- Email: [email protected]
Electron transfer studies of dithiolate complexes: effects of ligand variation and metal substitution
Inorganica Chimica Acta 341 (2002) 85 /90 www.elsevier.com/locate/ica
Electron transfer studies of dithiolate complexes: effects of ligand variation and metal substitution Anne E. McElhaney, Frank E. Inscore, Julien T. Schirlin, John H. Enemark * Department of Chemistry, The University of Arizona, Tucson, AZ 85721-0041, USA Received 28 April 2002; accepted 31 July 2002 Dedicated to Professor Kenneth N. Raymond on the occasion of his 60th birthday
Abstract Solution redox potentials and heterogeneous electron transfer rate constants have been measured for mono-ene-1,2-dithiolate complexes of the type (Tp*)M(E)(S S S) [Tp* /hydrotris(3,5-dimethyl-1-pyrazolyl)borate; M /Mo, W; E /O, NO]. The dithiolate ligands (S S S) equatorially coordinated to the central metal include: 1,2-benzenedithiolate (bdt); 3,6-dichloro-1,2-benzenedithiolate (bdtCl2); and 3,4-toluenedithiolate (tdt). Cyclic voltammograms reveal quasi-reversible one-electron reductions for all of the compounds; and, a one-electron quasi-reversible oxidation process is also observed for (Tp*)MoO(bdt), (Tp*)MoO(tdt), and (Tp*)WO(tdt). Electrochemical potentials are strongly dependent upon the nature of the substituents on the benzene ring of the dithiolate ligand. The nitrosyl /molybdenum complexes (E /NO) are more difficult to reduce than their corresponding oxomolybdenum complexes. The tungsten compound, (Tp*)WO(tdt), has the most negative electrochemical potentials among the complexes investigated. Heterogeneous electron transfer rates are insensitive to the variations of the central metal and ligands. These results support a highly covalent bonding interaction between the metal and the dithiolate ligands that modulates electron transfer reactions within these compounds. # 2002 Elsevier Science B.V. All rights reserved. Keywords: Electrochemistry; Molybdenum; Dithiolate; Heterogeneous electron transfer kinetics
1. Introduction Molybdenum-containing enzymes play essential metabolic roles in animals, plants, and microorganisms by catalyzing two-electron oxidations or reductions of various substrate molecules [1,2]. The molybdenum (Mo) center experiences oxidation states of VI, V, and IV during catalysis, and the initial oxidation state of the Mo active site is regenerated via two one-electron transfer reactions. Crystal structures obtained for several of these enzymes reveal common active site structural features, namely one (or two) pyranopterinene-1,2-dithiolate ligand and at least one terminal oxo ligand [3 /13]. The ene-dithiolate (Fig. 1) is believed to modulate electron transfer reactions within these en-
* Corresponding author. Tel.: /1-520-621 2245; fax: /1-520-626 8065 E-mail address: [email protected] (J.H. Enemark).
zymes, ensuring an efficient catalytic process [1,14]. The fundamental electrochemical properties of the Mo center may be investigated with well-characterized Modithiolate complexes, such as those with a [(Tp*)MoVO]2 center (Tp* /hydrotris(3,5-dimethyl1-pyrazolyl)borate) that mimic the Mo(V) active site in the sulfite oxidase family of molybdoenzymes [1,15]. Information regarding oxygen atom transfer (OAT) activity and Mo electron transfer reactions has been provided through spectroscopic studies of several (Tp*)MoO-dithiolate and (Tp*)WO-dithiolate complexes [14 /16]. Studying [(Tp*)MoII(NO)]2 analogues provides further insight due to redox activity involving the same Mo orbital as shown in the splitting diagram for the metal t2g orbitals for [MoNO]3 and [MoO]3 cores (Fig. 2) [17 /19]. Reduction of [MoNO]3 involves adding an electron to the empty Mo (dxy)0 orbital while reducing [MoO]3 necessitates the addition of an electron to Mo (dxy)1. The relative energy level of this
0020-1693/02/$ - see front matter # 2002 Elsevier Science B.V. All rights reserved. PII: S 0 0 2 0 - 1 6 9 3 ( 0 2 ) 0 1 2 0 6 - 9
86
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
Fig. 1. Structure of the pyranopterin-ene-dithiolate ligand of molybdoenzymes in its protonated form [3 /13].
reversible electrochemistry for one-electron reduction of the metal center. However, only (Tp*)MoO(bdt), (Tp*)MoO(tdt), and (Tp*)WO(tdt) could be reversibly oxidized in these experiments. The electrochemical potentials are highly dependent upon the nature of the ligands and the particular metal center, but the electron transfer kinetics of the complexes are similar to one another. In combination with previous studies of Modithiolate complexes, this work supports the special role of the ene-dithiolate ligand in modulating electron transfer reactions at the Mo center.
Fig. 2. Splitting diagram for the metal t2g orbitals of the [Mo(NO)]3 and [MoO]3 cores [17,18].
Mo redox orbital is determined by the nature of the other ligands. Electrochemistry is a powerful probe into the thermodynamics and kinetics of electron transfer reactions, which play key roles in the catalytic cycles of Mo enzymes. Investigating the variation of reduction potentials and heterogeneous electron transfer rate constants among closely related model complexes provides fundamental insight concerning the electron transfer reactions at the metal center. For example, Olson and Schultz studied the electrochemical differences among a series of (Tp*)MoO(X /Y) complexes, in which oxygen donor atoms were systematically replaced by sulfur donor atoms [20]. They observed that sulfur ligation of the oxo-Mo(V) center raises reduction potentials for the Mo(V/IV) reaction and facilitates higher rates of heterogeneous electron transfer. We recently reported that electrochemical potentials are strongly dependent upon the nature of the remote substituent among a series of oxo-Mo(V) complexes with phenoxide ligands and that the electron transfer rate constants correlate in a nearly linear manner with the reduction potentials [21]. Since the driving force for electron transfer is zero under electrochemical conditions, the magnitude of reorganization energy has a significant effect on heterogeneous electron transfer rates [22 /25]. In this study, electrochemical potentials and heterogeneous electron transfer rate constants were determined via cyclic voltammetric and chronocoulometric techniques for the dithiolate complexes shown (Fig. 3) to determine how variations of the axial ligand, benzene ring substituents, and metal center influence the electron transfer reactions. All of the complexes exhibited
Fig. 3. Structure of the [(Tp*)ME]2 system (M /Mo, W; E /O, NO) with coordinated equatorial ene-1,2-dithiolate ligands (S S S).
