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J . Am. Chem. Soc. 1992, 114, 3184-3792

3184

Organometallic Thermodynamics. Redox Couples Involving Metal-Metal Bonds J. Richard Pugh and Thomas J. Meyer* Contribution from the Department of Chemistry, The University of North Carolina at Chapel Hill, Chapel Hill, North Carolina 27599-3290. Received June 10, 1991

Abstract: Formal potentials for couples involving the oxidation or reduction of Mn2(CO)

[Fe(~5-CsH,)(CO)2]2, or [Mo(~s-C,Hs)(CO)3]2 have been measured. For two-electron couples, such as 2[Mn(CO)S(NCCH3)]+/Mn2(CO)lo (Eo’ = -0.30 f 0.03 V vs SSCE (saturated sodium chloride calomel electrode) in CH,CN, p = 0.1) or Mn2(CO),,/2[Mn(CO),]- (Eo’ = 4 . 6 9 f 0.01 V), the values were obtained by an equilibration technique. For one-electron couples such as Mn(CO),/[Mn(CO),]( E l l 2= -0.08 f 0.08 V), they were measured by fast-scan microelectrode cyclic voltammetry. The free energy changes for the homolytic dissociation of Mn2(CO),o(28 4 kcal/mol), [Fe($-C5H5)(CO)2]2(25 f 5 kcal/mol), and [Mo($-C,H,)(CO),]~ (22 3 kcal/mol) in CH’CN ( p = 0.1) were calculated by combining the potentials for the one- and two-electron couples. (Ell2 When these quantities are included with the potentials for intermediate one-electron couples such as [Mn2(CO)lo]+/o = 1.50 V), it is possible to construct complete or nearly complete Latimer diagrams which interrelate the various oxidation states. An overall pattern of reactivity toward electron transfer emerges from these diagrams. Oxidation-reduction mechanisms ~, by electron transfer are inaccessible on reasonable time scales in involving dissociation, Mn2(CO),o ~ M I I ( C O )followed Mn2(C0)9 solution at room temperature but could be important at higher temperatures although CO loss, Mn2(CO) + CO, is competitive. Once formed in CH’CN, the monomers [Mo(~~-C,H,)(CO)~] and Mn(CO), are unstable with respect to disproportionation, for example, into [Mn(CO),(NCCH,)]+ and [Mn(CO),]-, while [Fe(?S-C,H,)(CO)2] is stable. Large overvoltages exist in the two-electron oxidations of the metal-metal bonds, e.g., for Mn2(CO)lo 2CH3CN 2Mn(CO),(NCCH3)+ + 2e-, the overvoltage is +1.80 V, and in their reductions, Mn2(CO)lo+ 2e- 2[Mn(CO),]-, at electrodes, or in solution. These overvoltages exist because the reactions occur by one-electron steps to or from the electrode or to or from a reagent in solution and occur through intermediates, which are thermodynamically unstable, e.g., [Fe(q5-C,H5)(CO)2]2+ or [Fe($-C,H,)(CO),];. A related situation exists in the oxidation of the adjacent anions, to the metal-metal bonds, e.g., Mn2(CO)lo+ 2e-, or in the reduction of the adjacent cations, 2[Mn(CO),(NCCH3)]+ + 2e-- MII~(CO),~, 2[Mn(CO),]since these reactions occur via the unstable intermediates Mn(CO),, Fe($-C,H,)(CO),, or MO(~~-C,H,)(CO),.

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Introduction The systematic development of organometallic chemistry has been inhibited, in part, by the absence of comprehensive thermodynamic data. Where such data are available,’ insight has been gained into chemical reactivity and they have aided in the design of catalytic cycles. These data are difficult to acquire for reactions which are kinetically irreversible. This situation occurs frequently in organometallic redox chemistry, for example, where typically there are compositional and structural changes that lead to high kinetic barriers and slow reactions. Examples are the loss of metal-metal bonding with the gain or loss of electrons and oxidative addition, eqs 1 and 2. Because these reactions are typically slow at electrodes, electrochemical techniques such as cyclic voltammetry cannot be applied to the determination of reduction potentials in a direct manner.

