Dithiolate Ligands - American Chemical Society

Article pubs.acs.org/Organometallics

Redox Reactions of Nickel, Copper, and Cobalt Complexes with “Noninnocent” Dithiolate Ligands: Combined in Situ Spectroelectrochemical and Theoretical Study Peter Machata,†,‡ Peter Herich,† Karol Lušpai,† Lukas Bucinsky,† Stanislava Šoralová,§ Martin Breza,† Jozef Kozisek,† and Peter Rapta*,†,‡ †

Institute of Physical Chemistry and Chemical Physics, Faculty of Chemical and Food Technology, Slovak University of Technology in Bratislava, Radlinského 9, SK-812 37 Bratislava, Slovak Republic ‡ Center of Spectroelectrochemistry, Leibniz Institute for Solid State and Materials Research, Helmholtzstrasse 20, D-01069 Dresden, Germany § Department of Pharmaceutical Chemistry, Faculty of Pharmacy, Comenius University in Bratislava, Odbojárov 10, SK-832 32 Bratislava, Slovak Republic S Supporting Information *

ABSTRACT: The redox properties of copper, nickel, and cobalt complexes (MePh3P)[M(bdt)2] with the ligand benzene-1,2-dithiolate (bdt) and synthesized complexes (MePh3P)[M(bdtCl2)2] with the ligand 3,6-dichlorobenzene1,2-dithiolate (bdtCl2) have been studied by cyclic voltammetry and in situ EPR−UV/vis/NIR spectroelectrochemistry. The addition of chlorine substituents to the 3- and 6-positions of the benzene ring not only facilitates the reduction of [M(bdtCl2)2]− complexes but also leads to the remarkable stabilization of [M(bdtCl2)2]2− dianions in solution. In contrast to the EPR-silent copper complexes, the solutions of nickel samples exhibit a broad singlet EPR signal at room temperature which becomes anisotropic at 100 K with a characteristic rhombic pattern. Cathodic reduction of copper and cobalt complexes leads to paramagnetic species having an EPR signal with splitting from 63,65Cu for copper and from 59Co for cobalt samples, confirming a strong contribution of the central atom with substantial delocalization of the unpaired spin onto the central atom. B3LYP/6-311g*/pcm calculations of the monoanions as well as of their oxidized and reduced forms were performed. The spin density of all open-shell ground states calculated for the investigated complexes in different redox states corresponds well to the experimental spectroelectrochemical data.

■

INTRODUCTION Dithiolate complexes have attracted considerable attention because of their technical applications as superconductors,1,2 pesticides,3,4 resins, Q-switching dyes for IR spectroscopy,5,6 compounds with unusual magnetic properties,7 and biocatalysts.8,9 Dithiolate ligands, according to the Jørgensen classification scheme,10 are usually classified as “noninnocent” ligands, where these ligands are bonded with the central atom in such a manner that the distribution of the electron density between the central atom and the ligand is unclear. Consequently, there is a problem in determining the oxidation state of the central atom. In the case of “noninnocent” ligands the real physical (or spectroscopic) oxidation number differs from the formal one.11,12 The bis-dithiolate complexes are known to be in a variety of oxidation states which are largely determined by the nature of the ligand and by its variable contribution to the redox processes.13 Extended charge delocalization is typical for these complexes in different oxidation states. This delocaliza© XXXX American Chemical Society

tion makes it difficult to determine whether a metal-centered conversion between M(II) and M(III) or a ligand-centered redox process takes place. One example of such a structure is the (MePh3P)[Ni(bdt)2] complex14 with benzene-1,2-dithiolate ligand (bdt)2−, the redox properties of which are the subject of the present work. On the other hand, for similar square-planar complexes containing two bidentate O,O-, N,N-, and O,N-coordinated ligands, ligand-based redox processes leading to radical ligands with strong intramolecular antiferromagnetic coupling were reported,15 which has been explored for bdt and bdt-like ligands (bdtX) in [M(bdtX)2]q− complexes (M = Fe, Co, Ni, Cu, Pd, Pt, Au) by Neese, Wieghardt, and co-workers.16 Special Issue: Organometallic Electrochemistry Received: January 23, 2014

A

dx.doi.org/10.1021/om5000584 | Organometallics XXXX, XXX, XXX−XXX

Organometallics

■

Article

RESULTS AND DISCUSSION Cyclic Voltammetry. The redox properties of [M(bdt)2]− and [M(bdtCl2)2]− complex anions in solution were investigated by cyclic voltammetry in dichloromethane and N,Ndimethylformamide (DMF) containing tetrabutylammonium hexafluorophosphate (TBAPF6) as the supporting electrolyte. The resulting electrochemical data obtained in dichloromethane solutions are summarized in Table 1. In the cathodic

Results of the X-ray analysis of (MePh3P)[M(bdt)2] (M = Cu, Ni, Co) showed that the M−S bond lengths are shorter than the usual published values for M(II) complexes, supporting the assumption that the formal oxidation state in these structures is M(III).17 The results of UV/vis spectroscopy of nickel, copper, and cobalt (MePh3P)[M(bdt)2] complexes show absorption bands which are significantly different from those of typical M(II) complexes: in particular, for the (MePh3P)[Ni(bdt)2] complex the results of UV/vis and EPR spectroscopy correspond well to similar data reported for Ni(III) complexes.17 Sellmann et al. reported18 three different oxidation forms of nickel complexes coordinated by 3,5-di-tert-arylbutyl-1,2benzenedithiolate with the formal oxidation states of the central atom being Ni(II), Ni(III), and Ni(IV). The Ni−S bond length obtained by X-ray analysis indicated that the redox processes between these oxidation states take place predominantly within the central [NiS4] part. Recently, [M(bdt)2]q complexes of copper, nickel, and cobalt with benzene-1,2dithiolate ligands in different charge states (q = −1, −2, −3) were studied also by quantum chemical calculations, including atomic charges and d-electron populations.19 The real oxidation state of the central atom varies between M(III) and M(II) for [M(bdt)2]− and [M(bdt)2]2−. On the other hand, the spin density of paramagnetic species is centered mainly on the central atom. Although the redox processes of dithiolate complexes represented a subject of numerous studies, there is still a lack of experimental results in the literature characterizing a detailed structure of their different charge states, particularly those formed upon cathodic reduction. In the present work the redox properties of copper, nickel, and cobalt complexes [M(bdt)2]− with the benzene-1,2-dithiolate ligand and newly prepared [M(bdtCl2)2]− complexes (Scheme 1) with the 3,6-dichlor-

Table 1. Comparison of the Electrode Potentials vs Fc+/Fc in TBAPF6/CH2Cl2 and Absorption Maxima of Initial ([M(bdt)2]− and [M(bdtCl2)2]−) and Reduced Forms (red, [M(bdt)2]2− and [M(bdtCl2)2]2−) of the Investigated Complexes in CH2Cl2 E1/2red vs Fc+/Fc, V

Ep/2ox vs Fc+/Fc, V

[Cu(bdt)2]− [Ni(bdt)2]−

−1.02 −0.95

0.31 −0.03

[Co(bdt)2]−

−1.27

0.03

[Cu(bdtCl2)2]− [Ni(bdtCl2)2]− [Co(bdtCl2)2]−

−0.77 −0.65 −0.97

0.60 0.22 0.02

complex

a

λmax, nm λmax(red), nm 397a 362, 876a 361, 657a 324, 403 369, 856 369, 669

341, 458a 402a

339, 453 399

Measured in DMF.

region at a scan rate of 100 mV s−1 all investigated compounds give reversible cyclic voltammograms (Figure 1). This indicates

Scheme 1. Structure of the Investigated Complexes (a) [M(bdt)2]− and (b) [M(bdtCl2)2]−, where M = Cu, Ni, Co

Figure 1. Cyclic voltammograms of (a) [M(bdt)2]− and (b) [M(bdtCl2)2]−, where M = Cu, Ni, Co, for oxidation and reduction in TBAPF6/CH2Cl2 (scan rate v = 100 mV s−1).

the chemical stability of generated species (chemical reversibility) as well as a fast electron transfer (electrochemical reversibility) at the first cathodic peak, except for [Cu(bdt)2]−, as will be discussed below in more detail. The values of reduction potentials vs Fc+/Fc depend on the central atom of the studied complexes. In the case of [M(bdt)2]− samples the easiest reduction was found for the [Ni(bdt)2]− complex with E1/2 = −0.95 V vs Fc+/Fc. A more negative reduction potential with a potential shift of −70 mV was observed for investigated copper complex with E1/2 = −1.02 V vs Fc+/Fc. An even stronger potential shift was found for the cobalt complex [Co(bdt)2]− with the half-wave potential E1/2 = −1.27 V vs Fc+/Fc. The potential shifts observed in our work are in good agreement with already reported electrochemical data for the same structures measured

obenzene-1,2-dithiolate ligand (bdtCl2)2− are studied by cyclic voltammetry and in situ EPR−UV/vis/NIR spectroelectrochemistry. Special attention is focused on the characterization of the redox site for both oxidation and reduction of the complexes (MePh3P)[M(bdt)2] and (MePh3P)[M(bdtCl2)2] in different organic solvents. Finally, quantum-chemical calculations of investigated complexes at the B3LYP/6-311g* level of theory in different charge and spin states allow us to interpret the experimental spectroelectrochemical results. B