2. Experimental 2.1. Synthesis and characterization Reactions and operations required for preparing the compounds were performed under strict anaerobic conditions obtained by blanketing synthetic manipulations with pre-purified Ar gas and by utilizing standard Schlenk techniques, a high-vacuum/gas double line manifold, and an inert atmospheric glove bag. Glassware was oven-dried at 150 8C and repeatedly flushed with inert gas prior to use. The employed organic solvents were purified following standard procedures, distilled under nitrogen gas, degassed by freeze-thawpump cycles, and transferred to reaction vessels under inert gas via steel cannulae techniques. Reagents used in synthetic procedures were generally used as received. The following reagents (Aldrich) were dried and/or distilled in vacuo and stored under nitrogen gas prior to use: H2bdt (1,2-benzenedithiol), H2tdt (3,4toluenedithiol), and H2bdtCl2 (3,6-dichloro-1,2-benzenedithiol). Potassium hydrotris(3,5-dimethyl-1-pyrazolyl)borate (KTp*) [26] and the precursor complexes (Tp*)MoVOCl2 [26], (Tp*)MoII(NO)I2 [27], and (Tp*)WVOI2 [16,28] were prepared according to the literature. The compounds (Tp*)MoVO(bdt) [29,30], (Tp*)MoVO(tdt) [26], and (Tp*)MoVO(bdtCl2) [15] were synthesized as previously described. A modified procedure was employed for the preparation of (Tp*)MoII(NO)(tdt) [31] and, subsequently, (Tp*)MoII(NO)(bdt) and (Tp*)MoII(NO)(bdtCl2) [32,33]. The synthesis, isolation, purification, and characterization of (Tp*)WVO(tdt) involved procedures developed by Young and co-workers [16,28]. Reaction progress was monitored and product purity was confirmed by thin-layer chromatography (silica gel 60 F254 plastic sheets, EM Science). All of the dithiolate compounds were purified under argon by column chromatography (silica gel, 230/400 mesh) prior to spectroscopic and electrochemical characterization as previously discussed in the literature [15,16,26,28,33]. Multiple physical methods were utilized in the characterization of the compounds. Electronic absorption
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
spectra of samples solvated in 1,2-dichloroethane were collected on a modified Cary 14 (with OLIS interface, 250 /2600 nm) spectrophotometer. Infrared (IR) spectra (4000 /400 cm 1) were acquired in KBr disks or as dichloromethane solutions (between NaCl plates) on a Nicolet Avatar ESP 360 FT-IR spectrophotometer. Mass spectra were recorded on a JEOL HX110 highresolution sector instrument utilizing fast atom bombardment (FAB) ionization in a matrix of 3-nitrobenzyl alcohol (NBA). Electronic paramagnetic resonance (EPR) spectra of fluid solutions (298 K) or of frozen glasses (77 K) in dry degassed toluene were acquired at X-band frequency (approximately 9.1 GHz) with a Bruker ESP 300 spectrometer for those complexes containing a [(Tp*)MVO]2 (formally d1) center. 1H NMR spectra of the diamagnetic [(Tp*)MoII(NO)]2 (formally d4) complexes were measured on a Bruker DRX-500 spectrometer. The results of these spectroscopic studies were consistent with previous data and confirmed the purity of the samples submitted for electrochemical characterization [14 /16,33].
87
CV was used to determine electrochemical potentials for the reduction (and, in some cases, the oxidation) of each neutral compound. With the initial and final experimental potentials set 9/0.3 V of the redox potential, CC was performed over a range of pulse widths from 250 to 950 ms (in intervals of 100 ms). The slopes of the resulting Anson plots (charge vs. square root of time) were averaged and used to calculate diffusion coefficients upon application of the Cottrell equation [34]. The value for the electrode area employed in these calculations (0.022 cm2) was determined using CC data for ferrocene and the known diffusion coefficient of ferrocene in 1,2-DCE [35]. For each electron transfer reaction of interest, data from a series of CVs were recorded at scan rates ranging from 0.1 to 2 V s 1 (in intervals of 0.1 V s 1). Using the Nicholson method, the resulting peak-to-peak separations (DEp) at each scan rate were converted to a kinetic parameter (C) and heterogeneous electron transfer rate constants were calculated [36,37]. The diffusion coefficients described above were included in these rate constant calculations.
2.2. Electrochemical procedures Cyclic voltammetry (CV) and chronocoulometry (CC) were conducted at room temperature with a Bioanalytical Systems (BAS) CV-50W potentiostat. BAS supplied software provided scan acquisition control and data analysis capabilities. The 1,2-dichloroethane (DCE) employed for electrochemistry was of anhydrous grade (EM Science, DriSolv) and required no further purification. Electrochemical measurements were performed on degassed DCE sample solutions (0.5 /1.0 mM) over the potential range 9/1.5 V (vs. Ag/AgCl reference electrode, BAS) at a platinum-disk electrode (1.6 mm diameter, BAS). The platinum-disk electrodes were polished with 0.05 mm alumina (Buehler) and electrochemically cleaned in dilute H2SO4 prior to use. Solutions contained 0.1 /0.2 M dried tetra-n-butylammonium tetrafluroborate ([Bu4N][BF4], Aldrich) as the supporting electrolyte, and a platinum-wire auxiliary electrode (BAS) was used. CV scans of electrolyte solutions showed no indication of impurities within the utilized potential range. Ferrocene (Aldrich) was added to each solution upon completion of the experiments, and potentials are reported with respect to the ferrocene/ferrocenium (Fc/Fc ) redox couple. Each experiment was conducted at least twice for reproducibility purposes. Electrochemical experiments were also repeated using a homemade platinum-disk microelectrode (approximately 200 mm diameter) on MeCN sample solutions (0.5 /1.0 mM) with an EG&G potentiostat (Model 283, Princeton Applied Research). The MeCN (Aldrich) was freshly distilled and passed through a column of alumina immediately prior to use.
3. Results and discussion 3.1. Electrochemical potentials Cyclic voltammograms for all of the complexes displayed a well-resolved sets of peaks corresponding to a one-electron quasi-reversible reduction of the neutral molecule (Table 1). The CVs for (Tp*)MoO(bdt), (Tp*)MoO(tdt), and (Tp*)WO(tdt) also revealed a one-electron quasi-reversible oxidation process (example shown in Fig. 4) whereas oxidation Table 1 Reduction potentials (vs. Fc/Fc )
a
Compound
Mo(V/IV) E1/2 (mV)
Mo(VI/V) E1/2 (mV)
(Tp*)MoO(bdtCl2) (Tp*)MoO(bdt) (Tp*)MoO(tdt)
508 642 650
742 (Epa only) 534 462
Mo(II/I) E1/2 (mV) Irreversible oxidation (mV) (Tp*)MoNO(bdtCl2) 723 (Tp*)MoNO(bdt) 862 (Tp*)MoNO(tdt) 896
844 684 638
W(V/IV) E1/2 (mV) W(VI/V) E1/2 (mV) (Tp*)WO(tdt) a
1032
262
Experimental conditions for cyclic voltammetry at Pt-disk electrode: 100 mV s 1, 0.5 /1.0 mM solutions of sample in 0.1 /0.2 M [Bu4][NBF4]/1,2-DCE. Estimated error 92 mV.