+

-

M I I ~ ( C O ) , ~ 2e-

+

Ir(CO)C13(PPh3)2 2 2

2[Mn(CO),]-

(1)

pentadienyl)molybdenum]were purchased from Aldrich and purified by recrystallization from cyclohexane/isopentane. The purified compounds were stored in a freezer at 0 “C. Solutions containing the anions [Mn(CO),]-, [Fe($-C,H,)(CO),]-, or [Mo(qs-C,H,)(CO),]- were prepared by controlled-potentialelectrclysis. They were prepared by electrochemical reduction of 0.1 M [N(?-C4H9)4](PF6)/CH3CN solutions containing Mn2(CO)lo, [Fe(q5CSHS)(C0)2]2r or [ M O ( ~ ~ - C ~ H ~ ) ( C The O )reductions ~ ] ~ ~ were performed in a three-compartment electrolysis cell in a drybox at the peak potentials for the tweelectron reductions of the dimer. After the current had decreased to the baseline, UV-vis and FTIR spectra were used to confirm that the reactions were complete. The oxidation products [Mn(Co),(NCCH,)](PF6), [Fe(q5-C5H5)(co)2(NccH3)](PF6),and [Mo(q5-C5H5)(CO),(NCCH,)](PF6) were prepared by literature procedures.’ The polypyridyl complexes of osmium and ruthenium and all of the bipyridinium ions were prepared as hexafluorophosphate salts by literature procedure^.^ Spectrograde acetonitrile was purchased from Burdick and Jackson and used as received. Tetra-n-butylammonium hexafluorophosphate

Ir(CO)Cl(PPh3)2+2C1- (2)

We have developed a redox equilibration technique for measuring formal potentials for two-electron couples involving the oxidation or reduction of Mn2(CO)lo,[Fe($-C,H,)(CO),],, or [Mo(t&H,)(CO),] 2. Fast-scan, microelectrode cyclic voltammetry has been used to measure potentials for one-electron couples involving the monomers Mn(CO),, Fe($-C5Hs)(C0),, or Mo(q5-C5H5)(CO)3.When the results of the two are combined, a powerful quantitative basis is made available for describing the properties of these compounds toward electron transfer and for calculating free energy changes for their homolytic dissociation.2

Experimental Section Materials. The compound dimanganese decacarbonyl was purchased from Aldrich and purified by sublimation. The compounds bis[di~arbonyl(~~-cyclopentadienyl)iron]and bis[tricarbonyl($-cyclo(1) (a) Connor, J. A. Top. Curr. Chem. 1977, 71, 71. (b) See the entire issue of Polyhedron 1988, 7, 1. (c) Simh, J. A. M.; Beauchamp, J. L. Chem.

Rev. 1990, 90, 629. (2) Pugh, J. R.; Meyer, T. J. J . Am. Chem. Sac. 1988, 110. 8245.

(3) (a) Drew, D.; Darensbourg, D. J.; Darensbourg, M. Y. Inorg. Chem. 1975.7, 1579. (b) Ferguson, J. A.; Meyer, T. J. Inorg. Chem. 1971,lO. 1025. (4) (a) Sullivan, B. P.; Lumpkin, R. S.; Meyer, T.J. Inorg. Chem. 1987, 26,1248. (b) Takel-Takvoryan,N. F.; Hemingway, R. W.; Bard, A. J. J . Am. Chem. Soc. 1973,95,6582.(c) Neveux, P. W. Ph.D. Dissertation, University

of North Carolina, Chapel Hill, NC, 1987. (d) QQ2+= 6,7,8,9-tetrahydrodipyrido[ 1,2-a:2’,1’-c][ 1,4]diazocinediium dication. TQ2+= 7,8-dihydr*6Hdiprido[ 1,2-a:2’,1’-c][ 1,4]diazepinediium dication. DQ2+= 6,7-dihydrodipyrido[ 1,2-a:2’,1’-c]pyrazinediium dication. Homer, R. F.; Tomlinson, T. E. J . Chem. SOC. 1960,2498. Homer, R. F.;Tomlinson, T.E. Nature 1959,184, 2012. Salmon, R. T.; Hawkridge, R. M. J . Electroanal. Chem. Interfacial Electrochem. 1980, 112, 253. (e) Michaelis, L.; Hill, E. S.J . Gen. Physiol. 1933, 16, 859. Michaelis, L.; Hill, E. S. J . Am. Chem. SOC.1933, 55, 1491. Michaelis, L. J . Prakt. Chem. 1959, 9, 164. Michaelis, L. Eiochem. Z.1932, 250, 564. (f) Rosenblum, M. Chemistry of the Iron Group Metallocenes: Ferrocene, Ruthenocene, and Osmocene, John Wiley and Sons: New York, 1965; Vol. 2. (g) Buckingham, D. A.; Dwyer, F. P.; Goodwin, H. A.; Sargeson, A. M. Aust. J . Chem. 1964, 17, 325. (h) Sullivan, B. P.; Salmon, D. J.; Meyer, T. J. Inorg. Chem. 1978, 17, 3334. (i) Kealy, T. J.; Pauson, P. L. Nature 1951, 168, 1039. Miller, S. A.; Tebboth, J. A.; Tremaine, J. F. J. Chem. Sac. 1952, 632. Suji, T.; Aoyaqui, S. J. J. Electroanal. Chem. Interfacial Electrochem. 1975,58, 401. Kew, G.; DeArmond, K.; Hanck, K. J. J. Phys. Chem. 1981, 79, 541. Vlcek, A. A. Coord. Chem. Rev. 1982,43, 39.