Organometallics

Article

in DMF solutions using polarography.20 Small differences between the reported shifts and those observed in our experiments could be caused by other experimental conditions during the measurements, including solvents and electrochemical techniques. A significant change in the redox potentials for complexes with different central atoms indicates a strong contribution of the metal ion to the redox processes of the investigated samples. The incorporation of two chlorine substituents to the 3- and 6- positions of the benzene ring results in the shift of the reduction potential to the anodic region. A similar effect was observed for other electron-accepting groups attached to the benzene ring of the dithiolate ligand.13,19 The difference between the reduction potentials of [M(bdt)2]− type complexes in comparison to the [M(bdtCl2)2]− samples is about 300 mV for cobalt and nickel complexes and 250 mV for complexes with a copper central atom. Interestingly, a more reversible behavior for the first reduction step was found for new [M(bdtCl2)2]− complexes, as was proved by cyclic voltammetry at different scan rates (see Figures S1 and S2 in the Supporting Information). Even at the slow scan rate of 5 mV s−1 cyclic voltammograms with a ratio of the peak currents (Ipa/Ipc) close to 1 were observed, in contrast to the case for [M(bdt)2]− complexes, confirming the high chemical stability of [M(bdtCl2)2]2− dianion states. Additionally, when the scan rate is decreased, the difference in the anodic and cathodic peak potentials (peak-to-peak separation ΔEp) substantially decreases for [Cu(bdt)2]−, while for [Cu(bdtCl2)2]− a nearly constant ΔEp (a slight increase of the peak-to-peak separation at higher scan rates is due to iR drop) was found for all tested scan rates (Figures S1a and S2a). This indicates a slow heterogeneous electron transfer for [Cu(bdt)2]− reduction in comparison to the [Cu(bdtCl2)2]− reduction (see also the red lines in Figure 1). The more chemically stable [M(bdt)2]2− dianions generated were found using DMF solvent instead of CH2Cl2, as illustrated for [Ni(bdt)2]− in Figures S3 and S1b (Supporting Information). Therefore, in further spectroelectrochemical studies of [M(bdt)2]− anions DMF was used as the main solvent due to the higher stability of the reduced states for the investigated samples. In contrast to the cathodic reduction a more complex redox behavior was observed upon the anodic oxidation (Figure 1). The behavior of metallocomplexes with benzenedithiolate ligand in the anodic region is quite different. For the [Ni(bdt)2]− complex the voltammogram exhibits a redox couple at Ep/2 = −0.03 V vs Fc+/Fc (Ipc/Ipa = 0.5 at v = 100 mV s−1), with a strongly shifted oxidation potential of 280 mV in comparison to that for [Cu(bdt)2]−. For the complex with a copper central atom an irreversible oxidation peak was observed at the half-peak potential Ep/2 = 0.31 V vs Fc+/Fc. The cobalt complex [Co(bdt)2]− is oxidized at the lower oxidation potential Ep/2 = 0.03 V vs Fc+/Fc, being similar to the observed potential for a complex with a nickel central atom, but the oxidation is fully irreversible even at higher scan rates, confirming low chemical stability of oxidized [M(bdt)2]−. [M(bdtCl2)2]− redox behavior in the anodic region differs from that of [M(bdt)2]−. The oxidation of the complex with a copper central atom is chemically irreversible with the highest potential: Ep/2 = 0.60 V vs Fc+/Fc. For the complex with a nickel central atom, the introduction of chlorine substituents significantly decreases the ratio of the peak currents (Ipc/Ica) of the first anodic peak. In contrast to the case for [Ni(bdt)2]−, the oxidation process for [Ni(bdtCl2)2]− is completely

chemically irreversible and the oxidation of this complex is more difficult (Ep/2 = 0.22 V vs Fc+/Fc). As already stated in the reduction studies of the investigated samples, the introduction of chlorine substituents facilitates their reduction. As expected, a shift of oxidation potentials in the anodic region for the samples with a dichlorobenzenedithiolate ligand was observed, except for the [Co(bdtCl2)2]− sample. For this sample the peak current in the voltammetric back scan was higher than the peak current measured in the forward scan (Figure 1). Interestingly, the lowest shift of the oxidation potential was found (Ep/2 = 0.02 V vs Fc+/Fc). At lower scan rates (e.g., 5 mV s−1) the peak in the back scan is separated into two voltammetric peaks, indicating a complex redox mechanism with follow-up products (Figure S4, Supporting Information). After this process a thin blue film on the electrode was observed. The redox potentials summarized in Table 1 enable us to estimate the electrochemical energy gap (ΔEgap,ec) for all of the investigated complexes, which is obtained as the difference between the formal potential of the first oxidation step and the formal potential of the first reduction step (which are roughly equal to the corresponding E1/2 values). The largest similar ΔEgap,ec values of 1.33 and 1.37 V were determined for [Cu(bdt)2]− and [Cu(bdtCl2)2]−, respectively. A much smaller electrochemical gap was found for nickel complexes, namely 0.92 V for [Ni(bdt)2]− and 0.87 V for [Ni(bdtCl2)2]−. The electrochemical gap calculated for [Co(bdtCl2)2]− is 0.99 V, which is much smaller than ΔEgap,ec = 1.30 V for [Co(bdt)2]−. Generally, the first reduction potential of the cobalt complexes is more negative than that of the nickel and copper samples, suggesting their weaker electron-accepting ability. The differences in the electrochemical energy gaps as well as in the redox potentials for both reduction and oxidation for all investigated complexes will be discussed in detail concerning our quantum chemical calculations in a separate paragraph below. In Situ EPR−UV/Vis/NIR Spectroelectrochemistry. For monoanions of the investigated complexes with a copper central atom no EPR signal was observed at room and low temperatures. During the in situ study of the electroreduction of [Cu(bdt)2]− by EPR−UV/vis/NIR spectroelectrochemistry at a platinum-mesh electrode (scan rate v = 5 mV s−1) in DMF an electrochemically quasi-reversible process was observed at the first cathodic peak (inset in Figure 2a). The larger peak-topeak separation ΔEp of the in situ cyclic voltammogram is caused by the higher iR drop in the flat EPR-spectroelectrochemical cell. In the UV/vis region the absorption of [Cu(bdt)2]− at 397 nm decreases and new optical bands at 341 and 458 nm appear with isosbestic points at 353 and 437 nm in the forward scan (Figure 2b). Simultaneously, a new welldefined four-component room-temperature EPR signal (due to the nuclear spin I = 3/2 for 63,65Cu) with the hyperfine splitting from the Cu atom (g = 2.047) arises (inset in Figure 2c). This indicates that a large part of the unpaired spin density is localized over the central part of the molecule. Interestingly, the corresponding EPR signal can be simulated by considering a slightly different hyperfine splitting constant for each copper isotope (a(63Cu) = 77 G and a(65Cu) = 81 G). A characteristic mI -dependent line width is also visible in the corresponding EPR spectrum and had to be taken into account by its simulation (Figure S5, Supporting Information). The evolution of the EPR signal intensity (Figure 2c) corresponds well to the evolution of a new optical band at 458 nm (see Figure 2a). In dichloromethane a similar behavior was C


Organometallics

Article

Figure 3. In situ EPR−UV/vis/NIR spectroelectrochemistry of [Ni(bdt)2]−: (a) potential dependence of UV/vis/NIR spectra with corresponding cyclic voltammogram (0.1 M TBAPF6/DMF, v = 5 mV s−1); (b) evolution of UV/vis/NIR spectra in 2D projection; (c) potential dependence of double-integrated EPR intensity with corresponding EPR spectrum.

Figure 2. In situ EPR−UV/vis/NIR spectroelectrochemistry of [Cu(bdt)2]−: (a) potential dependence of UV/vis spectra with corresponding cyclic voltammogram (0.1 M TBAPF6/DMF, scan rate v = 5 mV s−1); (b) evolution of UV/vis spectra in 2D projection; (c) potential dependence of double -integrated EPR intensity with the corresponding EPR spectrum.

the complexes with a nickel central atom exhibit well-defined broad singlet EPR signals. For [Ni(bdt)2]− the g factor is 2.0836 and the line width ΔBpp is 14.56 G. Only a small decrease in the corresponding EPR intensity was observed upon reduction (Figure 3c). It should be noted that for initially paramagnetic compounds in solution it is more challenging to use in situ ESR spectroelectrochemistry, since a strong EPR signal is found in solution and remains to a certain extent during the whole spectroelectrochemical experiment because the bulk solution far from the electrode (but still in the EPR cavity) contains nonreacted paramagnetic species, resulting in an EPR signal. A marked decrease of the observed EPR signal for [Ni(bdt)2]− at the first reduction peak was confirmed unambiguously using a large Pt-mesh electrode under chronopotentiostatic conditions (Figure S7, Supporting Information). Analogously to [Ni(bdt)2]− in DMF, similar redox behavior in the region of the first cathodic peak is exhibited by [Ni(bdtCl2)2]− in CH2Cl2 (Figure S8, Supporting Information). Before reduction a similar broad single EPR line with ΔBPP = 15.76 G and a g factor value of 2.0888 was identified.