88
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
Fig. 4. Cyclic voltammogram of (Tp*)MoO(tdt). Experimental conditions at Pt-disk electrode with Ag/AgCl reference: 100 mV s 1, 0.72 mM sample, 0.13 M Bu4NBF4 in 1,2-dichloroethane.
appeared to be electrochemically irreversible for the other four complexes within the experimental potential range. Potentials for the three oxo-Mo compounds were recently reported with the synthesis and characterization of (Tp*)MoO(bdtCl2) [15]. The reactions were assigned as oxidations or reductions based on whether the observed redox potential (E1/2, the average of Epa and Epc) was positive or negative, respectively, of the measured rest potential for the sample solution. The electron transfer processes were determined to be quasireversible based on the data approaching theoretical reversibility criteria at relatively slow CV scan rates (0.1 V s 1): anodic and cathodic peak separation (Ep) approximately 0.059 V; ratio of anodic and cathodic peak heights (ipa/ipc) approximately 1; and linear dependence of current flow on the square root of the scan rate (ip/n1/2) [34,38]. For the oxo-Mo(V) complexes, one-electron oxidation and reduction processes correspond to Mo(VI/V) and Mo(V/IV), respectively, assuming that the electron transfer reactions are metal-based in nature. Other possibilities for the oxidative process, involving sulfur ligand redox activity, have been previously discussed [15]. Adding substituents to the benzene ring of the dithiolate ligand greatly affects the redox potentials of the Mo center. Fundamental molecular orbital aspects of these potential changes have been discussed in detail [15]. The electron-withdrawing ability of the two chlorine atoms on the (bdtCl2) ligand raises the Mo(V/IV) potential by approximately 130 mV, facilitating the reduction process. The Mo(VI/V) potential is shifted positively by more than 200 mV, making the CV peaks convoluted with the positive potential edge of the
useable solvent window. Changing the (bdt) ligand to the (tdt) ligand, on the other hand, slightly lowers the Mo oxidation and reduction potentials due to the electron-donating character of the remote methyl group. A parallel influence of benzene ring variations on reduction potentials was observed within the series of nitrosyl-Mo(II) complexes. All of the reduction potentials for Mo complexes with nitrosyl axial ligands are more negative than those with oxo axial ligands. For (Tp*)Mo(NO)(bdtCl2), (Tp*)Mo(NO)(bdt), and (Tp*)Mo(NO)(tdt), reduction involves adding an electron to a formally d4 metal, a process that is less favorable than reducing the formally d1 metal of an oxo-Mo core, based on electrostatic arguments. The ill-defined CV peaks for the oxidative process of each nitrosyl complex appear at more positive potentials than the observed oxidation potential for the corresponding oxo complex. For (Tp*)WO(tdt), potentials for the (VI/V) and (V/ IV) reactions are considerably more negative than those for (Tp*)MoO(tdt). Tungsten in relatively high oxidation states is typically more difficult to reduce than Mo in corresponding compounds [39]. Thus, oxidizing the neutral complex is thermodynamically easier for the oxo-W(V) center than for Mo(V), while reduction is much more difficult for tungsten. 3.2. Heterogeneous electron transfer kinetics The heterogeneous rate constants determined for the dithiolate complexes in this study show minor variations with changes in ligand and metal center (Table 2). The individual values are all within an order of magnitude of
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
each other and are significantly smaller than that found for (Tp*)MoO(tdt) in a related study by Olson and Schultz [20]. The apparently slower electron transfer kinetics measured for these systems can be attributed to the effects of uncompensated solution resistance, which was not fully accounted for in this study [40,41]. Nevertheless, the rate constants in Table 2 can be compared to one another because all of the experiments were conducted under identical conditions and the calculations were performed in the same manner. The precision of the reported values is 9/10%, based on the reproducibility of results from multiple experiments under the same experimental conditions. Microelectrodes are typically expected to improve electron transfer kinetic measurements due to reduced solution resistance [42,43]; however, rate constant values obtained using a microelectrode were not statistically different from those in Table 2. Many systems have been reported in which redox potentials and heterogeneous electron transfer rates of metal-centered complexes are influenced in a parallel manner by donor atoms and substituent groups [20,23,44,45]. For example, we observed that rate constants increase systematically with E1/2 for the Mo(VI/V) and Mo(V/IV) reactions of (Tp*)MoOCl(p OC6H4X) and (Tp*)MoO(p -OC6H4X)2 (X /OEt, OMe, Et, Me, H, F, Cl, Br, I, and CN) carried out in MeCN [21]. In the present investigation, however, there is no clear correlation between electrochemical potentials and kinetic parameters. The rate of electron transfer within the complexes appears to be insensitive to the variations of the axial ligand, benzene ring substituents, and metal center. It is interesting to note that the substitution of sulfur for the equatorial oxygens in (Tp*)MoO(catecholate) has been shown to result in a significant shift to more positive reduction potentials and greater electron transfer rates [20]. However, the rate of electron transfer in the phenoxide complexes [21] are relatively equal to or somewhat larger than those for Table 2 Heterogeneous electron transfer rate constants in cm s 1 Mo(V/IV)
Mo(VI/V)
(Tp*)MoO(bdtCl2) (Tp*)MoO(bdt) (Tp*)MoO(tdt)
0.020 0.015 0.033
0.0097 0.026
Mo(II/I)
Irreversible oxidation
(Tp*)WO(tdt) a
the ene-1,2-dithiolate complexes in the present work. This unexpected effect may be attributed to the more resistive 1,2-dichloroethane employed as the solvent in this study, which results in a larger uncompensated solution resistance, and hence, slower electron transfer rates. Evidence for a unique Mo /sulfur bonding interaction in (Tp*)MoO(bdt) and (Tp*)MoO(tdt) has been provided via electronic absorption, magnetic circular dichroism, and resonance Raman spectroscopic studies [14]. It has been proposed that a three-centered psuedos type bond between the redox active Mo orbital and dithiolate ligand orbitals results in a highly covalent Mo/S interaction (Fig. 5) [14]. This covalency modulates the electrochemical potential of the Mo (and, similarly, W [16]) core, and as a result of the in-plane nature of this interaction, contributions to the reorganizational energy involving distortions along the apical Mo/O bond are anticipated to be minimized within the complex during electron transfer reactions [14]. Marcus theory predicts slower rates of electron transfer when large reorganization energies result from molecular rearrangements [46,47]. In complexes where the redox orbital is highly metal-centered, such as in the (Tp*)MoOCl(p -OC6H4X) and (Tp*)MoO(p -OC6H4X)2 complexes with oxygen and chlorine ligand orbitals [21], metal oxidation state changes are expected to be accompanied by more bond length adjustments in the coordination environment. Also, in addition to differences in metal /ligand bond distances, reorganization energies are influenced by force constants (larger for M /O vs. M /S) [20]. These effects of inner-shell reorganization energy can lead to significant electron
a
Compound
(Tp*)MoNO(bdtCl2) (Tp*)MoNO(bdt) (Tp*)MoNO(tdt)
89
0.015 0.018 0.0083 W(V/IV)
W(VI/V)
0.017
0.012
Estimated precision from multiple experiments is 910%.
Fig. 5. Molecular orbital diagram showing the psuedo-s bonding and anti-bonding interactions of the Mo dxy orbital and dithiolate ligand orbitals [14].
90
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
transfer rate variations. The small amount of reorganization energy for electron transfer within the dithiolate complexes due to more covalent metal/ligand interactions facilitates the reaction, making rate constants essentially equivalent for this series of compounds. The smaller range of electron transfer rate values observed in the dithiolate systems (factor of 3) versus the phenoxide complexes (Ref. [21], factor of 10) is consistent with smaller differences in reorganization energy for the dithiolate complexes, and hence greater covalency in the metal/ligand bonds. PES studies of (Tp*)Mo(E)(tdt), where E /O, S, and NO, and of (Tp*)MoO(qdt), where (qdt) /2,3-quinoxalinedithiolate, have also revealed substantial Mo /S covalency and have indicated that Mo oxidation state changes are offset by the bonding properties of the dithiolate ligand [19,48]. This provides further explanation for the herein reported similarity of heterogeneous electron transfer rate constants with the variations of ligand and metal among the investigated dithiolate complexes.