u)

0002-7863/92/ 1514-3784%03.00/0 0 1992 American Chemical Society

Redox Couples Involving Metal-Metal Bonds

J . Am. Chem. SOC.,Vol. 114, No. 10, 1992 3185

Table I. E, ,2 Values in 0.1 M TBAH/CH,CN vs SSCE“

couple

ref -1.47 4a -1.35 4b Ru(bpy)32+/+ RU(~,~’-(CONE~~)~-~PY)~~’/’-1.18 4~ OS(~,~’-(CONE~~)~-~~~)(PP~,),(CO)H+/O -1.11 4a 0~(bpym)(PPh~),(CO)H’/~ -1.04 4a R~(5,5’-(COOEt)2-bpy),~’/’ -1.01 4c -0.95 4a 0~(bpy)(PPh,)~(C0)H+/~ 0~(5,5’-(C00Et)~-bpy)(PPh~)~(CO)H+I~ -0.80 4a QQ2+/+ b -0.62 4d TQ2+/+C -0.55 4d pQ2+/+d -0.44 4e DQ2+/+e -0.35 4d Fe(q5-CSMeS)2t/o -0.12 4f -0.04 4g OS(bpy)2C12+/’ +0.31 4h R~(bpy)2C12+/” +0.37 4i F~(CP)~+/O +0.48 4f Fe(p-COOEtCp)2t/0 O~(bpy),’+/~+ +0.85 4j +0.90 4h R~(bpy)~(PPh,)Cl~+/+ +1.05 4j Fe(b~y),)+/~+ “The ligand abbreviations are bpy = 2,2’-bipyridine, bpym = 2,2‘bipyrimidine, bpz = 2,2’-bipyrazine, Cp = qS-CsHs. 1,l’-Butylene2,2’-bipyridinium ion. 1,I’-Propylene-2,2’-bipyridiniumion. 1,l’Dimethyl-4,4’-bipyridinium ion. e 1,l’-Ethylene-2,2’-bipyridiniumion. Ell2,V

OS(bpy) (PPh3)2(CO)H+/O

(TBAH) was prepared by the metathesis of tetra-n-butylammonium bromide (Aldrich) and hexafluorophosphodcacid (Aldrich) in water, and recrystallized from hot ethanol and dichloromethane/diethyl ether followed by drying in vacuo at 80 OC for 2 days. Instrumentation. A PAR 273 potentiostat and a Hewlett-Packard 7015B X-Y recorder were used for all conventional electrochemical measurements. All potentials are reported versus the saturated sodium chloride calomel electrode, SSCE (0.236 V vs NHE). For rapid-scan microelectrode cyclic voltammetry a PAR 175 universal programmer and a home-built current follower were used in conjunction with a two-electrode cell configuration. The data were recorded on a Tektronix 2230 100-MHz digital storage oscilloscope. Microelectrodes were prepared by literature techniquess The IO-gm-diameter platinum wire was purchased from Goodfellow Metals. Infrared spectra were obtained by using a Nicolet 20DX FTIR spectrometer with a 0.1” path length NaCl solution cell. A Hewlett-Packard 8451A diode array spectrophotometer and an air-tight 0.1-cm quartz cell were used for UV-vis measurements. Procedures. All electrochemicalmeasurements were performed inside a Vacuum Atmospheres inert atmosphere drybox. Solutions for conducting the redox equilibration experiments were also prepared inside the drybox. In the redox equilibration experiments, solutions containing the two reagents were mixed and the resulting solution was allowed to sit for several hours until equilibrium had been reached. The equilibrated solutions were added to IR or UV-vis cells, sealed with rubber septa, and removed from the drybox for measurement. The details of a typical redox equilibration experiment were as follows. Stock solutions 1.27 X IO-’ M in Mn2(CO)loand 2.55 X low3M in TQ2+u(Table I) were prepared in 0.1 M TBAH/CH,CN. A solution 2.54 X IO-’ M in [Mn(CO)J was generated by controlled-potential electrolysis of at -1.5 V vs SSCE. Exhaustive electrolysis of the TQ2+-containingsolution at -0.6 V vs SSCE gave a solution 2.55 X M in TQ+. For the equilibration experiments, 2 mL of the stock solution containing Mn2(CO)lowas mixed with 2 mL of the solution containing TQ+. In a second solution, 2 mL of the [Mn(CO),]- solution was mixed with 2 mL of the stock solution containing TQ2+. The solutions were allowed to equilibrate for several hours. After that period, UV-vis or FTIR solution cells were introduced into the drybox and filled with the stock solution containing M ~ I ~ ( C Osealed ) ~ ~with , rubber septa, and removed from the drybox, and the UV-vis or FTIR spectrum was recorded. The cell was reintroduced into the drybox and rinsed with 0.1 M TBAH/CH3CN, and a second solution added. This procedure was repeated until the spectra of all six solutions had been recorded. Over ( 5 ) (a) Howell, J. 0.;Wightman, R. M. Anal. Chem. 1984,56,524. (b) Wipf, D. 0.; Kristensen, E. W.; Deakin, M. R.; Wightman, R. M. Anal. Chem. 1988,60, 306. (c) Montenegro, M. I.; Pletcher, D. J. J. Elecfroanal. Chem. Interfacial Electrochem. 1986, 200, 371. (d) Howell, J. 0.; Goncalves, J. M.; Amatore, C.; Klasinc, L.; Wightman, R. M.; Kochi, J. K. J. Am. Chem. SOC.1984,106,3960. (e) Wightman, R. M. Anal. Chem. 1981,53, 1125A. ( f ) Howell, J. 0.;Kuhr, N. G.; Ensman, R. E.; Wightman, R. M.J. Electroanal. Chem. Interfacial Electrochem. 1986, 209, 77.

