Oxidation of Dihydrogen by Iridium Complexes of Redox-Active

Mar 24, 2010 - Preparation of Ir(III), Rh(III), and Ru(III) Amidophenolates. The complexes ..... Simulations employed DigiSim 3.05 software (BAS, West...
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1956

Organometallics 2010, 29, 1956–1965 DOI: 10.1021/om9010593

Oxidation of Dihydrogen by Iridium Complexes of Redox-Active Ligands Mark R. Ringenberg, Mark J. Nilges,‡ Thomas B. Rauchfuss,* and Scott R. Wilson School of Chemical Sciences, University of Illinois at Urbana-Champaign, Urbana, Illinois 61801. ‡ The Illinois EPR Research Center, School of Molecular & Cell Biology Received December 10, 2009

Unsaturated organoiridium complexes were prepared with amidophenolate ligands derived from 2-(2-trifluoromethyl)amino-4,6-di-tert-butylphenol (H2tBAF) and 2-tert-butylamino-4,6-di-tert-butylphenol (H2tBAtBu). The following 16e complexes were characterized: Cp*M(tBAR) with M = Ir (1F and 1t-Bu), Rh (2F), and (cymene)Ru(tBAF) (3F). These complexes undergo two 1e oxidations at potentials of about 0 and -0.25 V vs Cp2Fe0/þ. The magnitude of ΔE1/2 is sensitive to the counteranions, and the reversibility is strongly affected by the presence of Lewis bases, which stabilize the oxidized derivatives. Crystallographic measurements indicate that upon oxidation the amidophenolate ligands adopt quinoid character, as indicated by increased alternation of the C-C bond lengths in the phenylene ring backbone and shortened C-N and C-O bonds. Unlike the charge-neutral precursors, the cationic [Cp*M(tBAR)]þ are Lewis acidic and form well-characterized adducts with PR3 (R = Me, Ph), CN-, MeCN (reversibly), and CO. In the absence of competing ligands, the cations oxidize H2. Coulommetry measurements indicate that H2 is oxidized by the monocations [Cp*M(tBAR)]þ, not the corresponding dications. Oxidation of H2 is catalytic in the presence of a noncoordinating base at potentials required for the generation of [Cp*M(tBAR)]þ. The rate decreases in the order [Cp*M(tBAF)]BArF4 > [Cp*M(tBAF)]PF6 > [Cp*M(tBAt-Bu)]PF6. The reduction of ferrocenium by H2 is catalyzed by Cp*M(tBAR).

Introduction Complexes of redox-active ligands have long attracted the attention of inorganic chemists.1-6 Whereas metal-centered redox dramatically alters the bonding of the metal-ligand ensemble, changes in the redox state of the ligands have subtler affects on the metal center.7-9 Redox-active ligands may lead to faster catalysts for the reaction of small inorganic molecules (e.g., H2, N2, O2) because reorganizational barriers for redox-active organic substituents are typically lower than for inorganic centers. Furthermore, redox-active ligands supplement the electrons transferred from the metal, which can facilitate reactions that require multielectron transfers. Redox-active ligands are pervasive in enzymes, *Corresponding author. E-mail: [email protected]. (1) Allgeier, A. M.; Mirkin, C. A. Angew. Chem., Int. Ed. 1998, 37, 895–908. (2) Butin, K. P.; Beloglazkina, E. K.; Zyk, N. V. Russ. Chem. Rev. 2005, 74, 531–553. (3) De Bruin, B.; Hetterscheid, D. G. H.; Koekkoek, A. J. J.; Gruetzmacher, H. Prog. Inorg. Chem. 2007, 55, 247–354. (4) Kaim, W.; Schwederski, B. Pure Appl. Chem. 2004, 76, 351–364. (5) Chakraborty, S.; Laye, R. H.; Paul, R. L.; Gonnade, R. G.; Puranik, V. G.; Ward, M. D.; Lahiri, G. K. J. Chem. Soc., Dalton Trans. 2002, 1172–1179. (6) Fenske, D. Ber. 1979, 112, 363–375. (7) Klein, R. A.; Elsevier, C.; Hartl, F. Organometallics 1997, 16, 1284–1291. (8) Hartl, F.; Stufkens, D. J.; Vlcek, A. Inorg. Chem. 1992, 31, 1687– 1695. (9) Lorkovic, I. M.; Wrighton, M. S.; Davis, W. M. J. Am. Chem. Soc. 1994, 116, 6220–6228. (10) Stubbe, J.; van der Donk, W. A. Chem. Rev. 1998, 98, 705–762. pubs.acs.org/Organometallics

Published on Web 03/24/2010

especially those associated with O2 processing.10,11 The advantages of redox-active ligands for the catalytic transformations of small molecules have been demonstrated.12-14 Many redox-active noninnocent ligands (NILs) are based on cis-disubstituted-1,2-alkene and 1,2-phenylene, i.e., X-C(R)dC(R)-X. These complexes include the dithiolenes (X = S), diimines (X = NR), and catecholates (X = O).15-21 In this work, we are interested in converting M-NIL ensembles into stable cationic Lewis acids. One problem we faced is that conventional NILs typically oxidize to cations at very high potentials (most redox interconversions involve anions). Furthermore, complexes of many NILs tend to dimerize (11) Jazdzewski, B. A.; Tolman, W. B. Coord. Chem. Rev. 2000, 200-202, 633–685. (12) Knijnenburg, Q.; Gambarotta, S.; Budzelaar, P. H. M. Dalton Trans. 2006, 5442–5448. (13) Ben-Ari, E.; Leitus, G.; Shimon, L. J. W.; Milstein, D. J. Am. Chem. Soc. 2006, 128, 15390–15391. (14) Frech, C. M.; Ben-David, Y.; Weiner, L.; Milstein, D. J. Am. Chem. Soc. 2006, 128, 7128–7129. (15) Fourmigue, M. Coord. Chem. Rev. 1998, 178-180, 823–864. (16) Stiefel, E. I. Dithiolene Chemistry, Synthesis, Properties, and Applications; Prog. Inorg. Chem.; John Wiley: Hoboken, NJ, 2004. (17) Hendrickson, D. N.; Pierpont, C. G. Valence Tautomeric Transition Metal Complexes. In Spin Crossover in Transition Metal Compounds II; Springer: Berlin/Heidelberg, Germany, 2004; pp 786-786. (18) Evangelio, E.; Ruiz-Molina, D. Eur. J. Inorg. Chem. 2005, 2957– 2971. (19) Poddel’sky, A. I.; Cherkasov, V. K.; Abakumov, G. A. Coord. Chem. Rev. 2009, 253, 291–324. (20) Pierpont, C. G.; Lange, C. W. Prog. Inorg. Chem. 1993, 41, 331-442. (21) Moussa, J.; Amouri, H. Angew. Chem., Int. Ed. 2008, 47, 1372–1380. r 2010 American Chemical Society

Article Scheme 1. Proton-Induced Lewis Acidity versus OxidationInduced Lewis Acidity of a 16e Metal Complex