Acknowledgements We thank Dr. Franklin Schultz for helpful discussions, Dr. Neal Armstrong for the use of his laboratory for the microelectrode experiments, and Mr. Chet Carter for assistance with those experiments. Mass spectra were recorded at the University of Arizona Mass Spectrometry Facility, and EPR studies were conducted under the direction of Dr. Arnold Raitsimring at the University of Arizona. Financial support by the National Institutes of Health (GM 37773 to J.H.E.) is gratefully acknowledged.
References [1] R. Hille, Chem. Rev. 96 (1996) 2757. [2] K.V. Rajagopalan, in: M. Coughlan (Ed.), Molybdenum and Molybdenum Containing Enzymes, Pergamon Press, New York, 1980, pp. 243 /272. [3] C. Kisker, H. Schindelin, A. Pacheco, W.A. Wehbi, R.M. Garrett, K.V. Rajagopalan, J.H. Enemark, D.C. Rees, Cell 91 (1997) 973. [4] H. Schindelin, C. Kisker, J. Huber, K. Rajagopalan, D. Rees, Science 272 (1996) 1615. [5] H. Li, C. Temple, K.V. Rajagopalan, H. Schindelin, J. Am. Chem. Soc. 122 (2000) 7673. [6] F. Schneider, J. Lowe, R. Huber, H. Schindelin, C. Kisker, J. Knablein, J. Mol. Biol. 263 (1996) 53. [7] A.S. McAlpine, A.G. McEwan, S. Bailey, J. Mol. Biol. 275 (1998) 613. [8] A.S. McAlpine, A.G. McEwan, A.L. Shaw, S. Bailey, J. Biol. Inorg. Chem. 2 (1997) 634. [9] R.C. Bray, B. Adams, A.T. Smith, B. Bennett, S. Bailey, Biochem. 39 (2000) 11258. [10] L.J. Stewart, S. Bailey, B. Bennett, J.M. Charnock, C.D. Garner, A.S. McAlpine, J. Mol. Biol. 299 (2000) 593.
[11] M.J. Roma˜o, M. Archer, I. Moura, J.J.G. Moura, J. LeGall, R. Engh, M. Schneider, P. Hof, R. Huber, Science 270 (1995) 1170. [12] R. Huber, P. Hof, R.O. Duarte, J.J.G. Moura, I. Moura, J. LeGall, R. Hille, M. Archer, M. Roma˜o, Proc. Natl. Acad. Sci. USA 93 (1996) 8846. [13] J.C. Boyington, V.N. Gladyshev, S.V. Khangulaov, T.C. Stadtman, P.D. Sun, Science 275 (1997) 1305. [14] F.E. Inscore, R. McNaughton, B.L. Westcott, M.E. Helton, R. Jones, I.K. Dhawan, J.H. Enemark, M.L. Kirk, Inorg. Chem. 38 (1999) 1401. [15] F.E. Inscore, H.K. Joshi, A.E. McElhaney, J.H. Enemark, Inorg. Chim. Acta 331 (2002) 246. [16] F.E. Inscore, J.P. Hill, C.G. Young, M.L. Kirk, manuscript in preparation. [17] B.L. Westcott, J.H. Enemark, Inorg. Chem. 36 (1997) 5404. [18] B.E. Bursten, R.H. Cayton, Organometallics 6 (1987) 2004. [19] B.L. Westcott, N.E. Gruhn, J.H. Enemark, J. Am. Chem. Soc. 120 (1998) 3382. [20] G.M. Olson, F.A. Schultz, Inorg. Chim. Acta 225 (1994) 1. [21] J.N. Graff, A.E. McElhaney, P. Basu, N.E. Gruhn, C.J. Chang, J.H. Enemark, Inorg. Chem. 41 (2002) 2642. [22] S.J. Lippard, J.M. Berg, Principles of Bioinorganic Chemistry, University Science Books, Mill Valley, CA, 1994. [23] F.A. Schultz, private communication. [24] B.S. Brunschwig, N. Sutin, Coord. Chem. Rev. 187 (1999) 233. [25] N. Sutin, B.S. Brunschwig, C. Creutz, J.R. Winkler, Pure Appl. Chem. 60 (1988) 1817. [26] W.E. Cleland, Jr., K.M. Barnhart, K. Yamanouchi, D. Collison, F.E. Mabbs, R.B. Ortega, J.H. Enemark, Inorg. Chem. 26 (1987) 1017. [27] S.J. Reynolds, C.F. Smith, C.J. Jones, J.A. McCleverty, Inorg. Synth. 23 (1985) 4. [28] C.G. Young, private communication. [29] I.K. Dhawan, A. Pacheco, J.H. Enemark, J. Am. Chem. Soc. 116 (1994) 7911. [30] I.K. Dhawan, J.H. Enemark, Inorg. Chem. 35 (1996) 4873. [31] N. Alobaidi, C.J. Jones, J.A. McCleverty, Polyhedron 8 (1989) 1033. [32] I.K. Dhawan, Ph.D. dissertation, The University of Arizona, Tucson, AZ, 1995. [33] H.K. Joshi, F.E. Inscore, J.T. Schirlin, I.K. Dhawan, M.D. Carducci, T.G. Bill, J.H. Enemark, Inorg. Chim. Acta 337 (2002) 275. [34] A.J. Bard, L.R. Faulkner, Electrochemical Methods, John Wiley, New York, 1980. [35] K.M. Kadish, J.Q. Ding, T. Malinkski, Anal. Chem. 56 (1984) 1741. [36] R.S. Nicholson, Anal. Chem. 37 (1965) 1351. [37] R.S. Nicholson, I. Shain, Anal. Chem. 36 (1964) 706. [38] D.T. Sawyer, A. Sobkowiak, J.L. Roberts, Electrochemistry for Chemists, John Wiley, New York, 1995. [39] F.A. Cotton, G. Wilkinson, C. Murillo, M. Bochmann, R. Grimes, Advanced Inorganic Chemistry, John Wiley, New York, 1999. [40] C.P. Andrieux, D. Garreau, P. Hapiiot, J.M. Saveant, J. Electroanal. Chem. Interfacial Electrochem. 248 (1988) 447. [41] J.O. Howell, R.M. Wightman, J. Phys. Chem. 88 (1984) 3915. [42] R.M. Wightman, Science 240 (1988) 415. [43] D.O. Wipf, E.W. Kristensen, M.R. Deakin, R.M. Wightman, Anal. Chem. 60 (1988) 306. [44] J.U. Mondal, J.G. Zamora, S. Siew, G.T. Garcia, E.R. George, M.D. Kinon, F.A. Schultz, Inorg. Chim. Acta 321 (2001) 83. [45] D. Lexa, J.M. Saveant, Acc. Chem. Res. 16 (1983) 235. [46] R.A. Marcus, J. Chem. Phys. 24 (1956) 966. [47] R.A. Marcus, Electrochim. Acta 13 (1968) 995. [48] M.E. Helton, N.E. Gruhn, R.L. McNaughton, M.L. Kirk, Inorg. Chem. 39 (2000) 2273.