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Volts (vs SSCE)

Figure 1. Cyclic voltammograms recorded in 0.1 M TBAH/CH3CN at a 0.2-cm-diameter Pt disk electrode at a scan rate of 200 mV S-I: (A) 2.0 X IO-, M Mn,(CO)lo, (B) 2.0 X lo-’ M [Mo(~~-C,H,)(CO),]~, (C)

2.0 X

M [Fe(q5-CsHs)(CO)2]2.

the time span of the experiment, the solutions were stable, and there was no leakage of air into the cell.u For measurements at temperatures other than ambient, the same procedure was followed. Solution compositions at equilibrium were determined by UV-vis measurements. The cell was sealed before removal from the drybox. It was allowed to equilibrate in a constant temperature bath for at least 1 h, at -15, 20, or 60 OC, before measurements were conducted. Complications arising from decomposition of Mn2(CO)loat 60 OC were overcome by using a kinetics program that monitored absorbances at the maxima for Mn2(CO)loat 270 and 380 nm at 10-min intervals. Decomposition led to absorbance decreases at both maxima. Equilibrium was taken to have been reached when the absorbances had reached their maximum values. Temperature-dependent electrochemical measurements were made in a thermostated, three-compartment ce1L6 The cell was constructed with a water bath surrounding the three compartments and a purge line for air-sensitive materials. A sufficiently long salt bridge to the SSCE reference electrode was used so that the reference couple was isothermal and unaffected by the temperature changes that occurred in the working electrode compartment of the cell. Rapid-scan, microelectrode cycle voltammetry was performed inside a drybox. The first scan on freshly polished electrodes was recorded. The electrodes became passive after a series of scans. This complication prevented the use of a flow cell and background subtractions.sf

Results Cyclic Voltammetry. The cyclic voltammetric properties of are well documented in the literature.’ The voltammogram in Figure IA is in agreement with previous reports. In 0.1 M TBAH/CH,CN, Mn2(CO)loundergoes a two-electron, irreversible reduction to give [ M ~ I ( C O ) ~at ] - E$ = -1.45 V vs SSCE. Irreversible oxidation of [Mn(CO)5]-occursat EpB = -0.12 V to give Mn2(CO)lo. At E; = 1.55 V, Mn2(CO)loundergoes an irreversible, two-electron oxidation to give [ M ~ I ( C O ) ~ (a) Yee, E. L.; Cave, R. J.; Gwyer, K. L.; Tyma, P. D.; Weaver, M. 1131. (7) (a) Lacombe, D. A.; Anderson, C. P.; Kadish, K. M.Inorg. Chem. 1986,25,2074. (b) Lemoine, P.;Giraudeau, A.; Gross, M. Electrochim. Acta 1976, 21, 1. (c) Picket, C. J.; Pletcher, D. J. J. Chem. SOC.,Dalton Trans. 1975, 079. (6)