upon oxidation to neutral or cationic states, quenching any coordinative unsaturation.22 The N-substituted 2-amidophenolates popularized by Wieghardt et al. alleviate the problems of harsh redox potentials and dimerization.23-26 Such complexes undergo redox at potentials around -0.3 V vs Cp2Fe0/þ. Furthermore the amidophenolates oxidize to give complexes that remain monomeric by virtue of their bulky substituents. The amidophenolates are easily prepared and derivatized. For example, catechol is efficiently converted to a 3,5-di-tert-butyl derivative,28 and the resulting disubstituted diol then reacts with primary amines to give a range of aminophenols.29 We have demonstrated how ligand-centered oxidation of a Pt-amidophenolate-olefin complex induces nucleophilic attack at the olefin.27 Given their attractive properties, the amidophenolates are promising candidates for reactions wherein ligandcentered oxidation enhances the Lewis acidity of an unsaturated metal center. Metal-centered oxidation is known to cause increases in coordination number at the metal. The subtle effects of ligand-centered oxidation are, however, more applicable to catalytic processes where dramatic changes at the metal can lead to large kinetic barriers. The concept of oxidation-induced Lewis acidity (OI-LA) is related to our previously studied protonation-induced Lewis acidity (PI-LA, see Scheme 1), which was demonstrated in the case of 16e- diamido iridium(III) complexes.30,31 Like protonation-induced Lewis acidity, oxidation-induced Lewis acidity is expected to be operative when the starting complex is formally coordinatively unsaturated.32 Since the π-donor orbitals on the amidophenolates are also (22) Guyon, F.; Lucas, D.; Jourdain, I. V.; Fourmigue, M.; Mugnier, Y.; Cattey, H. Organometallics 2001, 20, 2421–2424. (23) Ray, K.; Petrenko, T.; Wieghardt, K.; Neese, F. Dalton Trans. 2007, 1552–1566. (24) Kokatam, S.-L.; Chaudhuri, P.; Weyhermueller, T.; Wieghardt, K. Dalton Trans. 2007, 373–378. (25) Bill, E.; Bothe, E.; Chaudhuri, P.; Chlopek, K.; Herebian, D.; Kokatam, S.; Ray, K.; Weyhermueller, T.; Neese, F.; Wieghardt, K. Chem.;Eur. J. 2005, 11, 204–224. (26) Kokatam, S.; Weyhermueller, T.; Bothe, E.; Chaudhuri, P.; Wieghardt, K. Inorg. Chem. 2005, 44, 3709–3717. (27) Boyer, J. L.; Cundari, T. R.; DeYonker, N. J.; Rauchfuss, T. B.; Wilson, S. R. Inorg. Chem. 2009, 48, 638–645. (28) Khomenko, T.; Salomatina, O.; Kurbakova, S.; Il’ina, I.; Volcho, K.; Komarova, N.; Korchagina, D.; Salakhutdinov, N.; Tolstikov, A. Russ. J. Org. Chem. 2006, 42, 1653–1661. (29) Blackmore, K. J.; Ziller, J. W.; Heyduk, A. F. Inorg. Chem. 2005, 44, 5559–5561. (30) Heiden, Z. M.; Gorecki, B. J.; Rauchfuss, T. B. Organometallics 2008, 27, 1542–1549. (31) Heiden, Z. M.; Rauchfuss, T. B. J. Am. Chem. Soc. 2009, 131, 3593–3600. (32) Nomura, M.; Fujii, M.; Fukuda, K.; Sugiyama, T.; Yokoyama, Y.; Kajitani, M. J. Organomet. Chem. 2005, 690, 1627–1637.

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Scheme 2. Oxidation of H2 by a Metal Complex via a Heterolytic Process

the redox-active orbitals (HOMO),19,24,33 ligand-centered oxidation would be expected to enhance the Lewis acidity at the metal by simultaneously raising the formal positive charge on the complex and weakening the metal-ligand π-bonds. In this project we apply oxidation-induced Lewis acidity to the oxidation of dihydrogen. The oxidation of dihydrogen by soluble metal complexes has long been studied,34,35 and in recent years detailed mechanistic information has become available.36,37 Recognizing that H2 oxidation begins with the formation of dihydrogen complexes,38 we sought complexes that would give H2 adducts susceptible to deprotonation. Dihydrogen complexes generally resist oxidation, but the corresponding hydrides are redox-active.39-41 Redox-active complexes of H2/H- complexes are operative in the [NiFe]and [FeFe]-hydrogenases, and the mechanism in Scheme 2 shows the general pattern of their reactions.42-44 We have previously communicated that the reduction of Agþ by H2 is catalyzed by amidophenolates of Cp*Ir(III).45 In this paper, we demonstrate the scope of oxidation-induced Lewis acidity to other ligands, and we further probe the proposed pathway for H2 activation.

Results Preparation of Ir(III), Rh(III), and Ru(III) Amidophenolates. The complexes Cp*Ir(tBAt-Bu) (1t-Bu), Cp*Ir(tBAF) (1F), Cp*Rh(tBAF) (2F), Cp#Ir(tBAF) (Cp# = tetramethylcyclopentadienyl), and (cymene)Ru(tBAF) (3F) were prepared from the corresponding dichloride dimers and (33) Hartl, F.; Rosa, P.; Ricard, L.; Le Floch, P.; Zalis, S. Coord. Chem. Rev. 2007, 251, 557–576. (34) Chalk, A. J.; Halpern, J.; Harkness, A. C. J. Am. Chem. Soc. 1959, 81, 854–857. (35) Harrod, J. F.; Halpern, J. Can. J. Chem. 1959, 37, 1933–1935. (36) Ogo, S. Chem. Commun. 2009, 3317–3325. (37) Rakowski DuBois, M.; DuBois, D. L. Chem. Soc. Rev. 2009, 38, 62. (38) Kubas, G. J. Adv. Inorg. Chem. 2004, 56, 127–177. (39) Rocchini, E.; Rigo, P.; Mezzetti, A.; Stephan, T.; Morris, R. H.; Lough, A. J.; Forde, C. E.; Fong, T. P.; Drouin, S. D. J. Chem. Soc., Dalton Trans. 2000, 3591–3602. (40) Cheng, T.-Y.; Szalda, D. J.; Zhang, J.; Bullock, R. M. Inorg. Chem. 2006, 45, 4712–4720. (41) Bullock, R. M. Chem.;Eur. J. 2004, 10, 2366–2374. (42) Artero, V.; Fontecave, M. Coord. Chem. Rev. 2005, 249, 1518– 1535. (43) Fontecilla-Camps, J. C.; Volbeda, A.; Cavazza, C.; Nicolet, Y. Chem. Rev. 2007, 107, 4273–4303. (44) Gloaguen, F.; Rauchfuss, T. B. Chem. Soc. Rev. 2009, 38, 100– 108. (45) Ringenberg, M. R.; Kokatam, S. L.; Heiden, Z. M.; Rauchfuss, T. B. J. Am. Chem. Soc. 2008, 130, 788–789.

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Figure 1. Structures of 2F (left) and 3F (right). Thermal ellipsoids are shown at the 50% probability level. Hydrogen atoms have been omitted for clarity. Selected bond lengths are shown in Table 1. Scheme 3. Rotamers of (cymene)Ru(tBAF) (3F)

aminophenols, 3,5-(t-Bu)2C6H2-6-OH-1-(NRH) where R = C6H4-2-CF3 and t-Bu (eq 1).

The new complexes were obtained as deeply colored, airstable solids that are readily soluble in a range of organic solvents. The 1H NMR spectra of 1t-Bu, 1F, and 2F are straightforward; however, spectra for the corresponding Ru(II) compounds are more complex. For 3F, we observed two methyl doublets for the cymene ligand, indicating that the isopropyl methyl groups are diastereotopic. Upon warming the sample, these peaks broaden, but complete coalescence was not observed even at 85 C. This pattern indicates that rotation of the cymene ligand is hindered by the bulky 2-(trifluoromethyl)phenyl substituent on the amine (Scheme 3). Electrochemical Studies. The Rh and Ir amidophenolato complexes undergo two reversible one-electron oxidations (Figure 2), as analyzed by cyclic voltammetry in CH2Cl2 solutions. The first oxidation of the Ir complexes occurs about 0.090 V more positive than for the Rh analogues. The substituents on the nitrogen affect the redox potentials in the expected way. For 1t-Bu and 1F, E1/2(0/þ) = -0.015 and 0.05 V and E1/2(þ/2þ) = 0.160 and 0.355 V, vs Cp2Fe0/þ (Table 2). The potential and ΔE for the two redox steps depend strongly on the supporting electrolyte and solvent conditions.

Figure 2. Cyclic (solid line) and differential pulse (dashed line) voltammograms of 1t-Bu vs Cp2Fe0/þ. Conditions: CH2Cl2, 10-3 M 1t-Bu, 0.1 M Bu4NPF6, and scan rate = 0.100 V/s.