Electron transfer studies of dithiolate complexes: effects of ligand variation and metal substitution Anne E. McElhaney, Frank E. Inscore, Julien T. Schirlin, John H. Enemark * Department of Chemistry, The University of Arizona, Tucson, AZ 85721-0041, USA Received 28 April 2002; accepted 31 July 2002 Dedicated to Professor Kenneth N. Raymond on the occasion of his 60th birthday
Abstract Solution redox potentials and heterogeneous electron transfer rate constants have been measured for mono-ene-1,2-dithiolate complexes of the type (Tp*)M(E)(S S S) [Tp* /hydrotris(3,5-dimethyl-1-pyrazolyl)borate; M /Mo, W; E /O, NO]. The dithiolate ligands (S S S) equatorially coordinated to the central metal include: 1,2-benzenedithiolate (bdt); 3,6-dichloro-1,2-benzenedithiolate (bdtCl2); and 3,4-toluenedithiolate (tdt). Cyclic voltammograms reveal quasi-reversible one-electron reductions for all of the compounds; and, a one-electron quasi-reversible oxidation process is also observed for (Tp*)MoO(bdt), (Tp*)MoO(tdt), and (Tp*)WO(tdt). Electrochemical potentials are strongly dependent upon the nature of the substituents on the benzene ring of the dithiolate ligand. The nitrosyl /molybdenum complexes (E /NO) are more difficult to reduce than their corresponding oxomolybdenum complexes. The tungsten compound, (Tp*)WO(tdt), has the most negative electrochemical potentials among the complexes investigated. Heterogeneous electron transfer rates are insensitive to the variations of the central metal and ligands. These results support a highly covalent bonding interaction between the metal and the dithiolate ligands that modulates electron transfer reactions within these compounds. # 2002 Elsevier Science B.V. All rights reserved. Keywords: Electrochemistry; Molybdenum; Dithiolate; Heterogeneous electron transfer kinetics
1. Introduction Molybdenum-containing enzymes play essential metabolic roles in animals, plants, and microorganisms by catalyzing two-electron oxidations or reductions of various substrate molecules [1,2]. The molybdenum (Mo) center experiences oxidation states of VI, V, and IV during catalysis, and the initial oxidation state of the Mo active site is regenerated via two one-electron transfer reactions. Crystal structures obtained for several of these enzymes reveal common active site structural features, namely one (or two) pyranopterinene-1,2-dithiolate ligand and at least one terminal oxo ligand [3 /13]. The ene-dithiolate (Fig. 1) is believed to modulate electron transfer reactions within these en-
* Corresponding author. Tel.: /1-520-621 2245; fax: /1-520-626 8065 E-mail address: [email protected] (J.H. Enemark).
zymes, ensuring an efficient catalytic process [1,14]. The fundamental electrochemical properties of the Mo center may be investigated with well-characterized Modithiolate complexes, such as those with a [(Tp*)MoVO]2 center (Tp* /hydrotris(3,5-dimethyl1-pyrazolyl)borate) that mimic the Mo(V) active site in the sulfite oxidase family of molybdoenzymes [1,15]. Information regarding oxygen atom transfer (OAT) activity and Mo electron transfer reactions has been provided through spectroscopic studies of several (Tp*)MoO-dithiolate and (Tp*)WO-dithiolate complexes [14 /16]. Studying [(Tp*)MoII(NO)]2 analogues provides further insight due to redox activity involving the same Mo orbital as shown in the splitting diagram for the metal t2g orbitals for [MoNO]3 and [MoO]3 cores (Fig. 2) [17 /19]. Reduction of [MoNO]3 involves adding an electron to the empty Mo (dxy)0 orbital while reducing [MoO]3 necessitates the addition of an electron to Mo (dxy)1. The relative energy level of this
0020-1693/02/$ - see front matter # 2002 Elsevier Science B.V. All rights reserved. PII: S 0 0 2 0 - 1 6 9 3 ( 0 2 ) 0 1 2 0 6 - 9
86
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
Fig. 1. Structure of the pyranopterin-ene-dithiolate ligand of molybdoenzymes in its protonated form [3 /13].
reversible electrochemistry for one-electron reduction of the metal center. However, only (Tp*)MoO(bdt), (Tp*)MoO(tdt), and (Tp*)WO(tdt) could be reversibly oxidized in these experiments. The electrochemical potentials are highly dependent upon the nature of the ligands and the particular metal center, but the electron transfer kinetics of the complexes are similar to one another. In combination with previous studies of Modithiolate complexes, this work supports the special role of the ene-dithiolate ligand in modulating electron transfer reactions at the Mo center.
Fig. 2. Splitting diagram for the metal t2g orbitals of the [Mo(NO)]3 and [MoO]3 cores [17,18].
Mo redox orbital is determined by the nature of the other ligands. Electrochemistry is a powerful probe into the thermodynamics and kinetics of electron transfer reactions, which play key roles in the catalytic cycles of Mo enzymes. Investigating the variation of reduction potentials and heterogeneous electron transfer rate constants among closely related model complexes provides fundamental insight concerning the electron transfer reactions at the metal center. For example, Olson and Schultz studied the electrochemical differences among a series of (Tp*)MoO(X /Y) complexes, in which oxygen donor atoms were systematically replaced by sulfur donor atoms [20]. They observed that sulfur ligation of the oxo-Mo(V) center raises reduction potentials for the Mo(V/IV) reaction and facilitates higher rates of heterogeneous electron transfer. We recently reported that electrochemical potentials are strongly dependent upon the nature of the remote substituent among a series of oxo-Mo(V) complexes with phenoxide ligands and that the electron transfer rate constants correlate in a nearly linear manner with the reduction potentials [21]. Since the driving force for electron transfer is zero under electrochemical conditions, the magnitude of reorganization energy has a significant effect on heterogeneous electron transfer rates [22 /25]. In this study, electrochemical potentials and heterogeneous electron transfer rate constants were determined via cyclic voltammetric and chronocoulometric techniques for the dithiolate complexes shown (Fig. 3) to determine how variations of the axial ligand, benzene ring substituents, and metal center influence the electron transfer reactions. All of the complexes exhibited
Fig. 3. Structure of the [(Tp*)ME]2 system (M /Mo, W; E /O, NO) with coordinated equatorial ene-1,2-dithiolate ligands (S S S).