J. J. Am. Chem. SOC.1979, 101,

Pugh and Meyer

3786 J. Am. Chem. SOC.,Vol. 114, No. 10, 1992 1h,

1

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VOLTS (VI. SSCE) Figure 3. Same as Figure 2: (A) (solid line) 1 X

I 1 0.0

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VOLTS (vs. SSCE) Figure 2. Same as Figure 1 but with a scan rate of 100 mV s-l: (A) 1

1 -2.0

lo-’ M [Os(5,5’-

(CO,NEt,),-bpy)(PPh,),(CO)H]+, (dashed h e ) with 1 X l(r3 M added Mn2(CO)lo, (B) (solid line) 1 X lo-’ M [0~(bpyz)(PPh,)~(CO)H]+, (dashed line) with 1 X lo-’ M added Mn2(CO)lo, (C) (solid line) 1 X lo-’ M [0s(5,5‘-(CO2Et),-bpy)(PPh,),(CO)H]+, (dashed line) with 1 X M added Mn2(CO)lo. The ligand bpz is Z.Z’-bipyrazine.

X M Mn2(CO)lo,(B) (solid line) 1 X lo-’ M [Os(bpy)(PPh,),and 1 X (CO)H]+, (dashed line) 1 X lo-) M [0~(bpy)(PPh~)~(C0)H]+ lo-) M Mn2(CO)lo.

(NCCH,)]+. At E; = -1.1 V, [Mn(CO),(NCCH,)]+ is reduced irreversibly to give MII~(CO),~. The same pattern of waves was observed for [Mo(qs-C,H,)(C0),I2, with irreversible reduction occurring at E = -1.10 V, and irreversible oxidation at E; = 0.90 V,Figure 1b. Oxidation of [Mo(q5-C5H,)(CO),]-occurred at EpB = -0.01 V,and reduction at E$ = -0.68 V.8 of [Mo(qS-CSHs)(C0),(NCCH3)]+ For [Fe(qs-C5H5)(C0)2]2, there was an irreversible reduction at E; = -1.80 V and an irreversible oxidation at E,B = 0.55 V, Figure 1C. The anion [Fe(qS-C,H5)(CO),]- was oxidized irreversibly at EpB = -0.93 V, and the cation [Fe(qS-C,H,)(C0),(NCCH3)]+reduced irreversibly at E$ = -0.68 VO8 Redox Catalysis and Redox Equilibration. The irreversibilities of the two-electron reductions or oxidations precluded the measurement of El12values by cyclic voltammetry. The potentials of these couples were measured by redox equilibration. Couples of known potential were mixed with the couple to be measured, and the equilibrium constant was determined after the system had reached equilibrium. The couples that were available for the redox equilibration experiments and their EI12values are listed in Table I. These couples undergo rapid, reversible electron transfer at E I I 2values that extend from +1.05 to -1.47 V in stepwise intervals that vary from 0.37 to 0.03 V. Four different types of couples were available: (1) oxidative, M(III/II) couples of polypyridyl complexes of Ru, Os,or Fe, (2) the Fe(III/II) couples of ferrocene or ferrocene derivatives, (3) bipyridinium dication/cation couples, and (4) reductive, polypyridyl-based couples of complexes of Ru or Os. All potentials cited in Table I were measured in 0.1 M TBAH/CH3CN vs the SSCE reference electrode. An initial screening of the appropriate potential range for the redox equilibration experiments was made by searching for redox catalysis? Examples are shown in Figures 2 and 3. The effect (8)(a) Dessy, R. E.; Weissman, P. M.; Pohl, R. L. J . Am. Chem. Soc. 1966,88,5117. (b) Dessy, R.E.; Stary, F. E.; King, R. B.; Waldrop, M.J . Am. Chem. Soc. 1966,88,471.(c) Ferguson, J. A.; Meyer, T. J. Itwrg. Chem. 1971, 10, 1025. (d) Denisovitch, L. I.; Gubin, S. P.; Chapovskii, Y.A,; Ustynok, N. A. Bull. Acad. Sci. USSR,Diu. Chem. Sci. (Engl. Traml.) 1971, 20, 1851. (e) Kadish, K. M.; Lacombe, D. A.; Anderson, J. E. Inorg. Chem. 1986,25,2246.(f) Davies, D. G.; Simpson, S.J.; Parker, V. D. J . Chem. Soc., Chem. Commun. 1984,352. (g) Mcyer, T.J. Prog. Inorg. Chem. 1975,19, 1.