For CH2Cl2 solutions, the two one-electron oxidations of 1F are separated by 0.17 and 0.317 V when the electrolyte is Bu4NPF6 and Bu4NBArF4, respectively. As observed by Geiger and co-workers, the diminished value of ΔE for Bu4NPF6 reflects the stabilization of the higher oxidation state by ion pairing.46 From ΔEox (= Eox(1) - Eox(2)), one can calculate the disproportionation constant (Kdisp) for the monocations.47 For example, in the system [1F]nþ/PF6-, ΔEox = 0.172 V corresponds to Kdisp = 3.57  10-2 M. Thus, a solution of nominally 0.50 mM [1F]PF6 in CH2Cl2 will contain about 20% each (0.069 mM) of both 1F and [1F](PF6)2. For the [1F]nþ/BArF4- system, Kdisp is 50 times smaller at 6.99  10-4 M. Similarly, a nominally 0.50 mM solution of [1F]BArF4 will contain only ∼3% each (0.013 mM) of 1F and [1F](BArF4)2. The (arene)Ru complex 3F exhibits a third quasi-reversible oxidation at 0.570 V. We attribute this process to the RuII/III couple. As shown in this and related work,48 a single amidophenolate ligand supports only two ligand-based oxidations. [Cp*Ir(amidophenolate)]þ and Related Adducts. Oxidations of 1F and 1t-Bu were conducted on a preparative scale using AgPF6 and FcPF6, respectively. The BArF4- salts were obtained by extracting the cations into Et2O by salt exchange. The salts of [1F]PF6 and [1t-Bu]PF6 were obtained as analytically pure brown solids. The salt [1F]BArF4 was further characterized crystallographically. The Ir-N and Ir-O bonds (2.010(3) and 2.044(7) A˚, respectively) are elongated relative to Ir-N and Ir-O bonds (1.963(4) and 1.996(3) A˚, respectively) in 1F.45 The carbon-carbon distances in the phenylene backbone of the amidophenolate differ from those in the neutral 1F, which average 1.40 A˚ (Table 1). Thus, the C11-C12 and C14-C15 bonds are elongated by 0.05 A˚, and the C13-C14 and C15-C16 bonds are shortened by 0.04 A˚. The pattern of bond length alternation in the phenylene backbone is similar to that of other semiquinone complexes.49 (46) LeSuer, R. J.; Geiger, W. E. Angew. Chem., Int. Ed. 2000, 39, 248–250. (47) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; John Wiley & Sons Inc.: New York, 2001; p 856. (48) Fujihara, T.; Okamura, R.; Tanaka, K. Chem. Lett. 2005, 34, 1562–1563. (49) Pierpont, C. G.; Buchanan, R. M. Coord. Chem. Rev. 1981, 38, 45–87.

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Table 1. Selected Bond Lengths (A˚) of Amidophenolato Complexes

M1-O1 M1-N1 O1-C11 N1-C16 C11-C12 C12-C13 C13-C14 C14-C15 C15-C16 C16-C11

1F (M = Ir)45

2F (M = Rh)

3F (M = Ru)

[1F]BArF445

[1F(MeCN)](OTf)2

1.963(4) 1.996(3) 1.342(4) 1.393(5) 1.416(5) 1.396(6) 1.396(6) 1.388(5) 1.395(6) 1.407(5)

1.958(2) 2.0032(18) 1.332(3) 1.390(3) 1.419(4) 1.390(4) 1.403(4) 1.378(4) 1.398(4) 1.405(4)

1.9600(18) 2.0032(15) 1.338(3) 1.396(3) 1.415(3) 1.391(3) 1.414(3) 1.379(3) 1.399(3) 1.403(3)

2.010(3) 2.044(7) 1.287(5) 1.342(7) 1.451(9) 1.398(1) 1.363(5) 1.453(6) 1.370(8) 1.427(8)

2.103(5) 2.095(6) 1.250(8) 1.314(9) 1.445(10) 1.344(10) 1.457(10) 1.356(10) 1.434(10) 1.480(10)

Table 2. Redox Couples and Disproportionation Constants (M) for Selected Complexes F

1 , Bu4NPF6 2F, Bu4NPF6 3F, Bu4NPF6 1F, Bu4NBArF4 1t-Bu, Bu4NPF6 1t-Bu, Bu4NBArF4

E1/20/þ (V)a

E1/2þ/2þ (V)a

ΔEo

Kdisp

0.073 -0.040 -0.015 0.005 -0.015 -0.262

0.245 0.227 0.110 0.380 0.160 -0.060

0.172 0.267 0.125 0.375 0.175 0.202

3.57  10-2 5.67  10-3 8.88  10-2 6.99  10-4 3.36  10-2 2.02  10-2

a Potentials vs Cp2Fe0/þ; conditions: CH2Cl2 solution, 0.1  10-3 M in complex, 0.1 M in Bu4NPF6.

Although the neutral amidophenolate complexes show no tendency to bind ligands, the cations are Lewis acidic. Thus oxidation of 1F in MeCN solution gives the adduct [1F(NCMe)]PF6. Solid samples of the adduct [1F(NCMe)]PF6 were found to liberate MeCN, reverting to the naked salt [1F]PF6 upon washing with ether and heating under vacuum. Stable adducts arise from treatment of cations [1t-Bu]þ and [1F]þ with CO, phosphines, and cyanide. A deep brown solution of [1t-Bu]PF6 in CH2Cl2 rapidly forms a pale orange adduct with CO, [1t-Bu(CO)]PF6 (νCO = 2049 cm-1). In contrast to the MeCN adduct, the CO complex showed no tendency to dissociate under vacuum. The salts [1t-Bu]þ and [1F]þ were found to react with the phosphines PMe3 and PPh3 to give the adducts [1R(PR3)]PF6, which were characterized by EPR spectroscopy and electrochemically (see below). The cherry-red charge-neutral cyanide 1t-Bu(CN) (νCN(KBr) = 2119 cm-1) was obtained by treating [1t-Bu]PF6 with NaCN. The complex was also prepared by treating 1t-Bu with AgCN. Isotropic liquid EPR spectra for the naked complex [1t-Bu]PF6 are characterized by giso = 1.9718, close to the g value for an organic radical (2.0023). As a frozen glass in 1:3 CH2Cl2/toluene, [1t-Bu]PF6 exhibits an axial spectrum g|| = 2.001 and g^ = 1.948 with 14N coupling A|| = 23.5 and A^ = 63.6 MHz. Spectra for the adducts, [1t-Bu(PMe3)]PF6 and 1t-BuCN, suggest greater delocalization of the electron onto the metal: 1t-Bu(CN) giso = 1.9741, A(14N) = 31.6, 30.3, and A(1H) = 19.1 MHz. The isotropic spectrum of [1t-Bu(PMe3)]PF6 has giso = 1.972 and A(14N) = 53.9, A(1H) = 24.0, and A(31P) = 52.8 MHz. Examples of Ir(IV) phosphine complexes have been characterized previously including

[Cp*IrMe2(PMe3)]þ, which exhibits a rhombic spectrum centered at g = 2.183 (A(31P) was not resolved at -100 C in frozen solution).50 The magnetic moments of 1t-Bu(CN), [1t-Bu]BArF4, F [1 ]PF6,45 and [1t-Bu(PPh3)]PF6 were found to be 1.78, 1.59, 1.75, and 1.55 μB, respectively. These values, which were measured on solutions, are reasonable for typical Ir(IV) complexes51 or organic radicals. The magnetic susceptibility of [1t-Bu(PPh3)]PF6 obeys the Curie-Weiss dependence over the temperature range -15 to 40 C. The cyclic voltammograms of the amidophenolate complexes are strongly affected by donor solvents. Thus, the first oxidation of 1F in MeCN solution is quasi-reversible even at scan rates greater than 2 V/s. The ic for the [1F]þ/0 couple shifts cathodically and becomes smaller as the scan rate is increased to 5 V/s. The cathodic shift and lower ic are consistent with an EC process, whereby oxidation of 1F is followed by rapid coordination of solvent to form [1F(MeCN)]þ. The second oxidation is reversible, consistent with a couple where the coordination number at iridium remains unchanged. The voltammograms for the oxidation of 1F and 1t-Bu in MeCN were digitally simulated, being sensitive to the parameters, E1/2 for the [1F]þ/0 and [1F(NCMe)]þ/2þ couples, the rate of the reaction MeCN þ [1F]þ, and the binding constant of MeCN to [1F]þ (Supporting Information). The voltammogram of 1F and 1t-Bu with the noncoordinating solvent and electrolyte CH2Cl2/Bu4NBArF was shown to be very informative in developing the concept of OI-LA. The redox properties of 1F are strongly affected by the addition of PPh3, consistent with the formation of the adduct, [1F(PPh3)]þ. The subsequent addition of PPh3 causes the voltammogram shown in Figure 3; the presence of a common intersection of the i-V traces for several concentrations of PPh3 (isoalpha points52) indicates that the reaction [1F]þ þ PPh3 gives a single product. The CV measurements with variable amounts of PPh3 indicate that PPh3 (50) Diversi, P.; Iacoponi, S.; Ingrosso, G.; Laschi, F.; Lucherini, A.; Zanello, P. J. Chem. Soc., Dalton Trans. 1993, 351–352. (51) Cotton, S. A. Chemistry of Precious Metals; Chapman & Hall: London, 1997. (52) Sanecki, P. T.; Skital, P. M. Electrochim. Acta 2008, 53, 7711– 7719.

