in the studied molecular systems. NBO analysis was exploited to elucidate some crucial facets of the bonding interactions that prevail in the most stable complexes. The results illustrated that the interactions are mainly electrostatic in nature, arising from charge transfer from the atoms of the ligand molecules to the metal center, although a weak covalent nature of a few bonds was established during further analysis of bonding using AIM theory. Our results also show the influence of the solvent (water) reaction field on the nature of bonding in the complexes.

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ASSOCIATED CONTENT

S Supporting Information *

Bond lengths for some specific bonds of free ligands and the most stable M2+−MPA and M2+−MPAH, M2+−(MPA)2 and M2+−MPA−MPAH complexes (M = Zn, Cd, Hg) calculated at the B3LYP/BS level (Table S1). Enthalpy (ΔH =MIA) variations for the formation process of the most stable microhydrated M2+−MPA and M2+−MPAH (M = Zn, Cd, Hg) complexes computed at the B3LYP/BS level (Table S2). Results of NBO analysis of the most stable M2+−MPA, M2+− MPAH, M2+−(MPA)2 and M2+−MPA−MPAH complexes in the solution phase (Table S3). This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: +91-33-2473-4971 (ext. 1257), Fax: 91-33-24732805. Notes

The authors declare no competing financial interest. 1611



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(22) Moreau, J. W.; Weber, P. K.; Martin, M. C.; Gilbert, B.; Hutcheon, I. D.; Banfield, J. F. Extracellular Proteins Limit the Dispersal of Biogenic Nanoparticles. Science 2007, 316, 1600−1603. (23) Xia, T.; Kovochich, M.; Liong, M.; Mädler, L.; Gilbert, B.; Shi, H.; Yeh, J. I.; Zink, J. I.; Nel, A. E. Comparison of the Mechanism of Toxicity of Zinc Oxide and Cerium Oxide Nanoparticles Based on Dissolution and Oxidative Stress Properties. ACS Nano 2008, 2 (10), 2121−2134. (24) Mopper, K.; Taylor, B. F. Biogeochemical Cycling of Sulfur Thiols in Coastal Marine Sediments. ACS Symp. Ser. 1986, 305, 324− 339. (25) Vairavamurthy, M. A.; Manowitz, B.; Maletic, D.; Wolfe, H. Interactions of Thiols with Sedimentary Particulate Phase: Studies of 3-Mercaptopropionate in Salt Marsh sediments from Shelter Island, New York. Org. Geochem. 1997, 26, 577−585. (26) Vairavamurthy, M. A.; Goldenberg, W. S.; Ouyang, S.; Khalid, S. The Interaction of Hydrophilic Thiols with Cadmium: Investigation with a Simple Model, 3-Mercaptopropionic Acid. Mar. Chem. 2000, 70, 181−189. (27) Belcastro, M.; Marino, T.; Russo, N.; Sicilia, E. Structure and

ACKNOWLEDGMENTS S.B. gratefully acknowledges the authority of Netaji Subhash Engineering College, Garia, Kolkata, India, for allowing him to carry out research work at IACS. D.M. is grateful to the Council of Scientific and Industrial Research (CSIR), Government of India, for providing him a Senior Research Fellowship (SRF). Thanks are due to the reviewers for their constructive comments and suggestions to improve the manuscript.