2. Experimental 2.1. Synthesis and characterization Reactions and operations required for preparing the compounds were performed under strict anaerobic conditions obtained by blanketing synthetic manipulations with pre-purified Ar gas and by utilizing standard Schlenk techniques, a high-vacuum/gas double line manifold, and an inert atmospheric glove bag. Glassware was oven-dried at 150 8C and repeatedly flushed with inert gas prior to use. The employed organic solvents were purified following standard procedures, distilled under nitrogen gas, degassed by freeze-thawpump cycles, and transferred to reaction vessels under inert gas via steel cannulae techniques. Reagents used in synthetic procedures were generally used as received. The following reagents (Aldrich) were dried and/or distilled in vacuo and stored under nitrogen gas prior to use: H2bdt (1,2-benzenedithiol), H2tdt (3,4toluenedithiol), and H2bdtCl2 (3,6-dichloro-1,2-benzenedithiol). Potassium hydrotris(3,5-dimethyl-1-pyrazolyl)borate (KTp*) [26] and the precursor complexes (Tp*)MoVOCl2 [26], (Tp*)MoII(NO)I2 [27], and (Tp*)WVOI2 [16,28] were prepared according to the literature. The compounds (Tp*)MoVO(bdt) [29,30], (Tp*)MoVO(tdt) [26], and (Tp*)MoVO(bdtCl2) [15] were synthesized as previously described. A modified procedure was employed for the preparation of (Tp*)MoII(NO)(tdt) [31] and, subsequently, (Tp*)MoII(NO)(bdt) and (Tp*)MoII(NO)(bdtCl2) [32,33]. The synthesis, isolation, purification, and characterization of (Tp*)WVO(tdt) involved procedures developed by Young and co-workers [16,28]. Reaction progress was monitored and product purity was confirmed by thin-layer chromatography (silica gel 60 F254 plastic sheets, EM Science). All of the dithiolate compounds were purified under argon by column chromatography (silica gel, 230/400 mesh) prior to spectroscopic and electrochemical characterization as previously discussed in the literature [15,16,26,28,33]. Multiple physical methods were utilized in the characterization of the compounds. Electronic absorption
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
spectra of samples solvated in 1,2-dichloroethane were collected on a modified Cary 14 (with OLIS interface, 250 /2600 nm) spectrophotometer. Infrared (IR) spectra (4000 /400 cm 1) were acquired in KBr disks or as dichloromethane solutions (between NaCl plates) on a Nicolet Avatar ESP 360 FT-IR spectrophotometer. Mass spectra were recorded on a JEOL HX110 highresolution sector instrument utilizing fast atom bombardment (FAB) ionization in a matrix of 3-nitrobenzyl alcohol (NBA). Electronic paramagnetic resonance (EPR) spectra of fluid solutions (298 K) or of frozen glasses (77 K) in dry degassed toluene were acquired at X-band frequency (approximately 9.1 GHz) with a Bruker ESP 300 spectrometer for those complexes containing a [(Tp*)MVO]2 (formally d1) center. 1H NMR spectra of the diamagnetic [(Tp*)MoII(NO)]2 (formally d4) complexes were measured on a Bruker DRX-500 spectrometer. The results of these spectroscopic studies were consistent with previous data and confirmed the purity of the samples submitted for electrochemical characterization [14 /16,33].
87
CV was used to determine electrochemical potentials for the reduction (and, in some cases, the oxidation) of each neutral compound. With the initial and final experimental potentials set 9/0.3 V of the redox potential, CC was performed over a range of pulse widths from 250 to 950 ms (in intervals of 100 ms). The slopes of the resulting Anson plots (charge vs. square root of time) were averaged and used to calculate diffusion coefficients upon application of the Cottrell equation [34]. The value for the electrode area employed in these calculations (0.022 cm2) was determined using CC data for ferrocene and the known diffusion coefficient of ferrocene in 1,2-DCE [35]. For each electron transfer reaction of interest, data from a series of CVs were recorded at scan rates ranging from 0.1 to 2 V s 1 (in intervals of 0.1 V s 1). Using the Nicholson method, the resulting peak-to-peak separations (DEp) at each scan rate were converted to a kinetic parameter (C) and heterogeneous electron transfer rate constants were calculated [36,37]. The diffusion coefficients described above were included in these rate constant calculations.
2.2. Electrochemical procedures Cyclic voltammetry (CV) and chronocoulometry (CC) were conducted at room temperature with a Bioanalytical Systems (BAS) CV-50W potentiostat. BAS supplied software provided scan acquisition control and data analysis capabilities. The 1,2-dichloroethane (DCE) employed for electrochemistry was of anhydrous grade (EM Science, DriSolv) and required no further purification. Electrochemical measurements were performed on degassed DCE sample solutions (0.5 /1.0 mM) over the potential range 9/1.5 V (vs. Ag/AgCl reference electrode, BAS) at a platinum-disk electrode (1.6 mm diameter, BAS). The platinum-disk electrodes were polished with 0.05 mm alumina (Buehler) and electrochemically cleaned in dilute H2SO4 prior to use. Solutions contained 0.1 /0.2 M dried tetra-n-butylammonium tetrafluroborate ([Bu4N][BF4], Aldrich) as the supporting electrolyte, and a platinum-wire auxiliary electrode (BAS) was used. CV scans of electrolyte solutions showed no indication of impurities within the utilized potential range. Ferrocene (Aldrich) was added to each solution upon completion of the experiments, and potentials are reported with respect to the ferrocene/ferrocenium (Fc/Fc ) redox couple. Each experiment was conducted at least twice for reproducibility purposes. Electrochemical experiments were also repeated using a homemade platinum-disk microelectrode (approximately 200 mm diameter) on MeCN sample solutions (0.5 /1.0 mM) with an EG&G potentiostat (Model 283, Princeton Applied Research). The MeCN (Aldrich) was freshly distilled and passed through a column of alumina immediately prior to use.
3. Results and discussion 3.1. Electrochemical potentials Cyclic voltammograms for all of the complexes displayed a well-resolved sets of peaks corresponding to a one-electron quasi-reversible reduction of the neutral molecule (Table 1). The CVs for (Tp*)MoO(bdt), (Tp*)MoO(tdt), and (Tp*)WO(tdt) also revealed a one-electron quasi-reversible oxidation process (example shown in Fig. 4) whereas oxidation Table 1 Reduction potentials (vs. Fc/Fc )
a
Compound
Mo(V/IV) E1/2 (mV)
Mo(VI/V) E1/2 (mV)
(Tp*)MoO(bdtCl2) (Tp*)MoO(bdt) (Tp*)MoO(tdt)
508 642 650
742 (Epa only) 534 462
Mo(II/I) E1/2 (mV) Irreversible oxidation (mV) (Tp*)MoNO(bdtCl2) 723 (Tp*)MoNO(bdt) 862 (Tp*)MoNO(tdt) 896
844 684 638
W(V/IV) E1/2 (mV) W(VI/V) E1/2 (mV) (Tp*)WO(tdt) a
1032
262
Experimental conditions for cyclic voltammetry at Pt-disk electrode: 100 mV s 1, 0.5 /1.0 mM solutions of sample in 0.1 /0.2 M [Bu4][NBF4]/1,2-DCE. Estimated error 92 mV.