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Figure 3. Cyclic voltammogram of 1F þ 0, 0.5, and 1 equiv of PPh3 vs Cp2Fe0/þ. Conditions: CH2Cl2, 0.12  10-3 M 1F, 0.2 M Bu4NBArF4, and scan rate = 0.1 V s-1.

Figure 4. Cyclic voltammogram of a solution of 1F and PPh3 (black) vs Cp2Fe0/þ and simulation (dashed blue). Conditions: CH2Cl2, 0.12  10-3 M 1F, 0.24 mM PPh3, 0.2 M Bu4NBArF4, and scan rate = 0.1 V s-1.

forms a one-to-one adduct (Supporting Information). The height of the current response (ip) for observed redox couples has a linear dependence with respect to (scan rate)1/2 in the presence of 2 equiv of PPh3 below 1 V s-1. At scan rates above 1 V s-1 this dependence deviates from linearity, indicative of the electron transfer rates (i.e., [1F]0/þ) being faster than the binding of Lewis bases.53 The voltammogram of 1F with PPh3 was digitally simulated (Figure 4). The binding of PPh3 by [1F]þ occurs at ca. 105 s-1 M-1 with Kequlib = 2  108 M-1. Conversely, the simulation indicated that dissociation of PPh3 from 1F(PPh3) proceeds at a rate of 103 s-1 with K = 6.2  106 M. Cyclic voltammetry experiments on [1t-Bu]0/þ in the presence of PPh3 gave similar findings. The couples [1F(MeCN)]þ/2þ and [1F(PPh3)]þ/2þ differ by only 0.091 V. The insensitivity of the couple to the nature of the Lewis base suggests that the second oxidation is highly localized on the amidophenolate ligand. Doubly Oxidized Species [Cp*Ir(tBAR)L]2þ. Oxidation of F 1 with 2 equiv of AgOTf in MeCN solution gave the deep purple diamagnetic salt [1F(MeCN)](OTf)2, which was characterized by 1H and 19F NMR spectra. Crystallographic analysis of [1F(MeCN)](OTf)2 supports the assignment of the ligand to the iminoquinone oxidation state.24 Thus, the Ir1-N1 and Ir1-O1 distances of about 2.10 A˚ are ∼0.11 A˚ longer than in 1F (Table 1). The carbon-oxygen (O1-C11, 1.25 A˚) and carbon-nitrogen (N1-C12, 1.31 A˚) distances (53) Creutz, C. Prog. Inorg. Chem. 1983, 1–73.

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Figure 5. Structure of the cation in [1F(MeCN)](OTf)2 with thermal ellipsoids at 50% probability and hydrogen atoms omitted. Selected bond lengths are shown in Table 1.

correspond to double bonds, indicative of the description of this ligand as an iminoquinone.24 Recrystallization of [1F(MeCN)](OTf)2 from CH2Cl2/ ether as well as heating under vacuum did not liberate the MeCN, in contrast to the behavior of the monocationic derivatives [1F(MeCN)]þ. A MeCN solution of [1F(MeCN)](OTf)2 was found to convert to [Cp*Ir(MeCN)3]2þ over the course of hours at room temperature, indicating weak binding properties of the doubly oxidized ligand. In CH2Cl2 solution, 1t-Bu(CN) was found to reversibly oxidize at -0.28 V, which is 0.42 V more negative than the [1t-Bu]þ/2þ couple. Oxidation of 1t-Bu(CN) with [Cp2Fe]PF6 gave diamagnetic [1t-Bu(CN)]PF6. Catalytic Oxidation of H2. Solutions of the neutral complexes (1F, 1t-Bu, 2F, 3F) in CH2Cl2 solution are unreactive toward 1 atm of H2 even for extended reaction times. A solution of [1t-Bu(MeCN)]PF6 in CH2Cl2 was very slow to react with H2. Under strictly analogous conditions, [1t-Bu]PF6 reacts with H2 over the course of several minutes at room temperature, as indicated by a color change from deep red to pale yellow. The reaction affords [Cp*Ir(μ-H)3IrCp*]PF654 and free aminophenol, consistent with the stoichiometry in eq 2.

2½1t-Bu þ þ 4H2 f ½CpIrðμ-HÞ3 IrCpþ þ 2HOðt-BuðHÞNÞC6 H2 ðt-BuÞ2 þ Hþ

ð2Þ

The reaction was monitored by UV-vis spectroscopy (λmax = 460 nm for [1t-Bu]PF6), and the rate was found to be first order in both [1t-Bu]PF6 and PH2 with d[1t-Bu]/dt = 1.61  10-4 ((0.12  10-4) s-1 at 22 C with PH2 of 1 atm. The oxidation of H2 also proceeded well in the presence of a noncoordinating base; however hydrogenolysis of the complex was suppressed. Thus, reduction of [1t-Bu]PF6 in CH2Cl2 containing 5 equiv of 2,6-di-tert-butylpyridine (t-Bu2py) and 1 atm of H2 resulted in the clean formation of neutral 1t-Bu (eq 3).

2½1t-Bu þ þ H2 þ 2t-Bu2 py f 21t-Bu þ 2t-Bu2 pyHþ ð3Þ The rate of the hydrogenation reaction in the presence of t-Bu2py was evaluated by monitoring the disappearance of (54) Fettinger, J. C.; Pleune, B.; Poli, R. J. Am. Chem. Soc. 1996, 118, 4906–4907.

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Table 3. Rate Constants for Reduction of [(C5Me5-nHn)M(O(NR)C6H2(t-Bu)2)]þ by H2 (5 equiv t-Bu2Py, CH2Cl2 soln, 22 C unless otherwise noted) catalysta

rate, at 1 atm H2, s-1

[1t-Bu]PF6 [1F]PF6 [1F]BArF4 [Cp#Ir(tBAF)]PF6 [2F]BArF4 [3F]PF6 [Ni(PCy2NBz2)2](BF4)2 AgOAc in pyridine (70 C)

1.01  10-3 ((2.66  10-5) 1.85  10-3 ((1.06  10-4) 4.66  10-2 ((2.28  10-3) 1.13  10-3 ((8.69  10-4) 4.74  10-3 ((2.18  10-4) no reaction 10 (ref 55) 0.49 (ref 34)

a

Cp# = C5Me4H-. PCy2NBz2 = [C6H11PCH2N(CH2Ph)CH2]2.

[1t-Bu]PF6. The rate was unaffected by the concentration of base and exhibited a first-order dependence on both [1t-Bu]PF6 and pH2 (eq 4). The rate of reduction of [1t-Bu]PF6 by 1 atm of D2 in the presence of t-Bu2py gave kH/kD = 2.01.

d½1t-Bu =dt ¼ ð1:0110 -3 ÞðPH2 Þ½ð1t-Bu Þþ  ð22 CÞ

ð4Þ

Similarly, the corresponding reduction of [1F]PF6 by H2 with t-Bu2py is nearly twice as fast as [1t-Bu]PF6 with a kH/kD = 1.2. Changing from Cp* to C5Me4H only slightly affected the rate. The rate of reduction was strongly affected by the counteranion: reduction of [1F]BArF4 by H2 is about 30 times faster than [1F]PF6. Different metal cations were also evaluated: the analogous [2F]BArF4 was reduced by H2 but was 10 times slower than the iridium complex, and the cymene ruthenium complex, [3F]PF6, does not react with H2. A summary of the observed first-order rate constants at 1.0 atm of H2 is shown in Table 3. With 1t-Bu as a catalyst, [Cp2Fe]þ can be reduced with H2 (Figure 6). Thus, a blue-green H2-saturated solution consisting of equal parts [Cp2Fe]PF6 and t-Bu2py became pale yellow several hours after the addition of 2.5 mol % of 1t-Bu. The rate (d[Cp2Fe]/dt) was shown to be independent of the initial concentration of [Cp2Fe]þ. The overall rate constant for the reduction of Fcþ is 9.00  10-4 s-1, which is comparable to 1.01  10-3 s-1 for the reduction of [1t-Bu]PF6 in the absence of the [Cp2Fe]þ.