Wavenumbers (cm”) Figure 4. Fourier transform infrared absorption spectra of solutions in

0.1 M TBAH/CH3CN: (A) initially 1.27 X 10” M in Mn2(CO)loand 2.55 X lo-’ M in TQ+, (B)initially 2.54 X lo-, M in [Mn(CO)S]- and 2.55 X lo-’ M in TQ2+, (C) 2.54 X lo-’ M in Mn2(CO)lo,(D) 5.08 X lo-’ M in [Mn(CO)S]-. of added [0~(bpy)(PPh,)~(C0)H]+ on the reduction of Mn2(CO),ois shown in Figure 2B. This reduction occurs via a sharp, irreversible wave at E = -1.30 V consistent with electron-transfer The reversible wave at Ellz = -1.47 V is catalysis, eqs 3 and a r*(bpy)-based reduction in [Os(bpy)(PPh,),(CO)H]+. Catalysis occurs by reduction of [ 0 ~ ( b p y ) ( P P h ~ ) ~ ( C 0 ) Hto] +[Os(bpy)(PPh,),(CO)H]O at the electrode followed by its reduction of Mn2(CO),o.

8.

~[O~(~PY)(PP~~)~(C +O ) H I + 2 2 =2i[ 0 ~ ( b p y ) ( P P h ~ ) ~ ( C O )(3) H]~

-

~ [ O S ( ~ P ~ ) ( P P ~ , ) ~ ( C O+) HMnz(C0)io IO 2 [ O s ~ b ~ +~2[Mn(C0)51~ ~ ~ (4) ~ (9)(a) Astruc, D. Angew. Chem., Int. Ed. Engl. 1988, 27,643. (b) Evans, D. H.Chem. Rev. 1990, 90,739. (c) Andrieux, C.P.; Gdlis, L.; Medebielle, M.; Pinson. J.; SavCnt, J.-M. J . Am. Chem. Soc. 1990, 112, 3509.

~

~

~

Redox Couples Involving Metal-Metal Bonds

J . Am. Chem. SOC.,Vol. 114, No. 10, 1992 3787

Table 11. Equilibrium Constants and Formal Potentials in 0.1 M TBAH/CH,CN

reaction Mn2(CO)lo+ 2TQ+ * 2[Mn(CO)J + 2TQ2+ Mn2(CO)lo+ 2e- 2[Mn(CO)J Mn,(CO)lo 2PQ2++ 2CH,CN * 2[Mn(CO)5(NCCH3)]++ 2PQ+ ~MII(CO)~(NCCH,)+ + 2e- Mn2(CO)lo+ 2CH,CN [Fe(Cp)(C0),lz + 2[R~(bpy)~l+ + 2[Fe(C~)(C0)~1+ 2[Ru(bpy),I2+ [F~(CP)(CO)~I, + 2e- 2[Fe(Cp)(CO)21[Fe(Cp)(CO),], + 2DQ2++ 2CH,CN + 2[Fe(Cp)(C0),(NCCH3)]++ 2DQ+ [Fe(Cp)(CO),(NCCH,)]+ + 2e- [Fe(Cp)(CO),], + 2CHjCN [Mo(Cp)(CO),], + 2PQ+ ==Z[Mo(Cp)(CO),f- + 2PQ2+ [Mo(Cp)(CO),I2 + 2e- 2[Mo(Cp)(CO),I+ 2Os(bpy)2C12 [Mo(Cp)(CO),], + 2[Os(bpy),C12]' + 2CH3CN * ~[MO(CP)(CO),(NCCH,)~+ [MO(C~)(CO),(NCCH,)]~ + 2e- [Mo(Cp)(CO),],+ + 2CH,CN 'Calculated by using eq 7.

+

-

K 1.6 (f0.3)

-

Mnz(CO)lo+ 2TQ+ K=

2[Mn(CO)J

+ 2TQ2+

[~M~(CO)sl-lZ[TQ2+12

-0.30 (10.3) 1.7 (ho.3) x

-1.44 (k0.05) 10-5

-0.20 (fO.02) 1.19 (fO.O1) X lo4 -0.56 (fO.01) 4.2 (fO.3)

X

-0.07 (10.06)

',

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1

I

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I

I

300

350

400

450

500

550

600

Wavelength (nm) Figure 5. UV-vis absorption spectra of a solution that was initially 1.27 X lo-) M in Mn2(CO)loand 2.55 X lo-, M in TQ+ in 0.1 M TBAH/ CHJCN.