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REFERENCES

(1) Glusker, J. P. Structural Aspects of Metal Liganding to Functional Groups in Proteins. Adv. Protein Chem. 1991, 42, 1−76. (2) Waalkes, M. P., Oberdorster, G., Foulkes, E. C., Eds. Biological Effects of Heavy Metals; CRC Press: Boca Raton, FL., 1990, 2, 129. (3) Waalkes, M. P.; Rehm, S. Chronic Toxic and Carcinogenic Effects of Cadmium Chloride in Male DBA/2NCr and NFS/NCr Mice: Strain-Dependent Association with Tumors of the Hematopoietic System, Injection Site, Liver, and Lung. Fundam. Appl. Toxicol. 1994, 23, 21−31. (4) Beryllium, Cadmium and Mercury Exposures in the Glass Manufacturing Industry; International Agency for Research on Cancer (IARC): Lyon, France, 1993; Vol. 58, p 119. (5) Sanità di Toppi, L.; Gabbrielli, R. Effects of γ-irradiation on Elongation and Indole-3-Acetic Acid Level of Maize (Zea mays) Coleoptiles. Environ. Exp. Bot. 1999, 41, 105−130. (6) Wagner, G. Accumulation of Cadmium in Crop Plants and Its Consequences to Human Health. J. Adv. Agron. 1993, 51, 173−212. (7) Cabrera, C.; Ortega, E.; Lorenzo, M. L.; Lopez, M. Rev. Environ. Contam. Toxicol. 1998, 154, 55−81. (8) Duxbury, T. Adv. Microb. Ecol. 1985, 8, 185−236. (9) Nriagu, J. O. In Cadmium in the Environment. Part 1: Ecological Cycling; Nriagu, J. O., Ed.; Wiley: New York, 1980; p 35. (10) Morel, F. M. M.; Morel-Laurens, N. M. L. Trace metals in seawater. NATO Conf. Ser. 4: Mar. Sci. 1983, 9, 841−869. (11) Brand, L. E.; Sunda, W. G.; Guillard, R. R. L. Reduction of Marine Phytoplankton Reproduction Rates by Copper and Cadmium. J. Exp. Mar. Biol. Ecol. 1986, 96, 225−250. (12) Ahner, B. A.; Price, N. M.; Morel, F. M. M. Phytochelatin Production by Marine Phytoplankton at Low Free Metal Ion Concentrations: Laboratory Studies and Field Data from Massachusetts Bay. Proc. Natl. Acad. Sci. U.S.A. 1994, 91, 8433−8436. (13) Lee, J. G.; Beth, B. A.; Morel, F. M. M. Export of Cadmium and Phytochelatin by the Marine Diatom Thalassiosira Weissflogii. Environ. Sci. Technol. 1996, 30, 1814−1821. (14) Hare, L.; Tessier, A. Predicting Animal Cadmium Concentrations in Lakes. Nature 1996, 380, 430−432. (15) Sakamoto-Arnold, C. M.; Hanson, A. K., Jr.; Huizenga, D. L.; Kester, D. R. Spatial and Temporal Variability of Cadmium in Gulf Stream Warm Core Rings and Associated Waters. J. Mar. Res. 1987, 45, 201−230. (16) Muller, F. L.; Kester, D. R. Kinetic Approach to Trace Metal Complexation in Seawater: Application to Zinc and Cadmium. Environ. Sci. Technol. 1990, 24, 234−242. (17) Bruland, K. W. Complexation of Cadmium by Natural Organic Ligands in the Central North Pacific. Limnol. Oceanogr. 1992, 37, 1008−1017. (18) Bhattacharya, S. K.; Madura, R. L.; Uberoi, V.; Haghighi-Podeh, M. R. Toxic effects of cadmium on methanogenic systems. Water Res. 1995, 29, 2339−2345. (19) Donat, J. R. Presented at the ACS Spring Meeting, Anaheim, CA, Apr 2−7, 1995. (20) Muller, F. L. L. Interactions of Copper, Lead and Ccadmium with the Dissolved, Colloidal and Particulate Components of Estuarine and Coastal Waters. Mar. Chem. 1996, 52, 245−268. (21) Toxicological Profile for Zinc. Toxic Substances Portal; Agency for Toxic Substances and Disease Registry (ATSDR), U.S. Department of Health and Human Services: Washington, D.C., 2005.