Discussion The 2-amidophenolate ligands are strong π-donors and as such stabilize metal complexes that are formally electronically unsaturated. The oxidation-induced Lewis acidity in such complexes has not been widely exploited. In this work we show that oxidation of 16e amidophenolate derivatives of Cp*M2þ (M = Rh, Ir) enhances their electrophilicity sufficiently to cause an interaction with dihydrogen. The spectroscopic, crystallographic, and electrochemical measurements indicate that the oxidation is significantly localized on the amidophenolate ligand. Since the complexes undergo two oxidations, we conducted experiments aimed at determining which of the two cationic states is responsible for the oxidation of H2. The bimolecular kinetics are compatible with either (or both) the monocation and a dication being the active oxidant. (55) Wilson, A. D.; Newell, R. H.; McNevin, M. J.; Muckerman, J. T.; Rakowski DuBois, M.; DuBois, D. L. J. Am. Chem. Soc. 2006, 128, 358–366. (56) Kubas, G. Metal Dihydrogen and σ-Bond Complexes; Kluwer Academic/Plenum: New York, 2001.

Figure 6. Solution of 0.1 M [Cp2Fe]PF6, 0.4 M t-Bu2py, and 1t-Bu (catalyst) in 10 mL of CH2Cl2 before (left) and 18 h after pressurization with 3 atm of H2 (right).

Although the “hydrogenophilicity” of metal centers normally increases with charge,56 our data indicate that the singly oxidized species are better catalysts than the dication. The role of the singly oxidized species is supported by the finding that [1F]BAr4F is nearly 30 times faster than the PF6- salt. Being subject to a far higher comproportionation constant, solutions of the BArF4- salt contains less of the dication (and neutral) components. Coulommetry experiments are also consistent with the activation of H2 by the monocations, since the rate (current) of H2 oxidation is not enhanced at potentials above the second oxidation. The inactivity of the dications is attributed in part to their high Lewis acidity, which may result in their binding counteranions, which compete with the binding of H2. Matching a complex’s Lewis acidity and hydrogenophilicity is a key consideration in the design of catalysts for oxidation of H2.37 In terms of mechanism, the proposed catalytic cycle involves the intact complex, consistent with the fact that the reduction of [1t-Bu]þ by H2 occurs far faster than the H2induced hydrogenolysis to give free aminophenol. Substituents on the amine affect the rate: the trifluoromethylphenyl derivative [1F]þ is reduced by H2 nearly twice as fast as the tert-butyl derivative [1t-Bu]þ. Substituents on the cyclopentadienyl ligands (C5Me5 vs C5Me4H) and the identity of the metals (Rh vs Ir) also affect the rates. On the basis of the above analysis, we propose that H2 oxidation begins with formation of the complex [Cp*Ir(tBAR)(H2)]þ. This assumption is consistent with the inhibition of H2 oxidation by MeCN. In the next step, the H2 adduct undergoes deprotonation, consistent with the high Brønsted acidity of many cationic dihydrogen complexes.56,57 Deprotonation of this dihydrogen complex would give the hydride Cp*Ir(tBAR)H, which would be readily oxidized by either [Cp*Ir(tBAR)]þ or Fcþ. In terms of their redox potential, the (unobserved) hydride can be modeled to some extent by the cyanide 1t-BuCN, although the hydride is expected to be the better electron donor of the two.58 By comparing the couples for (57) Zhang, K.; Gonzalez, A. A.; Hoff, C. D. J. Am. Chem. Soc. 1989, 111, 3627–3632. (58) Lever, A. B. P. Inorg. Chem. 2002, 29, 1271–1285.

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Scheme 4. Proposed Pathway for the Oxidation of H2 by Iridium Amidophenolate Complexes

[1t-Bu]þ/2þ and [1t-BuCN]0/þ, we see that cyanide stabilizes the doubly oxidized state by 0.44 V. In fact the [1t-BuCN]0/þ couple is 0.260 V more negative than the [1t-Bu]0/þ couple; that is, 1t-BuCN would be readily oxidized by [1t-Bu]þ. Thus, it is apparent that [1t-BuH] would readily convert to [1t-BuH]þ in the presence of [1t-Bu]þ (eq 5).

½1t-Bu X þ ½1t-Bu þ f ½1t-Bu Xþ þ ½1t-Bu 

ð5Þ

ðX - ¼ H - , CN - ; for X ¼ CN, E ¼ 0:260 VÞ Oxidation of metal hydrides is known to enhance their acidities by as much as 10-20 pKa units.59 Although in the present case the oxidation is ligand-centered, the acidity of the resulting [1t-BuH]þ is apparently readily deprotonated by t-Bu2Py, which has a pKa in MeCN of 11.5.60 The overall proposed mechanism is shown in Scheme 4.

Experimental Section Unless otherwise indicated, reactions were conducted using Schlenk techniques, and all reagents were purchased from Sigma-Aldrich. Reagents and routine solvents were obtained commercially and were purified by standard methods.61 Unless otherwise indicated, all solvents were HPLC grade and purified using an alumina filtration system (Glass Contour, Irvine, CA). Reagents. [Cp2Fe]PF6 was recrystallized by extraction into acetone followed by precipitation with ethanol.62 Electrolyte, Bu4NPF6, was crystallized prior to use by extracting into a minimal amount of acetone followed by addition of an equal part of ethanol and cooling to -30 C. CD2Cl2 and CD3CN (Cambridge Isotope Laboratories) were dried with CaH2 and distilled under vacuum onto NMR samples. 3,5-Di-tert-butyl-2hydroxy-1-(2-tert-butlylamine)benzene (H2tBAt-Bu) was prepared according to Blackmore et al. (Supporting Information).29 Celite 545 was dried at 120 C prior to use. We previously described the syntheses of 1F and [1F]PF6.45 (59) Ryan, O. B.; Tilset, M.; Parker, V. D. J. Am. Chem. Soc. 1990, 112, 2618–2626. (60) Izutsu, K. Acid-Base Dissociation Constants in Dipolar Aprotic Solvents; Blackwell Scientific Publications: Oxford, 1990; Vol. 35, p 166. (61) Perrin, D. D.; Armarego, W. L. F. Purification of Laboratory Chemicals, 3rd ed.; Pergamon Press: New York, 1988. (62) Connelly, N. G.; Geiger, W. E. Chem. Rev. 1996, 96, 877–910.

Instrumentation. NMR spectra were recorded on Varian UNITY INOVA 500NB or UNITY 500 spectrometers. FT-IR spectra were recorded on a Mattson Infinity Gold FTIR spectrometer. UV-vis spectra were recorded on a Cary 50 UV/vis spectrometer. Electrochemical experiments were performed using a BAS CV-50W voltammetric analyzer. For a typical cyclic voltammetry (CV) experiment the cell consisted of a 10-3 M analyte solution containing 0.1 M of supporting electrolyte. The working electrode was a 1 mM glassy carbon tipped electrode, which was polished between experiments using 1.0 μM alumina polish (Buehler) in deionized water. The reference electrode was Ag/ AgCl in saturated aqueous KCl with a Vycor tip, and a Pt wire was used as the counter electrode. Simulations employed DigiSim 3.05 software (BAS, West Lafayette, IN). Magnetic moments were recorded by the Evans method by comparing the shifts of the residual 1H benzene signals (δ 7.37) for solutions of C6D6 or CDCl3 with and without the paramagnetic sample.63 Using the same method, we found that the moment for [Cp2Fe]PF6 was 2.49 μB, vs the literature value of 2.62 μB.64 X-band EPR spectra were collected on a Varian E-122 spectrometer. Variable-temperature spectra were recorded using an E-257 variable-temperature accessory using liquid nitrogen as a coolant. The magnetic fields were calibrated with a Varian NMR Gauss meter, and the microwave frequency was measured with an EIP frequency meter. EPR spectra were simulated with the program SIMPOW6.65 The sample was flamesealed under vacuum in a 3  4 mm quartz tube at liquid nitrogen temperatures. [Cp*Ir(tBAF)(NCMe)]PF6, [1F(NCMe)]PF6. This salt was prepared analogously to [1F]PF6 but in MeCN solution and can be isolated by crystallization from Et2O. The MeCN was removed by double recrystallization from CH2Cl2 by the addition of ether. The solid was heated to 40 C at reduced pressure. Anal. Calcd for C31H39F9IrNOP(CH3CN) (found): C, 45.20 (44.60); H, 4.82 (5.04); N, 3.19 (3.14). After washing Anal. Calcd for C31H39F9IrNOP (found): C, 44.54 (44.77); H, 4.67 (4.30); N, 1.67 (1.69). Cp*Ir(tBAt-Bu), 1t-Bu. A solution of 1 g (1.256 mmol) of [Cp*IrCl2]2 in 40 mL of CH2Cl2 was combined with a solution of 700 mg (2.512 mmol) of H2tBAt-Bu in 20 mL of CH2Cl2. After 5 min of stirring, 1 g (7.24 mmol) of K2CO3 and 10 mg (60.3 μmol) of Et4NCl were added. After about 1 h the solution color had turned from orange to red. The reaction solution was filtered through a course frit to remove K2CO3, and the solvent was removed under reduced pressure. The solid was extracted from the residue into 10 mL of pentane, and the pentane solution was passed through a plug of Celite to remove the Et4NCl. The resulting solution was cooled to -30 C for 12 h, and the resulting crystals were collected and washed with cold pentane. Yield: 1.127 g (75%). 1H NMR (CD2Cl2): δ 1.30 (s, 9H), 1.48 (s, 9H), 1.83 (s, 9H), 1.89 (s, 15H), 6.71 (d, 1H), 7.33 (d, 1H). UV-vis (CH2Cl2) λmax, nm (ε): 456 (9436). Anal. Calcd for C28H44IrNO (found): C, 55.78 (55.74); H, 7.36 (7.71); N, 2.32 (2.46). [Cp*Ir(tBAt-Bu)]PF6, [1t-Bu]PF6. A solution of 63 mg (0.249 mmol) of AgPF6 in 5 mL of CH2Cl2 was added to a solution of 150 mg (0.249 mmol) of 1t-Bu in 40 mL of CH2Cl2. Upon (63) Girolami, G. S.; Rauchfuss, T. B.; Angelici, R. J. Synthesis and Technique in Inorganic Chemistry, 3rd ed.; University of Science Books: Sausalito, CA, 1999; pp 47-53. (64) Gray, H. B.; Hendrickson, D. N.; Sohn, Y. S. Inorg. Chem. 2002, 10, 1559–1563. (65) EPR spectra were simulated with the automated fitting program, SIMPOW6, a program developed at the University of Illinois based on POW (Nilges, M. J. Ph.D. Thesis, University of Illinois, UrbanaChampaign, 1979) and MPOW (Chang, H.-R.; Diril, H.; Nilges, M. J.; Zhang, X.; Potenza, J. A.; Schugar, H. J.; Hendrickson, D. N.; Isied, S. S. J. Am. Chem. Soc. 1988, 110, 625-627).