1

-1

3.0 2.5

(6)

2.0

4-

1.5

-

Eo'(Mnz(CO)lo/2[Mn(CO)S]-) = Eo',ef+ (RT/nF) In K

7.8 (f0.6) X lo4

250

0.0

(5)

[Mnz(CO) io1 [TQ+l r=O.lM, 22f2OC) = 1.6 (f0.3)X was obtained regardless of the direction in which the equilibrium was approached. The formal potential for the Mn2(CO)lo/2[Mn(CO)s]- couple was calculated from the equilibrium constant and the potential of the equilibrating couple by using eq 7 in which Eo'ref(=Ellz)is the potential for the equilibrating couple, F is the Faraday constant, T is the temperature (K), and n is the electrochemical stoichiometry. Equilibrium constants and formal potentials are listed in Table 11.

V vs SSCE"

-0.69 (fO.O1)

5 ( ~ 5 x) 10-5

-

The same experiment was repeated but with the bpy-modified complexes mentioned in the caption to Figure 3 which had E l / 2 values for the first A* reduction at -1.10 or -0.80 V. As Ell2for this reduction was made more positive in the catalyst, the extent of catalysis at constant sweep rate decreased.9c For [Os(5,5'(C0,Et) z- bpy) (PPhJ 2( CO) H ] +,Figure 3C, slow catalysis occurred at the first ligand-based reduction and rapid catalysis at the second. From the results of these experiments it can be estimated that Eo < -0.8 V vs SSCE for the Mn2(CO)lo/2[Mn(CO)J couple. In order to define the potential more precisely, redox equilibration experiments were undertaken. An infrared spectrum of a solution initially containing TQ+ and Mn2(CO),o in 0.1 M TBAH/CH3CN in the v, region is shown in Figure 4A. The peaks at 2045, 201 1, and 1985 cm-I are v, stretches for Mn2(CO),,, and the peaks at 1900 and 1894 cm-l are v m stretches for [ M ~ I ( C O ) ~ ] - Under . the conditions of the experiment, equilibrium between Mnz(CO)loand [Mn(CO),]- was reached within 6 h. After equilibration, the concentrations of Mn2(CO)lo and [Mn(CO)J were determined by measuring the areas under the CO bands and comparing them with preestablished calibration curves. The concentrations of TQ+ and TQz+ were surmised by mass balance by assuming the equilibrium in eq 5 , and the four concentrations used to calculate the equilibrium constant in eq 6. The same value for the equilibrium constant, K(CH3CN,

EO',

X

8 8 e SI


lo0 V s-I)I" or in 1,2-di-

0.1 M TBAH/CH3CN, 22 f 2OC, in V vs SSCE

Discussion Electrochemistry. In CH3CN, Mnz(CO)loundergoes an irreversible electrochemical oxidation to give [Mn(CO),(NCCH,)]+ and an irreversible reduction to give [Mn(CO),]-. The loss of structural integrity is typical for metal-metal bonds where electrons are gained or lost from orbitals that are largely bonding or antibonding with regard to the metal-metal interaction. There was no evidence for an intermediate [Mn2(CO)lo]o/-couple by rapid-scan, cyclic voltammetry even at scan rates of up to 10000 V s-I. Likewise, the reduction of Mnz(CO)lois chemically irreversible even in propionitrile at low temperatures.'" For the irreversible reduction wave for Mn2(CO)lothat appears at E$ = -1.45 V, the peak to half-peak potential separation, IE - EPlzl= 136 mV, is considerably in excess of the 59 mV expectd for a diffusion-controlled reaction. Both the peak shape and scan rate dependence are consistent with initial slow electron transfer with a transfer coefficient, an = 0.4. Oxidation of [Mn(CO),]at the electrode is rapid, IEp- E 2) = 59 mV as expected for a diffusion-controlled reaction. $he reduction of [Mn(CO),(NCCH3)]+is also slow, )Ep- Ep/21= 94 mV. The mechanism that has been proposed for electrochemical reduction of Mnz(CO)lois shown in eqs 12-14.' The electron-

Mnz(CO),o + e[Mnz(CO)ioIMn(CO),

[Mnz(CO)101[MdCO)sI- + Mn(CO)s

+ e-

[Mn(CO),]-

=i

(12)