observed (not shown) but the generated dianions are less stable, as already indicated in cyclic voltammetric studies. Analogously, a very similar behavior at the first cathodic peak is exhibited by [Cu(bdtCl2)2]−. Since chlorine substituents significantly improve both the electrochemical and chemical reversibility for the cathodic reduction, well-defined optical spectra of [Cu(bdtCl2)2]2− were recorded also in dichloromethane (Figure S6, Supporting Information). Here, the absorption maxima of the monoanion [Cu(bdtCl2)2]− are at 324 and 403 nm and new absorption bands arise at 339 and 453 nm for the reduced state with isosbestic points at 356 and 441 nm (see Figure S6 and Table 1). The corresponding EPR spectra are more broadened and less resolved than in the EPR spectrum of [Cu(bdt)2]2−. At the first reduction peak of [Ni(bdt)2]− in DMF (inset in Figure 3a), optical absorption of the dominating band at 876 nm and the absorption at 362 nm decrease and a new optical transition at 402 nm arises via the isosbestic point at 388 nm (Figure 3b). In contrast to the case for the copper complexes D


Organometallics

Article

For the [Co(bdt)2]− complex a reversible cathodic reduction (the ratio of cathodic and anodic currents in the corresponding cyclic voltammogram is close to 1) was observed even at a scan rate of 5 mV s−1 in DMF, as shown in Figure 4a. For both

fast chemical reactions of their mono-oxidized forms in solution (not shown). However, for complexes [Cu(bdtCl2)2]− and [Co(bdtCl2)2]− more reversible spectroelectrochemical behavior was found. Anodic oxidation of [Cu(bdtCl2)2]− at the first anodic peak is accompanied by a decrease of the anion band at 404 nm and an increase of the new optical band at 485 nm (Figure S10a, Supporting Information). The potential evolution of UV/vis/NIR spectra at the first anodic peak in CH2Cl2 confirmed the low stability of [Cu(bdtCl2)2]0, because only a partial recovery of the optical bands of the corresponding monoanion [Cu(bdtCl2)2]− was observed during the voltammetric back scan (Figure S10b). Poor electrochemical reversibility with a large peak−peak separation accompanied by only partial rereduction of the oxidized species indicates both slow electron transfer and significant structural differences between the [Cu(bdtCl2)2]− and [Cu(bdtCl2)2]0 redox states. In the case of the [Co(bdtCl2)2]− oxidation, which is accompanied by a new broad band in the NIR region, an even higher recovery of the initial optical bands at 368 and 668 nm was observed in the back scan (Figure S11b, Supporting Information). At room temperature no new EPR spectra were observed upon oxidation of all investigated complexes. Only in the case of nickel complexes there was an irreversible decrease of the EPR signal visible during oxidation at the first anodic peak. EPR Measurements at Low Temperatures. A rhombic pattern in EPR spectra was observed for the frozen DMF solutions of Ni complexes at 100 K, as shown in Figure 5.

Figure 4. In situ UV/vis/NIR spectroelectrochemistry of [Co(bdt)2]−: (a) cyclic voltammogram (0.1 M TBAPF6/DMF, scan rate v = 5 mV s−1, platinum-mesh working electrode); (b) potential dependence of UV/vis spectra; (c) evolution of UV/vis spectra in 2D projection.

cobalt complexes studied, no EPR signal in the X-band region was detected even at low temperatures down to 77 K. During the in situ reduction at the first cathodic peak in DMF the dominating absorption band at 657 nm decreases via an isosbestic point at 459 nm. Simultaneously, a new band in the 350−450 nm region appears which is overlapped with a decrease of the monoanion optical band at 361 nm (Figure 4c). Additionally, during the voltammetric back scan a complete recovery of the monoanion [Co(bdt)2]− optical bands was observed similarly to [Cu(bdtCl2)2]− (see Figure S6a, Supporting Information). An analogous behavior is exhibited also by [Co(bdtCl2)2]− (Figure S9, Supporting Information). As already discussed in the section Cyclic Voltammetry, quasireversible or irreversible voltammetric peaks were found upon oxidation for all complexes, indicating the slow rate of the electrode reactions or the chemical follow-up reactions of the oxidized complexes. Anodic oxidation of the nickel complex [Ni(bdtCl2)2]− and all [M(bdt)2]− samples in CH2Cl2 leads to an irreversible decrease of the initial optical bands, indicating

Figure 5. X-band EPR spectra of (a) [Ni(bdt)2]− and (b) [Ni(bdtCl2)2]− measured in DMF solution at room temperature (cyan lines) and at 100 K (black lines). E


Organometallics

Article

in the region of the first reduction peak the final solution was inserted under an argon atmosphere into the EPR tube, which was then immediately immersed into liquid nitrogen, and EPR spectra were recorded at 100 K. Figure 6b shows the EPR spectra of the one-electron-reduced copper complexes measured in this way at 100 K. Both copper complexes exhibit rich split EPR spectra at low temperatures, in contrast to the four-line isotropic EPR signal measured at room temperature. At 100 K, this EPR spectrum becomes anisotropic with a rhombic pattern similar to that observed for nickel complexes. Additionally, the splitting from the copper central atom makes the pattern of the observed spectra very complex. Similarly to the copper complexes, the one-electron-reduced cobalt complexes at the first reduction peak exhibit at 100 K a new EPR signal with a rhombic structure (Figure 6c) with hyperfine splitting due to 59Co nuclei (nuclear spin I = 7/2). Theoretical Calculations. Electronic Structure. Before proceeding to the theoretical interpretation of electrochemical behavior of the studied complexes, it might be worthwhile to explore the theoretical electronic structure, which is also directly related to the EPR behavior and/or the shape of the particular spin densities. A theoretical insight into the electronic structure is offered from the perspective of localized orbitals, Mulliken d populations, and QTAIM analysis. B3LYP/6-311g*/pcm calculations found that the anions [Co(bdt)2]−, [Ni(bdt)2]−, and [Cu(bdt)2]− in triplet, doublet, and singlet spin states, respectively, were energetically preferred. Hence, these anions will be denoted in the following text as 3[Co(bdt)2]−, 2[Ni(bdt)2]−, and 1[Cu(bdt)2]−. Following the aforementioned notation, their oxidized and reduced forms can be labeled, including the energetically preferred spin states, as 2[Co(bdt)2]0, 1[Ni(bdt)2]0, 2[Cu(bdt)2]0 and 2[Co(bdt)2]2−, 1[Ni(bdt)2]2−, 2[Cu(bdt)2]2−. The same spin states have been found as energetically preferred also for the complexes containing the bdtCl2 ligands. First of all, the Mulliken d populations as well as the excess of s populations above six electrons (neglecting 1s, 2s, and 3s “core” populations) on the central metal cations are related to localized orbitals. The number of directly localized d orbitals on the central metal cation corresponds to the formal oxidation state, while the (Mulliken) populations gained from the localized (M←S) donor bonds have to be in an agreement with the excess in the Mulliken d and s populations on the central metal atom (see Table 2). Following the deformed square-planar geometry and the Mulliken d populations of the coordination polyhedron [MS4] (note that each geometry has been aligned to fit the coordination system appropriately to the crystal field concept) the (nearly) doubly occupied dxy and dz2 orbitals can be considered unaffected by the ligand field. The bonding/nonbonding/antibonding interactions within the canonical Kohn−Sham orbitals, which also can cause shifts in the orbital order in regard to the crystal field theory, are not considered herein. On the other hand, the dx2−y2 orbitals are important as the acceptors of the electron density via the formation of σ coordination bonds. The only exception has been found by the localization procedure in the α domains for 2[Cu(bdt)2]2− with no σ dative bond and for 2[Cu(bdtCl2)2]2− with only a weak σ dative bond, due to the already filled dx2−y2 shells (see Table 2). The dxz and dyz populations vary between 1 (half-filled) and 2 (fully filled) with respect to the central atom and/or the charge of a complex. Obviously the s populations as well as the total d populations are relatively unaffected by the charge of a specific

The g matrix determined from the observed EPR spectra of paramagnetic nickel complexes is g = [2.188, 2.043, 2.011] for [Ni(bdt)2]− and g = [2.201, 2.043, 2.010] for [Ni(bdtCl2)2]−. These values are characteristic of the low-spin 3d7 electronic configuration. Examples of authentic Ni(III) include complexes such as [Ni(tacn)2]3+ (tacn = 1,4,7-triazacyclononane), for which g⊥ = 2.128 and g∥ = 2.026 as a powder at 150 K,21 and [Ni(Me2[14]aneN4)Cl2]+, for which g⊥ = 2.181 and g∥ = 2.025 as a powder at ∼300 K.22 As already mentioned above in In Situ EPR−UV/Vis/NIR Spectroelectrochemistry, upon the cathodic reduction of the EPR-silent copper complexes [Cu(bdt)2]− and [Cu(bdtCl2)2]− a new EPR signal arises at room temperature with g = 2.047 for [Cu(bdt)2]2− and g = 2.043 for [Cu(bdtCl2)2]2− and the characteristic splitting from 63,65Cu, confirming a strong contribution of the central atom with substantial delocalization of the unpaired spin onto the orbitals of the copper ion (Figure 6a).

Figure 6. X-band EPR spectra of (a) [Cu(bdt)2]2− and [Cu(bdtCl2)2]2− in 0.1 M TBAPF6/DMF solution at room temperature, (b) [Cu(bdt)2]2− and [Cu(bdtCl2)2]2− in frozen DMF solution at 100 K, and (c) [Co(bdt)2]2− in frozen DMF solution at 100 K.