Article addition of AgPF6, the reaction solution color changed from bright orange to dark brick red. The reaction solution was allowed to stir for 5 h. The reaction mixture was filtered through Celite to remove precipitated silver, rinsing with about 20 mL of CH2Cl2 until the effluent ran colorless. The solvent was removed under reduced pressure, and the resulting brick-red solid was crystallized from 5 mL of CH2Cl2 overlayered with 15 mL of hexane. To remove any excess 1t-Bu, the resulting crystals were washed with 10 mL portions of ether until the effluent ran colorless. The solid was dried at 40 C at 0.5 mmHg and stored in an N2-filled glovebox. Yield: 37 mg (20%). UV-vis (CH2Cl2, nm (ε)) λmax: 460 nm (4211). Anal. Calcd for C28H44F6IrNOP 3 CH2Cl2 (found): C, 41.82 (41.20); H, 5.46 (5.38); N, 1.68 (2.29). [Cp*Ir(tBAR)]BArF4, [1R]BArF4. The salt, [1R]BArF4, was made by ion exchange in which the [1R]PF6 was overlayered with ether containing 0.9 equiv of KBArF4. The reaction solution was vigorously stirred for ∼1 h, at which point the ether solution color turned dark brown. The residual PF6- salts were removed via filtration, and the filtrate was collected. The solvent was removed under reduced pressure to yield a brown powder. The magnetic moment of [1t-Bu]BArF4 was determined to be 1.59 μB, by the Evans method in C6D6 solution. [Cp*Ir(tBAt-Bu)(PPh3)]PF6, [1t-Bu(PPh3)]PF6. A solution of 110 mg (0.335 mmol) of Cp2FePF6 in 10 mL of CH2Cl2 was added to a solution of 205 mg (0.340 mmol) of 1t-Bu in 10 mL of CH2Cl2. Upon addition of the Cp2FePF6, the solution color turned from orange to brown. The reaction solution was stirred for 2 h, at which point 89 mg (0.340 mmol) of PPh3 was added, and the reaction solution was stirred for an additional 3 h. The solvent was removed under reduced pressure, and the resulting brown solid was washed with 2  10 mL of ether and 10 mL portions of n-hexanes until the effluent ran colorless. The resulting brown powder was crystallized from 5 mL of CH2Cl2 overlayered with 10 mL of ether. The solid was isolated via filtration and dried 20 C at reduced pressure. Yield: 140 mg (40%). Anal. Calcd for C46H59F6IrNOP2 3 CH2Cl2 (found): C, 51.55 (50.89); H, 5.61 (5.65); N, 1.27 (1.47). Magnetic moment (291 K): The magnetic moment of [1t-Bu(PPh3)]PF6 was determined to be 1.55 μB, by the Evans method in CDCl3 solution. [Cp*Ir(tBAt-Bu)(PMe3)]PF6, [1t-Bu(PMe3)]PF6. A Schlenk tube was charged with 23 mg (30.7 μmol) of [1t-Bu]PF6 and 5 mL of CH2Cl2. To a frozen solution excess PMe3 was transferred via vacuum distillation. The solution was thawed and shaken periodically over the course of 40 min. The color did not change substantially. Solvent and excess PMe3 were removed under reduced pressure, resulting in a brick-red, oily solid. The residue was triturated with 10 mL portions of n-hexanes until the effluent ran colorless. The resulting powder was isolated via filtration. UV-vis (CH2Cl2, nm) λmax: 460. [Cp*Ir(tBAt-Bu)(CO)]PF6, [1t-Bu(CO)]PF6. A solution of 22 mg (29 μmol) of [1t-Bu]PF6 in 5 mL of CH2Cl2 was sparged with CO for 30 s. The color of the solution lightened from brick red to yellow. The solvent was removed under reduced pressure, and the resulting solid was rinsed with 10 mL portions of hexanes until the extract was colorless. The resulting powder was isolated via filtration and dried under vacuum. IR (KBr): νCO = 2038 cm-1. Cp*Ir(tBAt-Bu)(CN), 1t-Bu(CN). A solution of 70 mg (0.116 mmol) of 1t-Bu in 10 mL of CH2Cl2 was added to a slurry of 15.6 mg (0.116 mmol) of AgCN in 10 mL of MeCN. No immediate color change was observed. After stirring for 12 h, the reaction solution had turned from orange to cherry red. The mixture was filtered through Celite to remove precipitated silver. The solvent volume was concentrated to about 5 mL, and 20 mL of hexanes was added. The resulting mixture was cooled to -30 C to yield a cherry-red solid. The air-stable product was isolated via filtration. Yield: 44 mg (60%). UV-vis (CH2Cl2) λmax, nm (ε): 480 (3887). MS-ESIþ (relative intensity): 630.2 m/z (100%, [Cp*Ir(tBAt-Bu)(CN)]þ). IR