+

Coordination Modes in the Interaction between Cd2 and 3Mercaptopropionic Acid. J. Phys. Chem. A 2004, 108, 8407−8410. (28) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Scalmani, G.; Barone, V.; Mennucci, B.; Petersson, G. A.; Nakatsuji, H.; Caricato, M.; Li, X.; Hratchian, H. P.; Izmaylov, A. F.; Bloino, J.; Zheng, G.; Sonnenberg, J. L.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Vreven, T.; Montgomery, J. A., Jr.; Peralta, J. E.; Ogliaro, F.; Bearpark, M.; Heyd, J. J.; Brothers, E.; Kudin, K. N.; Staroverov, V. N.; Keith, T.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, J. M.; Klene, M.; Knox, J. E.; Cross, J. B.; Bakken, V.; Adamo, C.; Jaramillo, J.; Gomperts, R.; Stratmann, R. E.; Yazyev, O.; Austin, A. J.; Cammi, R.; Pomelli, C.; Ochterski, J. W.; Martin, R. L.; Morokuma, K.; Zakrzewski, V. G.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Dapprich, S.; Daniels, A. D.; Farkas, Ö .; Foresman, J. B.; Ortiz, J. V.; Cioslowski, J.; Fox, D. J. Gaussian 09, revision B.01; Gaussian, Inc.: Pittsburgh, PA, 2009. (29) Becke, A. D. Density-Functional Thermochemistry. III. The Role of Exact Exchange. J. Chem. Phys. 1993, 98, 5648−5652. (30) Lee, C.; Yang, W.; Parr, R. G. Development of the Colic-Salvetti Correlation-Energy Formula into a Functional of the Electron Density. Phys. Rev. B 1988, 37, 785−789. (31) Zhu, W.; Luo, X.; Mok Puah, C.; Tan, X.; Shen, J.; Gu, J.; Chen, K.; Jiang, H. The Multiplicity, Strength, and Nature of the Interaction of Nucleobases with Alkaline and Alkaline Earth Metal Cations: A Density Functional Theory Investigation. J. Phys. Chem. A 2004, 108, 4008−4018. (32) Krishnan, R.; Binkley, J. S.; Seeger, R.; Pople, J. A. SelfConsistent Molecular Orbital Methods. XX. A Basis Set for Correlated Wave Functions. J. Chem. Phys. 1980, 72, 650−654. (33) McLean, A. D.; Chandler, G. S. Contracted Gaussian Basis Sets for Molecular Calculations. I. Second Row Atoms, Z = 11−18. J. Chem. Phys. 1980, 72, 5639−5648. (34) Hay, P. J.; Wadt, W. R. Ab Initio Effective Core Potentials for Molecular Calculations. Potentials for the Transition Metal Atoms Sc to Hg. J. Chem. Phys. 1985, 82, 270−283. (35) Wadt, W. R.; Hay, P. J. Ab Initio Effective Core Potentials for Molecular Calculations. Potentials for Main Group Elements Na to Bi. J. Chem. Phys. 1985, 82, 284−298. (36) Hay, P. J.; Wadt, W. R. Ab Initio Effective Core Potentials for Molecular Calculations. Potentials for K to Au Including the Outermost Core Orbitals. J. Chem. Phys. 1985, 82, 299−310. (37) Russo, N.; Toscano, M.; Grand, A. Lithium Affinity for DNA and RNA Nucleobases. The Role of Theoretical Information in the Elucidation of the Mass Spectrometry Data. J. Phys. Chem. B 2001, 105, 4735−4741. (38) Rodgers, M. T.; Armentrout, P. B. Noncovalent Interactions of Nucleic Acid Bases (Uracil, Thymine, and Adenine) with Alkali Metal 1612