To investigate the reduced states of copper complexes at low temperatures, we performed an ex situ spectroelectrochemical experiment, where the corresponding samples were electrolyzed in TBAPF6/DMF in a special electrolytic cell using a large platinum-mesh working electrode. After a complete reduction F


Organometallics

Article

Table 2. Mulliken Orbital Populations and the Localized Orbital (Pipek−Mezey) Analysis for [M(bdt)2]q/[M(bdtCl2)2]q Complexesa Mulliken population

localized orbital

system

dz2

dxz

dyz

dxy

dx2−y2

sσ

dtotal

dformald

σα(M/S)

σβ(M/S)

[Cu(bdt)2]0 1 [Cu(bdt)2]− 2 [Cu(bdt)2]2− 1 [Ni(bdt)2]0 1 [Ni(bdt)2]0 2 [Ni(bdt)2]− 1 [Ni(bdt)2]2− 2 [Co(bdt)2]0 3 [Co(bdt)2]− 2 [Co(bdt)2]2− 2 [Cu(bdtCl2)2]0 1 [Cu(bdtCl2)2]− 2 [Cu(bdtCl2)2]2− 1 [Ni(bdtCl2)2]0 1 [Ni(bdtCl2)2]0 2 [Ni(bdtCl2)2]− 1 [Ni(bdtCl2)2]2− 2 [Co(bdtCl2)2]0 3 [Co(bdtCl2)2]− 2 [Co(bdtCl2)2]2−

1.94 1.94 1.93 1.89

1.98 2.00 2.00 1.83

1.98 2.00 2.00 1.83

1.96 1.96 1.96 1.92

1.48 1.46 1.52 1.06

0.75 0.72 0.64 0.61

9.34 9.35 9.39 8.53

1.78 1.98 1.21 1.19 1.52 1.98 2.00 2.00 1.81

1.78 1.98 1.21 1.19 1.52 1.98 2.00 2.00 1.81

1.93 1.91 1.93 1.93 1.89 1.95 1.95 1.95 1.92

1.10 0.86 1.11 1.12 0.78 1.49 1.47 1.53 1.08

0.60 0.59 0.64 0.63 0.62 0.76 0.74 0.65 0.61

8.48 8.58 7.36 7.34 7.54 9.34 9.35 9.40 8.52

1.90 1.85 1.92 1.91 1.84

1.76 1.98 1.20 1.19 1.51

1.76 1.98 1.20 1.19 1.51

1.92 1.90 1.92 1.92 1.88

1.13 0.88 1.13 1.13 0.81

0.61 0.59 0.64 0.63 0.62

8.47 8.59 7.36 7.34 7.54

0.28/0.72 0.27/0.73 0.00/0.82 0.20/0.80 0.18/0.78b,c 0.21/0.79 0.16/0.86 0.24/0.75 0.24/0.76 0.17/0.85 0.28/0.71 0.27/0.72 0.06/0.95 0.21/0.80 0.19/0.78b,c 0.22/0.78 0.17/0.85 0.25/0.75 0.25/0.75 0.17/0.84

0.28/0.72

1.90 1.85 1.92 1.91 1.84 1.94 1.94 1.93 1.90

4α/4β 4α/4β 5α/4β 3.8α/3.8β 3α/3βb 4α/3β 4α/4β 4α/2β 4α/2β 4α/3β 4α/4β 4α/4β 5α/4β 3.8α/3.8β 3α/3βb 4α/3β 4α/4β 4α/2β 4α/2β 4α/3β

2

πβ(M/S)

0.20/0.81

0.20/0.81

0.11/0.81

0.19/0.82 0.19/0.83 0.14/0.88 0.28/0.71

0.10/0.86 0.09/0.87

0.21/0.81

0.20/0.81

0.11/0.81

0.19/0.82 0.19/0.82 0.14/0.87

0.09/0.85 0.08/0.86 0.02/0.90

a

Mulliken populations show the d and sσ orbital populations as the excess above the six s electrons in the 1s, 2s, and 3s core shells. Localized orbitals show the number of localized d orbitals on the central atom (dformal) as well as the polarization ratio of the metal/sulfur (M/S) localized bond orbitals. bFoster−Boys localization scheme. cEight instead of four σ and four π orbitals have been localized in each of the α and β domains. dnα/mβ denotes the number n/m of localized α/β d orbitals.

Figure 7. Localized DAFH eigenvectors of 1[Ni(bdt)2]0, including labels and occupations in the Mulliken domain of the Ni atom. The isosurface value is 0.02 bohr−1.5.

hand, the localized orbital approach is, in addition to others (such as domain averaged Fermi holes,23 natural orbitals for chemical valence,24 natural bond orbitals,25 projection based population analysis26), suited to distinguish between the d electron configuration of the central atom and the d and/or s electron density received from the donor−acceptor bonds (although for the current analysis all these contributions are given in the realms of Mulliken populations). Thus, in the 2 [Cu(bdt)2]0, 1[Cu(bdt)2]−, and 2[Cu(bdt)2]2− complexes the electron configurations of copper are formally d8, d8, and d9, which correspond to the formal oxidation states Cu(III), Cu(III), and Cu(II), respectively. A similar situation is found

metal complex. There is also only a small effect of the ligand (bdt vs bdtCl2) on the Mulliken orbital populations and/or localized orbital analysis. Thus, following the Mulliken total d populations, all central atoms are in the oxidation state M(II) with the d electron configurations of d9, d8, and d7 for copper, nickel, and cobalt, respectively. This can be related to the physical oxidation state of the central atoms. If one formally defines the d electron configuration by only those d shells which do not participate in the coordination bonding, the ordinary Mulliken population analysis does not lead to straightforward results, because all (bonding and nonbonding) populations are included (although the use of overlap populations seems a plausible option). On the other G


Organometallics

Article

for the cobalt complexes: i.e., the formal oxidation state Co(III) with the d6 electron configuration in 2[Co(bdt)2]0 and 3[Co(bdt)2]− complexes, unlike the oxidation state Co(II) with a d7 electron configuration in 2[Co(bdt)2]2−. A very interesting situation is found for nickel complexes. The localized orbitals indicate d7 and d8 electron configurations in 2[Ni(bdt)2]− and 1 [Ni(bdt)2]2−, respectively. On the other hand, for 1[Ni(bdt)2]0 the Pipek−Mezey (PM) localization procedure yields a formal d population on nickel of 7.6, which is equally distributed over α and β domains (3.8 each). It indicates only 80% localization of one of the occupied d orbitals on the nickel atom. On the other hand, the Foster−Boys (FB) localization procedure yielded a d6 electron configuration of Ni, leading to the formal oxidation state Ni(IV). The FB localized bonding orbitals correspond to a contribution of ordinary σ and strong π interactions involving dx2−y2, dxz/dyz, and 4s shells of nickel. To explore the differences between the PM and FB localization procedures in 1[Ni(bdt)2]0 complex in more detail, domain averaged Fermi holes23 (DAFH) analysis has been performed to further elucidate the electron structure of 1[Ni(bdt)2]0. The DAFH eigenvectors and the particular occupation numbers for the nickel domain of 1[Ni(bdt)2]0 are compiled in Figure 7. Apparently, DAFH analysis is in close agreement with results of the FB localized orbitals procedure: i.e., the DAFH formal delectron configuration is d6. Furthermore, DAFH analysis yields a complementary picture of the Ni−S bonding. In addition to the σ donor bonds due to the interaction of empty dx2−y2 and s orbitals of Ni with the px/py orbitals of sulfur, the dxy/dxz π interaction of Ni with the pz orbitals of S becomes apparent (see Figure 7). As found for the PM localized orbitals this π interaction accounts for 1.6 d electrons of Ni and can be assigned to the noninnocent character of the bdt/bdtCl2 ligand. Thus, taking into account the aforementioned considerations, from a strictly rigorous point of view Ni in 1[Ni(bdt)2]0 has a d6 formal electron configuration. However, the favorable interaction with the pz orbitals of sulfur atoms (see Figure 7d) indicates an antiparallel spin−spin interaction between the central atom and ligands in the diamagnetic 1[Ni(bdt)2]0 complex.16 The shapes of DAFH eigenvectors retained in the Ni domain are worth pointing out (see Figure 7). These indicate how much the noninnocent character of the bdt ligand affects the nonbonding d orbitals on the central atom. Similar shapes are found for the localized nonbonding d orbitals in all of the studied systems. To complete the discussion of the electronic structure of the studied complexes, the QTAIM charges of M and S as well as bond critical points (BCP) characteristics and/or M−S bond lengths are considered (see Tables 3 and 4). The Co/Ni charges in 2[Co(X)2]2−/1[Ni(X)2]2− (X = bdt, bdtCl2) differ considerably from those of their neutral and monocharged analogues, while the charge of copper is not really affected by the total charge of the complex. The charge of the central atom is also not really different for the bdt or bdtCl2 ligands, similarly as in the case of spin populations (see Table 3). QTAIM charges of the central atoms reflect the changes of the formal d configuration upon the total charge of the complex. On the other hand, the sulfur atomic charges are sensitive not only to the charge of the particular complex but also to the presence of Cl in the ligand. Thus, the polarization in bdtCl2 affects more the energetics of the complexes than the steric interactions of the dative bonds and/or the total d populations of the central atoms. Analogously, the complex charge considerably affects not only the M−S bond lengths and/or the metal−BCP

Table 3. QTAIM Charges, Spin Populations, and Spin Squares ⟨S2⟩ in the Complexes under Study atom charge system 2