88
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
Fig. 4. Cyclic voltammogram of (Tp*)MoO(tdt). Experimental conditions at Pt-disk electrode with Ag/AgCl reference: 100 mV s 1, 0.72 mM sample, 0.13 M Bu4NBF4 in 1,2-dichloroethane.
appeared to be electrochemically irreversible for the other four complexes within the experimental potential range. Potentials for the three oxo-Mo compounds were recently reported with the synthesis and characterization of (Tp*)MoO(bdtCl2) [15]. The reactions were assigned as oxidations or reductions based on whether the observed redox potential (E1/2, the average of Epa and Epc) was positive or negative, respectively, of the measured rest potential for the sample solution. The electron transfer processes were determined to be quasireversible based on the data approaching theoretical reversibility criteria at relatively slow CV scan rates (0.1 V s 1): anodic and cathodic peak separation (Ep) approximately 0.059 V; ratio of anodic and cathodic peak heights (ipa/ipc) approximately 1; and linear dependence of current flow on the square root of the scan rate (ip/n1/2) [34,38]. For the oxo-Mo(V) complexes, one-electron oxidation and reduction processes correspond to Mo(VI/V) and Mo(V/IV), respectively, assuming that the electron transfer reactions are metal-based in nature. Other possibilities for the oxidative process, involving sulfur ligand redox activity, have been previously discussed [15]. Adding substituents to the benzene ring of the dithiolate ligand greatly affects the redox potentials of the Mo center. Fundamental molecular orbital aspects of these potential changes have been discussed in detail [15]. The electron-withdrawing ability of the two chlorine atoms on the (bdtCl2) ligand raises the Mo(V/IV) potential by approximately 130 mV, facilitating the reduction process. The Mo(VI/V) potential is shifted positively by more than 200 mV, making the CV peaks convoluted with the positive potential edge of the
useable solvent window. Changing the (bdt) ligand to the (tdt) ligand, on the other hand, slightly lowers the Mo oxidation and reduction potentials due to the electron-donating character of the remote methyl group. A parallel influence of benzene ring variations on reduction potentials was observed within the series of nitrosyl-Mo(II) complexes. All of the reduction potentials for Mo complexes with nitrosyl axial ligands are more negative than those with oxo axial ligands. For (Tp*)Mo(NO)(bdtCl2), (Tp*)Mo(NO)(bdt), and (Tp*)Mo(NO)(tdt), reduction involves adding an electron to a formally d4 metal, a process that is less favorable than reducing the formally d1 metal of an oxo-Mo core, based on electrostatic arguments. The ill-defined CV peaks for the oxidative process of each nitrosyl complex appear at more positive potentials than the observed oxidation potential for the corresponding oxo complex. For (Tp*)WO(tdt), potentials for the (VI/V) and (V/ IV) reactions are considerably more negative than those for (Tp*)MoO(tdt). Tungsten in relatively high oxidation states is typically more difficult to reduce than Mo in corresponding compounds [39]. Thus, oxidizing the neutral complex is thermodynamically easier for the oxo-W(V) center than for Mo(V), while reduction is much more difficult for tungsten. 3.2. Heterogeneous electron transfer kinetics The heterogeneous rate constants determined for the dithiolate complexes in this study show minor variations with changes in ligand and metal center (Table 2). The individual values are all within an order of magnitude of
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
each other and are significantly smaller than that found for (Tp*)MoO(tdt) in a related study by Olson and Schultz [20]. The apparently slower electron transfer kinetics measured for these systems can be attributed to the effects of uncompensated solution resistance, which was not fully accounted for in this study [40,41]. Nevertheless, the rate constants in Table 2 can be compared to one another because all of the experiments were conducted under identical conditions and the calculations were performed in the same manner. The precision of the reported values is 9/10%, based on the reproducibility of results from multiple experiments under the same experimental conditions. Microelectrodes are typically expected to improve electron transfer kinetic measurements due to reduced solution resistance [42,43]; however, rate constant values obtained using a microelectrode were not statistically different from those in Table 2. Many systems have been reported in which redox potentials and heterogeneous electron transfer rates of metal-centered complexes are influenced in a parallel manner by donor atoms and substituent groups [20,23,44,45]. For example, we observed that rate constants increase systematically with E1/2 for the Mo(VI/V) and Mo(V/IV) reactions of (Tp*)MoOCl(p OC6H4X) and (Tp*)MoO(p -OC6H4X)2 (X /OEt, OMe, Et, Me, H, F, Cl, Br, I, and CN) carried out in MeCN [21]. In the present investigation, however, there is no clear correlation between electrochemical potentials and kinetic parameters. The rate of electron transfer within the complexes appears to be insensitive to the variations of the axial ligand, benzene ring substituents, and metal center. It is interesting to note that the substitution of sulfur for the equatorial oxygens in (Tp*)MoO(catecholate) has been shown to result in a significant shift to more positive reduction potentials and greater electron transfer rates [20]. However, the rate of electron transfer in the phenoxide complexes [21] are relatively equal to or somewhat larger than those for Table 2 Heterogeneous electron transfer rate constants in cm s 1 Mo(V/IV)
Mo(VI/V)
(Tp*)MoO(bdtCl2) (Tp*)MoO(bdt) (Tp*)MoO(tdt)
0.020 0.015 0.033
0.0097 0.026
Mo(II/I)
Irreversible oxidation
(Tp*)WO(tdt) a
the ene-1,2-dithiolate complexes in the present work. This unexpected effect may be attributed to the more resistive 1,2-dichloroethane employed as the solvent in this study, which results in a larger uncompensated solution resistance, and hence, slower electron transfer rates. Evidence for a unique Mo /sulfur bonding interaction in (Tp*)MoO(bdt) and (Tp*)MoO(tdt) has been provided via electronic absorption, magnetic circular dichroism, and resonance Raman spectroscopic studies [14]. It has been proposed that a three-centered psuedos type bond between the redox active Mo orbital and dithiolate ligand orbitals results in a highly covalent Mo/S interaction (Fig. 5) [14]. This covalency modulates the electrochemical potential of the Mo (and, similarly, W [16]) core, and as a result of the in-plane nature of this interaction, contributions to the reorganizational energy involving distortions along the apical Mo/O bond are anticipated to be minimized within the complex during electron transfer reactions [14]. Marcus theory predicts slower rates of electron transfer when large reorganization energies result from molecular rearrangements [46,47]. In complexes where the redox orbital is highly metal-centered, such as in the (Tp*)MoOCl(p -OC6H4X) and (Tp*)MoO(p -OC6H4X)2 complexes with oxygen and chlorine ligand orbitals [21], metal oxidation state changes are expected to be accompanied by more bond length adjustments in the coordination environment. Also, in addition to differences in metal /ligand bond distances, reorganization energies are influenced by force constants (larger for M /O vs. M /S) [20]. These effects of inner-shell reorganization energy can lead to significant electron
a
Compound
(Tp*)MoNO(bdtCl2) (Tp*)MoNO(bdt) (Tp*)MoNO(tdt)
89
0.015 0.018 0.0083 W(V/IV)
W(VI/V)
0.017
0.012
Estimated precision from multiple experiments is 910%.
Fig. 5. Molecular orbital diagram showing the psuedo-s bonding and anti-bonding interactions of the Mo dxy orbital and dithiolate ligand orbitals [14].
90
A.E. McElhaney et al. / Inorganica Chimica Acta 341 (2002) 85 /90
transfer rate variations. The small amount of reorganization energy for electron transfer within the dithiolate complexes due to more covalent metal/ligand interactions facilitates the reaction, making rate constants essentially equivalent for this series of compounds. The smaller range of electron transfer rate values observed in the dithiolate systems (factor of 3) versus the phenoxide complexes (Ref. [21], factor of 10) is consistent with smaller differences in reorganization energy for the dithiolate complexes, and hence greater covalency in the metal/ligand bonds. PES studies of (Tp*)Mo(E)(tdt), where E /O, S, and NO, and of (Tp*)MoO(qdt), where (qdt) /2,3-quinoxalinedithiolate, have also revealed substantial Mo /S covalency and have indicated that Mo oxidation state changes are offset by the bonding properties of the dithiolate ligand [19,48]. This provides further explanation for the herein reported similarity of heterogeneous electron transfer rate constants with the variations of ligand and metal among the investigated dithiolate complexes.