Article

Ions. Threshold Collision-Induced Dissociation and Theoretical Studies. J. Am. Chem. Soc. 2000, 122, 8548−8558 and references therein.. (39) Hoyau, S.; Ohanessian, G. Interaction of Alkali Metal Cations (Li+−Cs+) with Glycine in the Gas Phase: A Theoretical Study. Chem.Eur. J. 1998, 4, 1561−1569. (40) Marino, T.; Russo, N.; Toscano, M. Gas-Phase Metal Ion (Li+, Na+, Cu+) Affinities of Glycine and Alanine. J. Inorg. Biochem. 2000, 79, 179−185. (41) Burda, J. V.; Sÿponer, J.; Hobza, P. Ab Initio Study of the Interaction of Guanine and Adenine with Various Mono- and Bivalent Metal Cations (Li+, Na+, K+, Rb+, Cs+; Cu+, Ag+, Au+ ; Mg2+, Ca2+, Sr2+, Ba2+; Zn2+, Cd2+, and Hg2+). J. Phys. Chem. 1996, 100, 7250− 7255. (42) Boys, S. F.; Bernardi, F. The Calculation of Small Molecular Interactions by the Differences of Separate Total Energies. Some Procedures with Reduced Errors. Mol. Phys. 1970, 19, 553−566. (43) Khaliullin, R.; Bell, A. T.; Head-Gordon, M. Electron Donation in the Water−Water Hydrogen Bond. Chem.Eur. J. 2009, 15, 851− 855. (44) Klamt, A.; Schuurmann, G. COSMO: A New Approach to Dielectric Screening in Solvents with Explicit Expressions for the Screening Energy and its Gradient. J. Chem. Soc., Perkin Trans. 1993, 2, 799−805. (45) Reed, A. E.; Curtiss, L. A.; Weinhold, F. Intermolecular Interactions from a Natural Bond Orbital, Donor−Acceptor Viewpoint. Chem. Rev. 1988, 88, 899−926. (46) Glendening, D. E.; Reed, A. E.; Carpenter, J. E.; Winhold, F. NBO, version 3.1; Gaussian, Inc.: Pittsburgh, PA, 1992. (47) Bader, R. F. W. Atoms in Molecules. A Quantum Theory; Oxford University Press: New York, 1994. (48) Cioslowski, J.; Nanayakkara, A.; Challacombe, M. Rapid Evaluation of Atomic Properties with Mixed Analytical/Numerical Integration. Chem. Phys. Lett. 1993, 203, 137−142. (49) Cioslowski, J.; Nanayakkara, A. A New Robust Algorithm for Fully Automated Determination of Attractor Interaction Lines in Molecules. Chem. Phys. Lett. 1994, 219, 151−154. (50) Bader, R. F. W. A Bond Path: A Universal Indicator of Bonded Interactions. J. Phys. Chem. A 1998, 102, 7314−7323. (51) Abboud, J. L. M.; Mo, O.; de Paz, J. L. G.; Yanez, M.; Esseffar, M.; Bouab, W.; El-Mouhtadi, M.; Mokhlisse, R.; Ballesteros, E. Thiocarbony Versus Carbonyl Compounds: A Comparison of Intrinsic Reactivities. J. Am. Chem. Soc. 1993, 115, 12468−12476. (52) Hocquet, A. Intramolecular Hydrogen Bonding in 2′Deoxyribonucleosides: An AIM Topological Study of the Electronic Density. Phys. Chem. Chem. Phys. 2001, 3, 3192−3199. (53) Rozas, I.; Alkorta, I.; Elguero, J. Behavior of Ylides Containing N, O, and C Atoms as Hydrogen Bond Acceptors. J. Am. Chem. Soc. 2000, 122, 11154−11161. (54) Ortiz Alba, J. C.; Jané, C. B. XAIM, version 1.0; Universitat Roviral Virgili, Tarragona, Spain; available at http://www.quimica.urv. es/XAIM, 1998. Accessed September, 2012. (55) Pearson, G. R. Hard and Soft Acids and Bases. J. Am. Chem. Soc. 1963, 85, 3533−3539. (56) Rulíšek, L.; Havlas, Z. Theoretical Studies of Metal Ion Selectivity. 1. DFT Calculations of Interaction Energies of Amino Acid Side Chains with Selected Transition Metal Ions (Co2+, Ni 2+, Cu2+, Zn2+, Cd2+, and Hg2+). J. Am. Chem. Soc. 2000, 122, 10428−10439. (57) Frenkel, A. I.; Vairavamurthy, M.; Newville, M. A Study of the Coordination Environment in Aqueous Cadmium−Thiol Complexes by EXAFS Spectroscopy: Experimental vs Theoretical Standards. J. Synchrotron Radiat. 2001, 8, 669−671.

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Interaction Between Group IIb Divalent Transition-Metal Cations and 3

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