0

[Cu(bdt)2] [Cu(bdt)2]− 2 [Cu(bdt)2]2− 1 [Ni(bdt)2]0 2 [Ni(bdt)2]− 1 [Ni(bdt)2]2− 2 [Co(bdt)2]0 3 [Co(bdt)2]− 2 [Co(bdt)2]2− 2 [Cu(bdtCl2)2]0 1 [Cu(bdtCl2)2]− 2 [Cu(bdtCl2)2]2− 1 [Ni(bdtCl2)2]0 2 [Ni(bdtCl2)2]− 1 [Ni(bdtCl2)2]2− 2 [Co(bdtCl2)2]0 3 [Co(bdtCl2)2]− 2 [Co(bdtCl2)2]2− 1

spin population sulfur

⟨S2⟩

metal

sulfur

metal

0.762 0.749 0.763 0.664 0.703 0.591 0.904 0.938 0.685 0.765 0.752 0.762 0.672 0.718 0.591 0.914 0.911 0.686

−0.160 −0.322 −0.500 −0.121 −0.300 −0.448 −0.191 −0.359 −0.471 −0.094 −0.246 −0.402 −0.052 −0.225 −0.353 −0.125 −0.282 −0.375

0.021

0.156

0.7537

0.464

0.128

0.7527

0.515

0.093

0.7569

1.760 1.843 1.045 0.025

−0.119 0.028 −0.009 0.163

1.3542 2.0498 0.7642 0.7540

0.456

0.127

0.7527

0.568

0.085

0.7586

1.782 1.873 1.046

−0.130 0.019 −0.009

1.3672 2.0524 0.7646

Table 4. QTAIM Characteristics of M−S Bonds in the Complexes under Study: M−S Bond Length RM−S, M−BCP Distance RM−BCP, BCP Electron Density ρBCP, Its Laplacian ΔρBCP, and Ellipticity ε

2

[Cu(bdt)2]0 1 [Cu(bdt)2]− 2 [Cu(bdt)2]2− 1 [Ni(bdt)2]0 2 [Ni(bdt)2]− 1 [Ni(bdt)2]2− 2 [Co(bdt)2]0 3 [Co(bdt)2]− 2 [Co(bdt)2]2− 2 [Cu(bdtCl2)2]0 1 [Cu(bdtCl2)2]− 2 [Cu(bdtCl2)2]2− 1 [Ni(bdtCl2)2]0 2 [Ni(bdtCl2)2]− 1 [Ni(bdtCl2)2]2− 2 [Co(bdtCl2)2]0 3 [Co(bdtCl2)2]− 2 [Co(bdtCl2)2]2−

RM−S, bohr

RM−BCP, bohr

ρBCP, e bohr−3

ΔρBCP, e bohr−5

ε

4.174 4.206 4.420 4.089 4.140 4.211 4.165 4.204 4.239 4.168 4.196 4.391 4.078 4.129 4.187 4.153 4.187 4.213

1.923 1.937 1.999 1.865 1.893 1.890 1.924 1.923 1.914 1.920 1.932 1.987 1.863 1.891 1.882 1.920 1.935 1.906

0.0876 0.0842 0.0654 0.0972 0.0922 0.0825 0.0944 0.0885 0.0819 0.0882 0.0852 0.0676 0.0986 0.0936 0.0849 0.0951 0.0914 0.0846

0.146 0.137 0.135 0.215 0.188 0.195 0.183 0.189 0.200 0.149 0.141 0.143 0.218 0.189 0.203 0.185 0.173 0.207

0.061 0.064 0.042 0.014 0.002 0.034 0.153 0.064 0.254 0.058 0.060 0.037 0.014 0.006 0.034 0.157 0.164 0.249

distances but also their BCP electron densities ρBCP and corresponding Laplacians ΔρBCP (see Table 4). The M−S bond length always increases with the charge of the complex, and a similar trend holds for the M−BCP distance in the case of copper complexes. Analogously to this, ρBCP exhibits a reverse trend. In the bdtCl2-containing complexes, the M−S and M−BCP lengths are shorter and the ρBCP characteristics are larger than in the bdt complexes, which might indicate stronger dative interactions in the case of bdtCl2 complexes. H


Organometallics

Article

Figure 8. Calculated spin density of all open-shell ground states for the investigated complexes in different redox states. The isosurface value is 0.005 e bohr−3.

Figure 9. Frontier molecular (spin) orbitals of [Cu(bdt)2]q complexes. The isosurface value is 0.05 bohr−1.5 (the particular orbital energies in hartree are given in parentheses).

Spectroscopy and Electrochemistry Interpretations. We have compared theoretical and experimental data for all investigated complexes in different charge states to find the most plausible explanation for the observed voltammetric and spectroscopic results described above. The results of theoretical calculations agree with the EPR behavior of the studied species. The EPR signal of 2[Ni(bdt)2]− can be assigned to a d7 configuration (seven localized d orbitals have been found on Ni; see Table 2). The triplet spin state of 3 [Co(bdt)2]− cannot be easily detected experimentally in Xband EPR. The inability to see an EPR signal in our experiment

might be due to the properties of the triplet spin state being unsuitable for observation by conventional EPR, as the zerofield splitting might be larger than the X-band microwave energy and the temperature limit is 100 K in our laboratory. The singlet spin state complex 1[Co(bdt)2]− has not been favored energetically at the B3LYP level of theory as mentioned above, although it has to be stressed that the preference of a low-spin state over a high-spin state can be dependent on the portion of exact exchange in the employed DFT functional. In accord with the theoretical prediction, the reduced form of the cobalt complex 2[Co(bdt)2]2− exhibited a typical Co(II) EPR I


Organometallics

Article

Figure 10. Frontier molecular (spin) orbitals of [Ni(bdt)2]q− complexes. The isosurface value is 0.05 bohr−1.5 (the particular orbital energies in hartree are given in parentheses).

the parent monoanions clearly indicate that the redox site for the reduction is mainly located on the moiety built from the central atom and the sulfur atoms of the [MS4] coordination polyhedron. It is noteworthy, with respect to the redox process, that the α-HOMO and β-LUMO shapes of 2[Cu(bdt)2]2− are the same as the LUMO shape of the 1[Cu(bdt)2]− anion (see Figure 9). Hence, the electron enters the LUMO during the reduction of 1[Cu(bdt)2]−, which becomes the α-HOMO and β-LUMO (i.e., SOMO) of 2[Cu(bdt)2]2−. Similarly to the oxidation of 1 [Cu(bdt)2]−, the β-LUMO of 2[Cu(bdt)2]0 (representing the SOMO) resembles exactly the HOMO of 1[Cu(bdt)2]− (Figure 9). The same is found also for the oxidation and reduction of 2 [Ni(bdt)2]− (Figure 10). This agreement between the LUMO of the oxidized forms and the HOMO of the reduced forms is quite remarkable. The shape and (more precisely) the order of the unrestricted α and β canonical orbitals are often different, which makes the identification of the SOMO less straightforward (essentially the β-LUMO) for transition-metal complexes. In this regard, the β-LUMO of 2[Co(bdt)2]0 and α-HOMO of 3 [Co(bdt)2]− are considerably different (Figure 11). Nevertheless, in the case of 2[Co(bdt)2]0 the α-LUMO (see Figure 11) has a considerably lower energy (by about 0.0280 hartree = 0.762 eV = 73.6 kJ/mol) than the β-LUMO and has moreover the same shape as the α-HOMO of 3[Co(bdt)2]−. This is in agreement with the fact that the reduction of 2 [Co(bdt)2]0 leads to 3[Co(bdt)2]− in a triplet spin state: i.e., the electron enters the α-LUMO. The shapes of the β-LUMO of 3[Co(bdt)2]− and α-HOMO of 2[Co(bdt)2]2− differ in the charge redistribution between the central atom and the ligand after reduction. Note that the highest occupied orbital of 2 [Co(bdt)2]2− is actually the β-HOMO with an energy of −0.1312 hartree; thus, the reduction of 2[Co(bdt)2]2− leads to a triplet state of [Co(bdt)2]−. The shapes of the frontier orbitals of the bdt ligand containing systems also remain the same for analogous bdtCl2 complexes, and the same holds for the following interpretation. All corresponding frontier orbitals for [M(bdtCl2)2]q species (q = 0, −1, −2) are collected in Figure S12 (Supporting Information). Furthermore, it is fair to stress that any issues connected with the choice of a DFT

signal at 100 K, as proved by EPR spectroelectrochemistry (see Figure 6c). Interestingly, the spin density distribution in the isoelectronic doublet spin structures 2[Co(bdt)2]2−, 2[Ni(bdt)2]−, and 2[Cu(bdt)2]0 is moved from the central atom to the ligands in the sequence 2[Co(bdt)2]2− < 2[Ni(bdt)2]− < 2[Cu(bdt)2]0, as indicated in Figure 8 and Table 3. In addition to QTAIM spin populations, the actual QTAIM charges of central metal atoms and sulfur atoms, as well as the spin contamination (measurable as the difference between the values of calculated and ideal spin squares ⟨S2⟩) for all studied species, are shown in Table 3. Spin densities of bdtCl2-containing species are collected in Figure S11 (Supporting Information). Generally, most of the unrestricted DFT spin densities are very well resolved by the particular β-LUMO (the lowest unoccupied molecular spin− orbital, actually representing the SOMO, the single occupied molecular orbital), especially for Cu and Ni species (see Figures 8−10). In the case of 3[Co(bdt)2]− both β-LUMO and β-LUMO+1, formally representing the open-shell orbitals SOMO and SOMO+1, have to be summed up to obtain the shape of the particular spin density (compare Figures 11 and 8). A similar situation is found also for 2[Co(bdt)2]0, although the 2[Co(bdt)2]0 species is affected by a considerable spin contamination. The theoretical value of ⟨S2⟩ for one unpaired electron is 0.75, while its actual value for 2[Co(bdt)2]0 is 1.354. The large spin contamination is further confirmed by the Co spin density of 1.782 and has to be counterbalanced by the negative spin density populations on sulfur atoms (altogether −0.520) as well as on the remaining atoms (altogether −0.262; see Table 3). The actual spin density distribution of 2[Co(bdt)2]0 seems very similar to the spin density of the 3[Co(bdt)2]− system. Additionally, the β-LUMO+1 of 2[Co(bdt)2]0 resembles exactly the shape of β-LUMO+1 of the 3[Co(bdt)2]− anion (see Figure 11). The shape of LUMO orbitals is not only related to spin densities and/or the SOMO in the case of open-shell systems but is in many cases a qualitative indicator of the reduction site for a given species. Figures 9−11 show the LUMOs (as well as HOMOs) for all investigated complexes in their initial, monoreduced, and mono-oxidized forms. The presented LUMOs of J