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(KBr) νmax: 2119 cm-1 (CN). Anal. Calcd for C29H44IrN2O 3 CH2Cl2: C, 50.48 (49.62); H, 6.49 (6.50); N, 3.92 (4.01). Magnetic moment (291 K): The magnetic moment of 1t-BuCN was determined to be 1.78 μB, by the Evans method in C6D6 solution. [Cp*Ir(tBAt-Bu)CN]PF6, [1t-Bu(CN)]PF6. A solution of 24 mg (0.078 mmol) of [Cp2Fe]PF6 was added to a solution of 50 mg (0.079 mmol) of 1t-Bu(CN) in 10 mL of MeCN. The solution color turned from cherry red to brown. The reaction solution was stirred for 1 h, after which 30 mL of Et2O was added to precipitate a brown solid. The product was crystallized from 2 mL of CH2Cl2 and 10 mL of Et2O, followed by washing with 10 mL of hexanes. Yield: 47 mg (76%). 1H NMR (CD2Cl2): δ 1.01 (s, 9H), 1.35 (s, 9H), 1.48 (s, 9H), 1.58 (s, 15H), 6.15 (d, 1H), 7.35 (d, 1H). Anal. Calcd for C29H46IrN2OPF6 3 2(CH2Cl2): C, 39.45 (38.00); H, 5.12 (5.11); N, 2.96 (2.90). Cp*Rh(tBAF), 2F. A solution of 240 mg of H2tBAF (0.65 mmol) in 10 mL of MeCN was added to a slurry of 200 mg (0.32 mmol) of [Cp*RhCl2]2 in 15 mL of MeCN followed by 0.2 mL (14 mmol) of Et3N. The reaction solution was stirred for 4 h, and the solvent was removed under vacuum. The air-stable purple microcrystals were extracted into ca. 50 mL of n-hexane, and this extract was filtered to remove Et3NHCl. Evaporation of filtrate afforded a purple powder. Yield: 0.18 g (92%). Slow evaporation of a MeCN solution of 2F gave purple X-ray quality crystals. 1H NMR (CD2Cl2): δ 1.09 (s, 9H), 1.53 (s, 9H), 1.67 (s, 15H), 6.16 (d, 1H), 6.86 (d, 1H), 7.24 (d, 1H), 7.38 (t, 1H), 7.65 (t, 1H), 7.78 (d, 1H). 19F NMR (CD2Cl2): δ -60.5 (s). FD-MSþ: m/z = 601. UV-vis (CH2Cl2) λmax, nm (ε): 579 (3110), 353 (410), 288 (1300). Anal. Calcd for C31H39NF3ORh (found): C, 61.95 (61.45); H, 6.54 (6.73); N, 2.33 (2.48). [Cp*Rh(tBAF)]BArF4, [2F]BArF4. A solution of 105 mg (0.17 mmol) of 2 in 10 mL of CH2Cl2 was added a solution of 170 mg (0.169 mmol) of [Cp2Fe]BArF4 in 10 mL of CH2Cl2. An immediate color change from purple to green was observed. After 5 h, the solvent was removed under vacuum. The oily green residue was crystallized by extracting into 7 mL of CH2Cl2 and layered with 30 mL of hexane. The resulting green solid was washed with 10 mL of hexane and dried at 0.5 mmHg. A MeCN solution of the [Cp*Rh(tBAF)]BArF4 upon cooling to -30 C afforded brown X-ray quality crystals of the adduct [Cp*Rh(tBAF)(NCMe)]BArF4. Yield: 0.196 g (80%). ESIMSþ, m/z: 601 ([Cp*Rh(tBAF)]þ). UV-vis (CH2Cl2) λmax, nm (ε): 653 (4000), 437 (7400). UV-vis (MeCN) λmax, nm (ε): 793 (800), 575 (1300), 447 (3700), 333 (sh). Anal. Calcd for C63H51NBF27ORh (found): C, 53.44 (53.90); H, 3.63 (3.25); N, 1.00 (1.38). (cymene)Ru(tBAF), 3F. To a solution of 0.24 g (0.65 mmol) of H2tBAF in 10 mL of MeCN was added a solution of 0.2 g (0.33 mmol) of [(cymene)RuCl2]2 in 15 mL of MeCN followed by 0.19 mL (13 mmol) of Et3N. The reaction solution was stirred for 4 h, and the solvent was evaporated under reduced pressure. The red residue was extracted into 100 mL of hexanes, and the solid Et3NHCl was removed via filtration. The solvent volume was reduced to about 20 mL and by using a slow stream of argon, and the remaining solvent was evaporated to yield X-ray quality crystals. Yield: 170 mg (87%). 1H NMR (CD2Cl2): δ 1.08 (9H, s), 1.36 (6H, dd), 1.53 (9H, s), 1.73 (3H, s), 2.76 (1H, m), 4.98 (1H, d), 5.13 (1H, d), 5.38 (1H, d), 5.45 (1H, d), 6.20 (1H, d), 6.50 (1H, d), 6.89 (2H, s), 7.03 (1H, s), 7.29 (2H, s), 7.34 (1H, d), 7.44 (1H, d), 7.54 (1H, t), 7.64 (1H, t), 7.78 (1H, d). 19F NMR (CD2Cl2): δ -59.36 (s). ESIþ: m/z = 599 ([C31H38RuNO]þ). Anal. Calcd for C31H38NORu (found): C, 68.73 (67.82); H, 7.07 (7.12); N, 2.58 (2.44). [(cymene)Ru(tBAF)]PF6, [3F]PF6. A solution of 120 mg (0.36 mmol) of [Cp2Fe]PF6 in 15 mL of CH2Cl2 was added to 195 mg (0.33 mmol) of 3F in 15 mL of CH2Cl2. The solution color changed from red to purple. The reaction solution was stirred for 3 h, and the solvent was evaporated under reduced pressure. The resulting solid was dissolved in 2 mL of CH2Cl2 and layered with 10 mL of hexanes to afford purple crystals. The solid was

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Scheme 5. Parameters Used in Simulation of Voltammograms of 1F and 1t-Bu in the Presence of PPh3 and in MeCNa

a Operating conditions for 1R þ MeCN (R = F, t-Bu): MeCN, 0.2 mM 1R, 0.20 M Bu4NPF6, and scan rate = 0.1 V/s. Operating conditions for 1F þ PPh3: CH2Cl2, 0.12 mM 1F, 0.2 M Bu4NBArF4, and scan rate = 0.1 V/s.

washed with 2  10 mL of hexanes. Yield: 170 mg (70%). ESIMS: m/z = 743 ([(cymene)Ru(tBAF)]þ). Anal. Calcd for C31H38F6NOPRu (found): C, 50.01 (49.88); H, 5.15 (5.21); N, 1.88 (2.16). Anal. Calcd for C31H39F9IrNOP(Ph3P) (found): C, 51.20 (48.67); H, 4.85 (4.51); N, 1.27 (1.35). [Cp*Ir(tBAF)(MeCN)](OTf)2, [1F(MeCN)](OTf)2. To a solution of 50 mg (0.072 mmol) of 1F in 10 mL of CH2Cl2 with 0.1 mL of MeCN was added 37 mg (0.144 mmol) of AgOTf in 5 mL of CH2Cl2. The solution color turned from orange to dark red upon addition of the AgOTf solution. The reaction solution was stirred at room temperature for 5 h. The solvent volume was reduced to about 1 mL and overlayered with 10 mL of Et2O to yield a brick-red solid. The solid was isolated via filtration and crystallized from 1 mL of CH2Cl2 and 7 mL of Et2O at -30 C to yield X-ray quality crystals. Yield: 31 mg (42%) 1H NMR (CD2Cl2): δ 1.147 (9H, s), 1.427 (9H, s), 1.588 (15H, s), 2.743 (3H, s), 6.034 (1H, s), 7.405 (1H, s), 7.472 (1H, d), 7.870 (1H, d), 8.022 (2H, t). 19F NMR (CD2Cl2): δ -59.5. The decomposition of [1F(MeCN)](OTf)2 was monitored by adding 0.02 mL of CD3CN to the CD2Cl2 solution, and the 19F NMR signal at δ -59.5 was observed to disappear after about 5 h. The 1H NMR spectrum was complicated. Oxidation of 1F and 1t-Bu in the Presence of MeCN and PPh3. Solution of 8.5 mg (0.0123 mmol) of 1F in 10 mL of CH2Cl2/ Bu4NPF6, 0.12 mM 1F (0.24 mM 1t-Bu) in 0.1 M Bu4NPF6/ MeCN, and 18.1 mM PPh3 in CH2Cl2/Bu4NPF6 were prepared and sparged with solvent-saturated N2. Before each electrochemical experiment, the working electrode was polished using alumina,66 and the reaction solution was briefly sparged with N2. The electrochemical experiment was performed by the addition of varying equivalents of the PPh3 solution at a scan rate of 0.1 V/s. The final scan was performed after the addition of ∼10 mg of Cp2Fe as an internal standard. The voltammogram of 1F in MeCN was simulated according to the mechanism and using the parameters in Scheme 5. With very similar parameters, we also simulated the CV for 1t-Bu in MeCN. The parameters E1MeCN-F and E1MeCN-Bu matched those observed directly for the couple [1F]0/þ and [1t-Bu]0/þ. The simulations indicate that affinities of both [1t-Bu]þ and [1F]þ for MeCN are low, consistent with the ease with which they can be desolvated. The t-Bu derivative’s affinity is 10 times higher of the two, which was unexpected. For these simulations involving MeCN, it was unnecessary to include E3 and the fit was insensitive to Keq 2. The simulation for the oxidation of 1F in the presence of PPh3 indicated a very high affinity (108-1) of [1F]þ for the phosphine. Because of this high affinity, the couple [1F(PPh3)]þ/0 was observable (E3PP3-F = -0.984 V) and incorporated into the (66) Tilset, M. Organometallic Electrochemistry: Thermodynamics of Metal-Ligand Bonding. In Comprehensive Organometallic Chemistry III; Crabtree, R. H., Mingos, D. M. P., Eds.; Elsevier: Oxford, 2007; pp 279-305.