Acknowledgements We thank Dr. Franklin Schultz for helpful discussions, Dr. Neal Armstrong for the use of his laboratory for the microelectrode experiments, and Mr. Chet Carter for assistance with those experiments. Mass spectra were recorded at the University of Arizona Mass Spectrometry Facility, and EPR studies were conducted under the direction of Dr. Arnold Raitsimring at the University of Arizona. Financial support by the National Institutes of Health (GM 37773 to J.H.E.) is gratefully acknowledged.
References [1] R. Hille, Chem. Rev. 96 (1996) 2757. [2] K.V. Rajagopalan, in: M. Coughlan (Ed.), Molybdenum and Molybdenum Containing Enzymes, Pergamon Press, New York, 1980, pp. 243 /272. [3] C. Kisker, H. Schindelin, A. Pacheco, W.A. Wehbi, R.M. Garrett, K.V. Rajagopalan, J.H. Enemark, D.C. Rees, Cell 91 (1997) 973. [4] H. Schindelin, C. Kisker, J. Huber, K. Rajagopalan, D. Rees, Science 272 (1996) 1615. [5] H. Li, C. Temple, K.V. Rajagopalan, H. Schindelin, J. Am. Chem. Soc. 122 (2000) 7673. [6] F. Schneider, J. Lowe, R. Huber, H. Schindelin, C. Kisker, J. Knablein, J. Mol. Biol. 263 (1996) 53. [7] A.S. McAlpine, A.G. McEwan, S. Bailey, J. Mol. Biol. 275 (1998) 613. [8] A.S. McAlpine, A.G. McEwan, A.L. Shaw, S. Bailey, J. Biol. Inorg. Chem. 2 (1997) 634. [9] R.C. Bray, B. Adams, A.T. Smith, B. Bennett, S. Bailey, Biochem. 39 (2000) 11258. [10] L.J. Stewart, S. Bailey, B. Bennett, J.M. Charnock, C.D. Garner, A.S. McAlpine, J. Mol. Biol. 299 (2000) 593.
[11] M.J. Roma˜o, M. Archer, I. Moura, J.J.G. Moura, J. LeGall, R. Engh, M. Schneider, P. Hof, R. Huber, Science 270 (1995) 1170. [12] R. Huber, P. Hof, R.O. Duarte, J.J.G. Moura, I. Moura, J. LeGall, R. Hille, M. Archer, M. Roma˜o, Proc. Natl. Acad. Sci. USA 93 (1996) 8846. [13] J.C. Boyington, V.N. Gladyshev, S.V. Khangulaov, T.C. Stadtman, P.D. Sun, Science 275 (1997) 1305. [14] F.E. Inscore, R. McNaughton, B.L. Westcott, M.E. Helton, R. Jones, I.K. Dhawan, J.H. Enemark, M.L. Kirk, Inorg. Chem. 38 (1999) 1401. [15] F.E. Inscore, H.K. Joshi, A.E. McElhaney, J.H. Enemark, Inorg. Chim. Acta 331 (2002) 246. [16] F.E. Inscore, J.P. Hill, C.G. Young, M.L. Kirk, manuscript in preparation. [17] B.L. Westcott, J.H. Enemark, Inorg. Chem. 36 (1997) 5404. [18] B.E. Bursten, R.H. Cayton, Organometallics 6 (1987) 2004. [19] B.L. Westcott, N.E. Gruhn, J.H. Enemark, J. Am. Chem. Soc. 120 (1998) 3382. [20] G.M. Olson, F.A. Schultz, Inorg. Chim. Acta 225 (1994) 1. [21] J.N. Graff, A.E. McElhaney, P. Basu, N.E. Gruhn, C.J. Chang, J.H. Enemark, Inorg. Chem. 41 (2002) 2642. [22] S.J. Lippard, J.M. Berg, Principles of Bioinorganic Chemistry, University Science Books, Mill Valley, CA, 1994. [23] F.A. Schultz, private communication. [24] B.S. Brunschwig, N. Sutin, Coord. Chem. Rev. 187 (1999) 233. [25] N. Sutin, B.S. Brunschwig, C. Creutz, J.R. Winkler, Pure Appl. Chem. 60 (1988) 1817. [26] W.E. Cleland, Jr., K.M. Barnhart, K. Yamanouchi, D. Collison, F.E. Mabbs, R.B. Ortega, J.H. Enemark, Inorg. Chem. 26 (1987) 1017. [27] S.J. Reynolds, C.F. Smith, C.J. Jones, J.A. McCleverty, Inorg. Synth. 23 (1985) 4. [28] C.G. Young, private communication. [29] I.K. Dhawan, A. Pacheco, J.H. Enemark, J. Am. Chem. Soc. 116 (1994) 7911. [30] I.K. Dhawan, J.H. Enemark, Inorg. Chem. 35 (1996) 4873. [31] N. Alobaidi, C.J. Jones, J.A. McCleverty, Polyhedron 8 (1989) 1033. [32] I.K. Dhawan, Ph.D. dissertation, The University of Arizona, Tucson, AZ, 1995. [33] H.K. Joshi, F.E. Inscore, J.T. Schirlin, I.K. Dhawan, M.D. Carducci, T.G. Bill, J.H. Enemark, Inorg. Chim. Acta 337 (2002) 275. [34] A.J. Bard, L.R. Faulkner, Electrochemical Methods, John Wiley, New York, 1980. [35] K.M. Kadish, J.Q. Ding, T. Malinkski, Anal. Chem. 56 (1984) 1741. [36] R.S. Nicholson, Anal. Chem. 37 (1965) 1351. [37] R.S. Nicholson, I. Shain, Anal. Chem. 36 (1964) 706. [38] D.T. Sawyer, A. Sobkowiak, J.L. Roberts, Electrochemistry for Chemists, John Wiley, New York, 1995. [39] F.A. Cotton, G. Wilkinson, C. Murillo, M. Bochmann, R. Grimes, Advanced Inorganic Chemistry, John Wiley, New York, 1999. [40] C.P. Andrieux, D. Garreau, P. Hapiiot, J.M. Saveant, J. Electroanal. Chem. Interfacial Electrochem. 248 (1988) 447. [41] J.O. Howell, R.M. Wightman, J. Phys. Chem. 88 (1984) 3915. [42] R.M. Wightman, Science 240 (1988) 415. [43] D.O. Wipf, E.W. Kristensen, M.R. Deakin, R.M. Wightman, Anal. Chem. 60 (1988) 306. [44] J.U. Mondal, J.G. Zamora, S. Siew, G.T. Garcia, E.R. George, M.D. Kinon, F.A. Schultz, Inorg. Chim. Acta 321 (2001) 83. [45] D. Lexa, J.M. Saveant, Acc. Chem. Res. 16 (1983) 235. [46] R.A. Marcus, J. Chem. Phys. 24 (1956) 966. [47] R.A. Marcus, Electrochim. Acta 13 (1968) 995. [48] M.E. Helton, N.E. Gruhn, R.L. McNaughton, M.L. Kirk, Inorg. Chem. 39 (2000) 2273.