Organometallics

Article

Figure 11. Frontier molecular (spin) orbitals of [Co(bdt)2]q complexes. The isosurface value is 0.05 bohr−1.5 (the particular orbital energies in hartree are given in parentheses). In the case of UB3LYP calculations of a triplet state β-LUMO+1 should be regarded as the SOMO+1 (or the ”upper” SOMO).

Table 5. Comparison of Theoretical Redox Potentials and Δgaptheor vs Δgapexpa system

−EA

IP

gaptheor

Δgaptheor

Δgaptheorb

gapexptl

Δgapexptl

Δgapexptlb

[Cu(bdt)2]− [Ni(bdt)2]− 3 [Co(bdt)2]− 1 [Cu(bdtCl2)2]− 2 [Ni(bdtCl2)2]− 3 [Co(bdtCl2)2]−

−1.62 −1.88 −2.06 −1.46 −1.59 −2.06

−0.42 −0.52 −0.48 0.04 0.15 0.06

1.20 1.36 1.58 1.50 1.74 2.12

−0.30 −0.38 −0.54

0.00 −0.08 −0.24

1.33 0.92 1.30 1.37 0.87 0.99

0.04 −0.05 −0.31

0.00 −0.09 −0.35

1 2

a

See text. All values are in volts (electronvolts). bΔgap aligned to the highest shift: i.e., 1[Cu(bdt)2]− vs 1[Cu(bdtCl2)2]−.

appropriate total energies). For instance, the lowest reduction potential observed for the 3[Co(bdt)2]− anion can be explained from a theoretical point of view by the lowest −EA value for 3 [Co(bdt)2]− in comparison to those for 2[Ni(bdt)2]− and 1 [Cu(bdt)2]− (Table 5). A similar conclusion is found for the Co complex of bdtCl2. The second lowest −EA value was

functional, and its impact on the order of the frontier orbitals, have been left out of any consideration. Herein a theoretical inspection of the reduction and oxidation potentials will employ the negative electron affinities (−EA) and the values of ionization potentials (IP), respectively, relative to the Fc+/Fc couple (using the difference of K


Organometallics

Article

preferred also for the complexes containing the bdtCl2 ligands. Furthermore, for the monoanionic complexes one can see that the central atom is reduced, as indicated by a change in the formal d electron configuration. On the other hand, the formal d electron configuration of the central metal remains rather unaffected by the oxidation, which means that prevailingly the ligands become the redox-active sites. The monoanionic nickel samples exhibited a broad singlet EPR signal at room temperature which became anisotropic at 100 K with a characteristic rhombic pattern. The results of B3LYP calculations agree with the EPR behavior of the studied species. The EPR signal of 2[Ni(bdt)2]− and 2[Ni(bdtCl2)2]− can be assigned to a d7 electron configuration of the Ni central atom. Cathodic reduction of copper and cobalt complexes leads to paramagnetic species having an EPR signal with splitting from 63,65Cu for copper and from 59Co for cobalt samples, confirming a strong contribution of the central atom with a substantial delocalization of the unpaired spin onto the central atom. Both copper complexes exhibit unusually rich split anisotropic EPR spectra with different g values at low temperatures, in contrast to the four-line isotropic EPR signal measured at room temperature. At the first nearly reversible reduction peak of [Ni(bdt)2]− and [Ni(bdtCl2)2]− an optical absorption of the dominating band in the NIR region decreases and a new optical transition in the visible region appears via isosbestic points. The same holds for the cobalt complexes, where upon reduction a reversible behavior in the corresponding optical spectra was confirmed using in situ EPR−UV/vis/ NIR spectroelectrochemistry, and the reduction was accompanied by remarkable reversible color changes of the corresponding solutions.

found for copper and the least negative value for the nickel monoanionic complex (see Table 5). This agrees well with the order of the experimental reduction potentials of the monoanionic complexes shown in Table 1 (see also Figure 1). In the case of anodic oxidations, the ionization potentials explain well the trend in the oxidation potentials found for nickel and cobalt samples (Table 5), while for the copper sample the estimated value does not correspond to the observed oxidation potential in the anodic region. Nevertheless, it is relatively hard to assess the exact reduction/ oxidation potentials due to the presence of different effects (e.g., explicit solvent effects, interactions within the electrode interface) which are not accounted for in the theoretical treatment. On the other hand, the theoretical “electrochemical” gap for the difference between the IP and −EA values and experimental electrochemical gaps fit each other qualitatively well, although the theoretical gap for copper complexes is still underestimated. Even better agreement is obtained for the theoretical and experimental ligand-based differences of the gaps (Δgap), defined as Δgaptheor = gap([M(bdtCl 2)2 ])theor − gap([M(bdt)2 ])theor (1)

Δgapexptl = gap([M(bdtCl 2)2 ])exptl − gap([M(bdt)2 ])exptl (2)

After the gap differences (Δgap) are aligned to the highest value of Cu, the agreement between theory and experiment is semiquantitative (see Table 5). Although it is not discussed in any detail, by an inspection of the differences between the LUMO(ox) and HOMO(red) orbital energies the same trends are found as for the ionization potentials.

■

■

CONCLUSION The redox properties of copper, nickel, and cobalt complexes (MePh3P)[M(bdt)2] with the benzene-1,2-dithiolate ligand (bdt) and (MePh3P)[M(bdtCl2)2] with the 3,6-dichlorobenzene-1,2-dithiolate ligand (bdtCl2) were studied by cyclic voltammetry and in situ EPR−UV/vis/NIR spectroelectrochemistry. Half-wave potentials for the reduction of the corresponding anions in solution increase in the sequence [Co(bdt)2]− < [Cu(bdt)2]− < Ni(bdt)2]−. The lowest reduction potential observed for the [Co(bdt)2]− anion can be explained by the lowest −EA value for 3[Co(bdt)2]− in comparison to those for [Ni(bdt)2]− and [Cu(bdt)2]−. A similar conclusion is found for the complexes of bdtCl2. The electron-accepting chlorine substituent facilitates the reduction, and so the electrode potential is shifted to the anodic region, retaining the sequence of redox potentials for different central atoms. The chemical stability of one-electron-reduction products [M(bdt)2]2− in dichloromethane increases in the sequence [Ni(bdt)2]2− < [Cu(bdt)2]2− < [Co(bdt)2]2−. On the other hand, for newly prepared [M(bdtCl2)2]− complexes all cyclic voltammograms are chemically reversible and all reduction products are kinetically stable. B3LYP/6-311g*/ pcm calculations of the monoanions in the triplet spin state for cobalt, doublet spin state for nickel, and singlet spin state for copper complexes as well as for their one-electron-oxidized and -reduced products were performed. According to our calculations, the energetically preferred oxidized and reduced forms (labeled by spin states) are 2[Co(bdt)2]0, 1[Ni(bdt)2]0, 2 [Cu(bdt)2]0 and 2[Co(bdt)2]2−, 1[Ni(bdt)2]2−, 2[Cu(bdt)2]2−. The same spin states have been found to be energetically

EXPERIMENTAL SECTION

Materials and Physical Measurements. For syntheses, the following reagents were used: Na (metallic) from Ferak Berlin; benzene-1,2-dithiol (≥95%), 3,6-dichlorobenzene-1,2-dithiol (≥95%), NiCl2·6H2O (≥98%), CoCl2·6H2O (≥98%), CuCl2·2H2O (≥99%), and methyltriphenylphosphonium bromide (≥98%) from SigmaAldrich. All solvents were products of mikroCHEM of centralchem p.a. grade. For cyclic voltammetry experiments commercially available dichloromethane (CH2Cl2), N,N-dimethylformamide (DMF), and ferrocene (Fc) purchased from Sigma-Aldrich were used without further purification. Tetrabutylammonium hexafluorophosphate (TBAPF6) of purissimum quality (Fluka) was dried under reduced pressure at 70 °C for 24 h and stored in a glovebox. Cyclic voltammograms were recorded in CH2Cl2 or in DMF with 0.1 M TBAPF6 as the supporting electrolyte using a one-compartment electrochemical cell with platinum wires as the working and counter electrodes and a silver wire as the pseudoreference electrode. Samples under study were dissolved at a concentration of 0.2 mM in dichloromethane. All electrochemical measurements were performed under an inert nitrogen atmosphere. Cyclic voltammograms were measured using an Autolab electrochemical analyzer equipped with a PGSTAT 100 potentiostat or a HEKA PG 390 potentiostat. The redox potentials are stated against the ferrocenium/ferrocene couple (Fc+/ Fc). In situ EPR−UV/vis/NIR spectroelectrochemical experiments were performed in the optical EPR cavity (ER 4104OR, Bruker, Germany). EPR spectra were recorded with an EMX X-band CW spectrometer (Bruker, Germany). UV/vis/NIR spectra were measured using an Avantes AvaSpec-2048x14-USB2 spectrometer with a CCD detector and AvaSpec-NIR256-2.2 instrument with an InGaAs detector and applying the AvaSoft 7.5 software. Both the EPR spectrometer and the UV/vis/NIR spectrometer are linked to a HEKA PG 390 potentiostat, which triggers both spectrometers. Triggering is performed by the L