simulation. The fit was quite sensitive to the rate of dissociation of PPh3 from [1F(PPh3)]þ and the associated equilibrium constant (Scheme 5). The heterogeneous electron transfer rate and symmetry parameter (R) were left at the default value of 1  105 cm s-1 and 0.5. For [1F]þ/0 and [1F]þ/2þ, the default values for the symmetry parameter and electron transfer rates are consistent with the observed ic/ia ≈ 1 for scan rates 0.025 to 5 V/s.67 The diffusion coefficients for 1R and [1RL]þ were adjusted to be proportional to the slopes of the (linear) plots of ip vs ν1/2, and electrochemically coupled species (e.g., 1F and [1F]þ) were assigned the same diffusion coefficients of 4.533  10-5 cm2 s-1.68 The diffusion coefficient of Cp2Fe in MeCN is 2.37  10-5 cm2 s-1.69 Reduction of [1t-Bu]PF6 with H2. A 50 mL stock solution of t-Bu [1 ]PF6 (8.03  10-5 M) in CH2Cl2 was prepared. To an ovendried UV-vis cell that was fitted with an airtight Teflon stopcock was added 4 mL of the stock solution under argon. An initial UV-vis spectrum was recorded. Taking care not to introduce air, the solution was sparged with H2 for 20 s, and the cell was sealed under an active purge of H2 to establish approximately 1 atm of H2. The solution was shaken and placed in the spectrometer, and the initial spectrum was obtained 48 s after H2 was introduced. Spectra were recorded at 10 s intervals for 40 min (296 K). The solution was not stirred for the duration of the experiment. The rate was monitored by following the peak at λmax = 460 nm. Data from t = 0 to 360 s were analyzed. The observed first-order rate constant, an average of four experiments, was 1.01  10-3 ( 2.66  10-5 s-1. Reduction of [1t-Bu]PF6 with H2 in the Presence of Base. A 10 mL portion of stock solution of [1t-Bu](PF6) (1.01 mM) and t-Bu2py (2.26 mM) in CH2Cl2 was prepared. To the UV-vis reaction cell was added 1 mL of the stock solution followed by 10 mL of CH2Cl2. The initial UV-vis spectrum was recorded, and the sample was frozen. The frozen solution was evacuated and then pressurized with 1.02 atm of H2. The sample was thawed by shaking the flask under running tepid tap water; the initial spectrum was recorded at time = 124 s. Time = 0 was assigned to the moment the sample was completely thawed. The solution color was pale brown upon insertion into the spectrometer. Spectra were recorded at 20 s intervals for 40 min (296 K). The sample was not stirred for the duration of the experiment. Rates were determined by monitoring the disappearance of the peak at 460 nm. The data used was from time = 124 to 900 s. The average of four runs has a rate of reduction of 1.702  10-4 ( 2.04  10-6 s-1. When a solution of [1t-Bu]PF6 in CH2Cl2 was treated with t-Bu2py, the optical spectrum remained unchanged. (67) Wang, R. L.; Tam, K. Y.; Compton, R. G. J. Electroanal. Chem. 1997, 434, 105–114. (68) Taylor, A. W.; Qiu, F.; Hu, J.; Licence, P.; Walsh, D. A. J. Phys. Chem. B 2008, 112, 13292–13299. (69) Jacob, S. R.; Hong, Q.; Coles, B. A.; Compton, R. G. J. Phys. Chem. B 1999, 103, 2963–2969.

Article NMR analysis indicated the formation of 1t-Bu, [Cp*Ir(μH)3IrCp*]þ, t-Bu2pyHþ, and H2tBAt-Bu. The reductions of [1F]PF6 and [2F]BArF4 were performed in the same manner. Reduction of Cp2Feþ Salts. A stock solution of 115 mg (0.347 mmol) of [Cp2Fe]PF6, 5.2 mg (8.67 μmol) 1t-Bu, and 76 μL (0.347 mmol) of t-Bu2py was prepared in 100 mL of CH2Cl2. To the UV-vis reaction cell was added 10 mL of the stock solution. An initial UV-vis spectrum was recorded, and the sample was frozen. The frozen solution was evacuated and then pressurized with 1.02 atm of H2. The sample was thawed by shaking the flask under running tepid tap water; the initial spectrum was recorded at time = 100 s. Time = 0 was assigned to the moment when the sample was completely thawed. The sample was blue upon insertion into the spectrometer. Spectra were recorded at 60 s intervals for 500 min (296 K). The sample was not stirred for the duration of the experiment. Rates were determined by monitoring the disappearance of the peak at 625 nm. In a control experiment, a solution equimolar in [Cp2Fe]PF6 and t-Bu2py under H2 was shown to be stable for days at room temperature. Reaction of 2F, 3F, and of [2F]þ, [3F]þ with H2. The following ligand hydrogenolysis experiments were conducted in CD2Cl2 solutions under ∼1 atm of H2 and analyzed by 19F NMR spectroscopy (H2tBAF -62 ppm). The yields for the reaction of [2F]0/þ with H2 were calculated from the ratio of anion 19F NMR signal (-59 ppm, BArF4-) to free ligand signal. Treatment of [3F]0/þ with H2 did not yield significant amounts of free ligand. Yield of H2tBAF (time): 2F 20% (7 days), [2F]BArF4 100% (∼2 h), 3F none (7 days), [3F]PF6 trace (7 days). Crystallography. The compound [1F(NCMe)](OTf)2 was crystallized from CH2Cl2 overlayered with hexanes. The structure was solved by direct methods. The proposed model includes 1.5 idealized hexane solvate molecules per unit cell disordered about the inversion center, one idealized CH2Cl2 solvate per unit cell sharing space with the disordered hexane, two triflate anions, one ordered and one disordered over two sites, and one host cation. Similar geometry was imposed on the two disordered anion sites using a standard deviation of 0.01 A˚. Similar, rigid-bond restraints (esd 0.01) were imposed on displacement parameters for overlapping disordered sites. Ordered methyl H atom positions, R-CH3, were optimized by rotation about R-C bonds with idealized C-H, R-H, and H-H distances. Remaining H atoms were included as riding idealized contributors. Methyl H atom U’s were assigned as 1.5Ueq of the carrier atom; remaining H atom U’s were assigned as 1.2Ueq.

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The compound [2F(MeCN)]BArF4 was crystallized from a saturated MeCN solution at -30 C. The structure was solved by direct methods. The crystal contained two MeCN solvates per formula unit; one of these was disordered about an inversion center. Atoms of the disordered solvent were refined as a rigid, idealized group, and their displacement parameters were restrained to be isotropic with an effective standard deviation of 0.01. The Cp* ligand was also found to have a 2-fold disorder; the atoms modeling the disorder were restrained to have the same anisotropic displacement parameters with an effective standard deviation of 0.01. For BArF4-, all CF3 substituents were found to be disordered over two sites. These F3C-C(aryl) groups were restrained to similar, ideal geometry (esd 0.01 A˚ for bond lengths and 0.02 for angles). H atoms were included as riding idealized contributors. Methyl H atom U’s were assigned as 1.5Ueq of the carrier atom; remaining H atom U’s were assigned as 1.2Ueq. The complex 2F was crystallized from a saturated MeCN solution at -30 C. The structure of 2F was solved by direct methods. Methyl H atom positions, R-CH3, were optimized by rotation about R-C bonds with idealized C-H, R-H, and H-H distances. Remaining H atoms were included as riding idealized contributors. U’s for -CH3 atoms were assigned as 1.5Ueq of the carrier atom. Remaining H atom U’s were assigned as 1.2 carrier Ueq. The compound 3F was crystallized from a saturated MeCN solution at -30 C. The structure of 3F was solved by direct methods. Methyl H atom positions, R-CH3, were optimized by rotation about R-C bonds with idealized C-H, R-H, and H-H distances. Remaining H atoms were included as riding idealized contributors. U’s for -CH3 atoms were assigned as 1.5Ueq of the carrier atom. Remaining H atom U’s were assigned as 1.2 carrier Ueq.

Acknowledgment. This research was sponsored by the U.S. Department of Energy. We thank Dr. Swarna Latha Kokatam for preliminary studies on the Ru complexes. Supporting Information Available: Kinetic plots, spectra, rate calculations. Crystallographic information files (cif) for [1F(MeCN)](OTf)2, 2F, [2F(MeCN)]BArF4, and 3F. This material is available free of charge via the Internet at http:// pubs.acs.org.