Organometallics

Article

visualized via a modified Gaussian03 fchk file using the IQmol package.41

software package PotMaster v2x40 (HEKA Electronik). EPR spectra were simulated using the standard Bruker program WinEPR SimFonia (Bruker). For standard in situ EPR−UV/vis/NIR spectroelectrochemical experiments an EPR flat cell was used. Laminated platinum mesh as the working electrode, silver wire as the pseudoreference electrode, and a platinum wire as the counter electrode were used in all spectroelectrochemical experiments in dichloromethane. TBAPF6 supporting electrolyte (0.2 M) was used to improve the electrochemical response in a flat spectroelectrochemical cell, with a larger iR drop in comparison to the conventional electrochemical cell. Synthesis of (MePh3P)[M(bdt)2] and (MePh3P)[M(bdtCl2)2] (M = Co, Ni, Cu). For syntheses of complexes with the bdt and bdtCl2 ligands, the modified procedure of Mrkvová et al.16 was used. In the preparation of (Me(Ph)3P)[Cu(bdt)2], a solution of Na (0.08 g, 3.3 mmol) in MeOH (10 cm3) was added to benzene-1,2-dithiol (bdtH2; 0.23 g, 1.6 mmol). To this mixture CuCl2·2H2O (0.13 g, 0.76 mmol) in MeOH (10 cm3) was added. Finally, methyltriphenylphosphonium bromide (Me(Ph)3PBr; 0.57 g, 1.6 mmol) in MeOH (10 cm3) was added. The resulting solution was stirred for 1 h. The complex was precipitated by the slow addition of water, with vigorous stirring. The powder was filtered off and washed with diethyl ether. The product was extracted with acetone from the crude product (yield 90%). Anal. Calcd (found): C, 59.93 (59.70); H, 4.22 (4.35); S, 20.64 (20.37). The same procedure was used for the preparation of the complex (Me(Ph)3P)[Ni(bdt)2]: NiCl2·6H2O (0.18 g, 0.76 mmol) (yield 89%). Anal. Calcd (found): C, 60.40 (60.33); H, 4.25 (4.20); S, 20.81 (20.67), The same procedure was also used for (Me(Ph)3P)[Co(bdt)2]: CoCl2·6H2O (0.18g, 0.76 mmol) (yield 85%). Anal. Calcd (found): C, 60.37 (60.30); H, 4.25 (4.21); S, 20.79 (20.47). Complete characterization and X-ray structures (CCDC refcode: AWIHET, AWIHAP, AWIHIX) of bdt complexes were reported by Mrkvová et al.16 An analogous procedure was used for the synthesis of the new complexes (Me(Ph)3P)[M(bdtCl2)2] (M = Cu, Ni, Co) with shorter stirring time of the resulting solution (5 min) and a larger weight of ligand (0.34 g, 1.6 mmol). The reaction yield was 55% for (Me(Ph)3P)[Cu(bdtCl2)2] (CuCl2·6H2O), 45% for (Me(Ph)3P)[Ni(bdtCl2)2] (NiCl2·6H2O), and 75% for (Me(Ph)3P)[Co(bdtCl2)2] (CoCl2·6H2O). The structures of the prepared complexes (Me(Ph)3P)[M(bdtCl2)2] (M = Cu, Ni, Co) have been confirmed by Xray single-crystal analysis.27 Computational Details. All calculations have been performed using Gaussian0328 and Gaussian0929 software. The geometry optimization of the systems under study has been performed with the B3LYP hybrid functional30 with 6-311g* basis sets for H, C, S ,and Cl atoms31 and Wachters basis sets with f polarization functions32 for Cu, Co, and Ni atoms. The solvent effect of CH2Cl2 solution has been approximated within the polarized continuum model (pcm).33 The stability of the optimized structures has been tested by vibrational analysis (no imaginary vibrations). The Molekel package34 has been used for the visualization of molecular orbitals and spin densities. Electronic structure parameters have been evaluated in the terms of QTAIM35 (quantum theory of atom in molecule) analysis using the AIMAll package.36 Alternatively, localized orbital analysis was performed via the Pipek−Mezey (PM)37 and Foster−Boys (FB)38 methods in the ORCA package.39 Canonical orbitals with energies below −2.0 hartrees were excluded from the localization procedure. The ORCA package was used also to evaluate the Mulliken orbital populations and for localized orbital population analysis. In addition, the domain averaged Fermi holes (DAFH) approach23 has been employed in the cases of a qualitative disagreement between the results of PM and FB methods. The DAFH analysis based on the Mulliken-like definition of atoms has been performed by the WinFermi program40 using the wave function produced by the Gaussian03 software. Within the central atom domain the electron pairs retained in the domain (localized atomic orbitals, lone pairs with occupation numbers close to 2) and the broken valences (representing the coordination bonds with occupation numbers between ca. 0.2 and 1.8) are analyzed. The obtained DAFH eigenvectors have been

■

ASSOCIATED CONTENT

■

AUTHOR INFORMATION

S Supporting Information *

Cyclic voltammograms for the reduction of [Cu(bdt)2]− and [Ni(bdt)2]− recorded at different scan rates in TBAPF6/ CH2Cl2 (Figure S1), cyclic voltammograms for the reduction of [Cu(bdtCl2)2]− and [Ni(bdtCl2)2]− recorded at different scan rates (Figure S2), cyclic voltammograms for the reduction of [Ni(bdt)2]− in TBAPF6/DMF (Figure S3), cyclic voltammograms of [Co(bdtCl2)2]− for the oxidation in TBAPF6/CH2Cl2 recorded at scan rates from 5 to 100 mV s−1 (Figure S4), experimental and simulated EPR spectra of [Cu(bdt)2]2− at room temperature (Figure S5), in situ EPR−UV/vis/NIR spectroelectrochemistry of [Cu(bdtCl2)2]− (Figure S6), EPR spectra of [Ni(bdt)2]− in 0.1 M TBAPF6/DMF before and after the cathodic reduction at the first reduction peak (Figure S7), in situ EPR−UV/vis/NIR spectroelectrochemistry of [Ni(bdtCl2)2]− (Figures S8 and S10), in situ UV/vis/NIR spectroelectrochemistry of [Co(bdtCl2)2]− (Figures S9 and S11), HOMO and LUMO orbitals of all [M(bdtCl2)2]q complexes (Figure S12), spin density of all open shell ground states calculated for the investigated complexes in different redox states (Figure S11), and xyz file giving the computed Cartesian coordinates of all of the molecules reported in this study. This material is available free of charge via the Internet at http://pubs.acs.org. Corresponding Author

*E-mail for P.R.: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This paper is dedicated to the memory of Prof. Dr. Lothar Dunsch, Dr. h.c. (1948−2013). The financial support of the Science and Technology Assistance Agency (contract No. APVV-0202-10) and Slovak Grant Agency VEGA (contracts 1/ 0679/11, 1/0289/12, 1/0327/12, and 1/0307/14), project SKAT-0027-12, and IFW Dresden is gratefully acknowledged. We thank the HPC center at the Slovak University of Technology in Bratislava, which is a part of the Slovak Infrastructure of High Performance Computing (SIVVP project No. 26230120002, funded by the European Region Development Funds) and the Vienna Supercomputing Center for computing facilities.

■

REFERENCES

(1) Zuo, J. L.; You, F.; You, X. Z.; Fun, H. K. Polyhedron 1997, 8, 1465−1469. (2) Cassoux, P.; Valade, L.; Kobayashi, H.; Kobayashi, A.; Clark, R. A.; Underhill, A. E. Coord. Chem. Rev. 1991, 110, 115−119. (3) Szolnai, T. Die chemoterapeutischen und pesticiden Wirkungen der Thiolreagenzien; Akad. Kiado: Budapest, Hungary, 1975; p 302. (4) Geiger, W. E.; Minest, E. Inorg. Chem. 1975, 14, 2141−2147. (5) Bigoli, F.; Deplano, P.; Devillanova, F. A.; Ferraro, J. R.; Lippolis, V.; Lukes, P. J.; Mercuri, M. L.; Pellinghelli, M. A.; Trogu, E. F.; Williams, J. M. Inorg. Chem. 1997, 36, 1218−1226. (6) Drexhage, K. H.; Mueller-Westerhoff, U. T. IEEE Quant. Electron. 1972, 8, 759−759. (7) Coomber, A. T.; Beljonne, D.; Friend, R. H.; Bredas, J. L.; Charlton, A.; Robertson, N.; Underhill, A. E.; Kurmoo, M.; Day, P. Nature 1996, 380, 144−146.

M


Organometallics

Article

(8) Cammack, R.; Fernandez, V. M.; Schneider, K. Nickel in hydrogenases from sulfate-reducing, photosynthetic, and hydrogenoxidizing bacteria. In The Bioinorganic Chemistry of Nickel; Lancaster, J. R., Ed.; VCH: New York, 1988; pp 167−190. (9) Sellmann, D.; Haussinger, D.; Knoch, F.; Moll, M. J. Am. Chem. Soc. 1996, 118, 5368−5374. (10) Jørgensen, C. K. Coord. Chem. Rev. 1966, 1, 164−178. (11) Chaudhuri, P.; Nazari Verani, C.; Bill, E.; Bothe, E.; Weyhermüller, T.; Wieghardt, K. J. Am. Chem. Soc. 2001, 123, 2213−2223. (12) Shimazaki, Y.; Yajima, T.; Tani, F.; Karasawa, S.; Fukui, K.; Naruta, Y.; Yamauchi, O. J. Am. Chem. Soc. 2007, 129, 2559−2568. (13) Cerdeira, A. C.; Afonso, M. L.; Santos, I. C.; Pereira, L. C. J.; Coutinho, J. T.; Rabaça, S.; Simão, D.; Henriques, R. T.; Almeida, M. Polyhedron 2012, 44, 228−237. (14) Zanello, P.; Grigotti, E. Homoleptic, Mononuclear Transition Metal Complexes of 1,2-Dithiolenes: Updating their Electrochemicalto-Structural Properties. In Trends in Molecular Electrochemistry; Pombeiro, A. J. L., Amatore, C., Eds.; Fontis Media-Marcel Dekker: New York, 2004; pp 3−70. (15) Chaudhuri, P.; Verani, C. N.; Bill, E.; Bothe, E.; Weyhermű ller, T.; Wieghardt, K. J. Am. Chem. Soc. 2001, 123, 2213−2223. (16) (a) Bachler, V.; Olbrich, G.; Neese, F.; Wieghardt, K. Inorg. Chem. 2002, 41, 4179−4193. (b) Ray, K.; Weyhermüller, T.; Neese, F.; Wieghardt, K. Inorg. Chem. 2005, 44, 5345−5360. (c) Ray, K.; Begum, A.; Weyhermüller, T.; Piligkos, S.; van Slageren, J.; Neese, F.; Wieghardt, K. J. Am. Chem. Soc. 2005, 127, 4403−4415. (17) Mrkvová, K.; Kameníček, J.; Šindelár,̌ Z.; Kvítek, L.; Mrozinski, J.; Nahorska, M.; Ž ák, Z. Transition Met. Chem. 2004, 29, 238−244. (18) Sellmann, D.; Binder, H.; Häußinger, D.; Heinemann, F. W.; Sutter, J. Inorg. Chem. Acta 2000, 300, 829−836. (19) Šoralová, S.; Breza, M.; Gróf, M. Polyhedron 2011, 30, 307−314. (20) Baker-Hawkes, M. J.; Billig, E.; Gray, H. B. J. Am. Chem. Soc. 1966, 88, 4870−4875. (21) Wieghardt, K.; Walz, W.; Nuber, B.; Weiss, J.; Ozarowski, A.; Statemeier, H.; Reinen, D. Inorg. Chem. 1986, 25, 1650−1654. (22) Gore, E. S.; Busch, D. H. Inorg. Chem. 1973, 12, 1−3. (23) (a) Ponec, R. J. Math. Chem. 1997, 21, 323−333. (b) Ponec, R.; Roithova, J.; Girones, X.; Frenking, G. Organometallics 2004, 23, 1790−1796. (c) Ponec, R.; Feixas, F. J. Phys. Chem. A 2009, 113, 5773−5779. (24) (a) Mitoraj, M.; Michalak, A. J. Mol. Model. 2007, 13, 347−355. (b) Michalak, A.; Mitoraj, M.; Ziegler, T. J. Phys. Chem. A 2008, 112, 1933−1939. (c) Mitoraj, M. P.; Michalak, A.; Ziegler, T. J. Chem. Theory Comp. 2009, 5, 962−975. (25) (a) Foster, J. P.; Weinhold, F. J. Am. Chem. Soc. 1980, 102, 7211−7218. (b) Reed, A. E.; Weinhold, F. J.; Curtiss, L. A.; Pochatko, D. J. Chem. Phys. 1986, 84, 5687−570. (c) Reed, A. E.; Curtiss, L. A.; Weinhold, F. Chem. Rev. 1998, 88, 899−926. (26) Bast, R.; Koers, A.; Gomes, A. S. P.; Iliaš, M.; Visscher, L.; Schwerdtfeger, P.; Saue, T. Phys. Chem. Chem. Phys. 2011, 13, 864− 876. (27) (a) Herich, P.; Kožíšek, J. Acta Crystallogr., Sect. C (in preparation). (b) Herich, P.; Fronc, M.; Lemée-Cailleau, M.-H.; Mason, S. A.; Kožíšek, J. Acta Crystallogr., Sect. B (in preparation). (28) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb, M. A.; Cheeseman, J. R.; Montgomery, J. A., Jr.; Vreven, T.; Kudin, K. N.; Burant, J. C.; Millam, J. M.; Iyengar, S. S.; Tomasi, J.; Barone, V.; Mennucci, B.; Cossi, M.; Scalmani, G.; Rega, N.; Petersson, G. A.; Nakatsuji, H.; Hada, M.; Ehara, M.; Toyota, K.; Fukuda, R.; Hasegawa, J.; Ishida, M.; Nakajima, T.; Honda, Y.; Kitao, O.; Nakai, H.; Klene, M.; Li, X.; Knox, J. E.; Hratchian, H. P.; 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.; Ayala, P. Y.; Morokuma, K.; Voth, G. A.; Salvador, P.; Dannenberg, J. J.; Zakrzewski, V. G.; Dapprich, S.; Daniels, A. D.; Strain, M. C.; Farkas, O.; Malick, D. K.; Rabuck, A. D.; Raghavachari, K.; Foresman, J. B.; Ortiz, J. V.; Cui, Q.; Baboul, A. G.; Clifford, S.; Cioslowski, J.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.; Komaromi, I.;

Martin, R. L.; Fox, D. J.; Keith, T.; Al-Laham, M. A.; Peng, C. Y.; Nanayakkara, A.; Challacombe, M.; Gill, P. M. W.; Johnson, B.; Chen, W.; Wong, M. W.; Gonzalez, C.; Pople, J. A. Gaussian 03, Revision C.02; Gaussian, Inc., Wallingford, CT, 2004. (29) 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.; Kobayashi, R.; Normand, J.; Raghavachari, K.; Rendell, A.; Burant, J. C.; Iyengar, S. S.; Tomasi, J.; Cossi, M.; Rega, N.; Millam, N. J.; 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 D.01; Gaussian, Inc., Wallingford, CT, 2009. (30) (a) Becke, A. D. J. Chem. Phys. 1993, 98, 5648−5652. (b) Lee, C.; Yang, W.; Parr, R. G. Phys. Rev. B 1988, 37, 785−789. (c) Vosko, S. H.; Wilk, L.; Nusair, M. Can. J. Phys. 1980, 58, 1200−1211. (d) Stephens, P. J.; Devlin, F. J.; Chabalowski, C. F.; Frisch, M. J. J. Phys. Chem. 1994, 98, 11623−11627. (31) (a) Krishnan, R.; Binkley, J. S.; Seeger, R.; Pople, J. A. J. Chem. Phys. 1980, 72, 650−654. (b) McLean, A. D.; Chandler, G. S. J. Chem. Phys. 1980, 72, 5639−5648. (32) (a) Wachters, A. J. H. J. Chem. Phys. 1970, 52, 1033−1036. (b) Bauschlichter, C. W., Jr.; Langhoff, S. R.; Barnes, L. A. J. Chem. Phys. 1989, 91, 2399−2411. (33) Miertuš, S.; Scrocco, E.; Tomasi, J. Chem. Phys. 1981, 55, 117− 129. (34) Varetto, U. MOLEKEL 5.4.0.8; Swiss National Supercomputing Centre, Lugano, Switzerland; http://molkel.cscs.ch/wiki/pmwiki.php. (35) Bader, R. F. W. Atoms in Molecules: A Quantum Theory; Clarendon Press: Oxford, U.K., 1990. (36) Keith, T. A. AIMAll, v. 11.12.19, TK Gristmill Software, Overland Park, KS, 2011; http://aim.tkgristmill.com. (37) Pipek, J.; Mezey, P. G. J. Chem. Phys. 1989, 90, 4916−4926. (38) Foster, S.; Boys, S. F. Rev. Mod. Phys. 1960, 32, 303−304. (39) Neese, F. ORCA program system. Wiley Interdiscip. Rev.: Comput. Mol. Sci. 2012, 2, 73−78. (40) Feixas, F.; Girones, X.; Ponec, R.; Roithova, J. WinFermi, v2.5; ICPF, AS CR, Prague, Czech Republic, 2008 (available upon request, contact [email protected]). (41) Gilbert, A. IQmol 2.3.0, 2014; http://iqmol.org/.

N


Dithiolate Ligands - American Chemical Society

Recommend Documents