fluorobenzene a t moderate scan rates.', On the basis of the electrochemical results, it can be inferred that the reductions of Mnz(CO)loor [Mn(CO),(NCCH,)]+ at the electrode occur by slow, rate-limiting electron transfer while the oxidations of Mn2(CO),o or [Mn(CO),]- are rapid, diffusion-controlled processes. The value of E,,, = -0.08 V vs SSCE for the Mn(CO),/ [Mn(CO),]- couple is within experimental error of the value of -0.143 V obtained by Tilset and Parker by using derivative cyclic voltammetry and kinetic shift factors.16 A value of 0.08 V was reported by Kuchynka and Kochi, by fast-scan, microelectrode cyclic v01tammetry.l~~The discrepancy between our value and theirs may lie in the waveform distortions that exist a t high scan rates. Both [MO(~~-C,H,)(CO),]~ and [Fe(~5-C,H,)(CO)z]z have also been the subject of detailed electrochemical investigations.8 Early results obtained at mercury electrodes may have been complicated by adsorption and the formation of mercury compounds." More recently, reversible, one-electron oxidations14band r e d ~ c t i o n s ' ~ ~ have been reported for [Fe(~5-C,Hs)(C0)z]2. Attempts to establish potentials for the one-electron couples involving [MO(~~-C,H,)(CO),] by rapid-scan, microelectrode voltammetry were complicated by an overlap in potentials. Oxidation of [Mo(q5-CSHs)(CO),]- gave both [Mo($-C,HS)(CO),(NCCH,)]+ and [Mo(.rlS-C,H,)(CO),],and their reduction waves at EpC = -0.30 V and at EpC = -0.20 V overlapped and could not be resolved. The acetonitrile complex appears because Mo($-C,H,)(CO), is unstable with respect to disproportionation and at any potential is oxidized (as MO(~~-C~H~)(CO),(NCCH,))'~ required to oxidize the anion (see below). Because of the overlap problem, the E l I 2value determined by Tilset and Parker for the [MO(~~-C,H,)(CO),]~/couple was adopted.I6 The estimate made of for the MO(~~-C,H,)(CO)~/couple by Tilset and Parker utilized the rate constant for dimerization of Mo(q5-C5HS)(CO),in order to use kinetic shift factors. This could lead to some error in their calculated E l I 2values because the dimerization rate constants have not been measured in 0.1 M TBAH/CH3CN. From the results of flash photolysis experiments these rate constants are solvent dependent.24 Despite this, our results for Mn(CO), and Fe($-C5HS)(CO)zagree well with C O ) , to be within theirs, and their value for M O ( ~ ~ - C ~ H ~ ) (appears the experimental uncertainty of the value that we could have determined directly in the absence of the overlap problem. The apparent reversibility of the [Fe(tlS-C,H,)(CO),],/[Fe(~s-C,H,)(CO)z]couple a t scan rates as low as 879 V s-I is inconsistent with the reported dimerization rate constant of 3 X lo9 M-' s-l determined by flash photolysis,19 and even less consistent with reversibility at rates as low as 1 V s-' repofi.4 earlier.8f

(13) (14)

transfer reaction in eq 15 could also play a role depending on the rate of cleavage of the Mn-Mn '/zbond in [Mnz(CO),o]-. The reactions involving Mn(CO), are in competition with recombi(14) (a) Dalton, E. F.; Ching, S.; Murray, R. W. Inorg. Chem. 1991,30, 2642. (b) Bullock, J. P.; Palazotto, M. C.; Mann, K. R. Inorg. Chem. 1991, 30, 1284. (c) Dalton, E. F. Ph.D. Dissertation, University of North Carolina, Chapel Hill, NC, 1990.

(15) (a) O'Twle, T. R.; Younathan, J. N.; Sullivan, B. P.; Meyer, T. J. Inorg. Chem. 1989, 28, 3923. (b) O'Twle, T. R. Private communication. (16) Tilset, M.; Parker, V. D. J . Am. Chem. SOC.1989, 1 1 1 , 6711. (17) Moran, M.; Cuadrado, I.; Losada, J. J . Orgammet. Chem. 1987,320, 317. (18) (a) Stiegman, A. E.; Stieglitz, M.; Tyler, D. R. J . Am. Chem. SOC. 1983,105,6032. (b) Goldman. A. S.;Tyler, D. R. J . Am. Chem. SOC.1984, 106, 4066. (c) Fei, M.; Sur, S. K.; Tyler, D. R. Organometallics 1991, 10, 419. (d) Zhang, Y.; Gosser, D. K.; Rieger, P. H.; Sweigart, D. A. J . Am. Chem. SOC.1991,1131, 4062. (e) Ryan, 0. B.;Tilset, M.; Parker, V. D. J . Am. Chem. Soc. 1990,112, 2618. (19) Caspar, J. V.; Meyer, T. J. J . Am. Chem. Soc. 1980, 102, 7794.

Pugh and Meyer

3790 J . Am. Chem. SOC.,Vol. 114, No. 10, 1992 Table 111. Free Energies of Dissociation in 0.1 M TBAH/CH,CN a t 22 i 2 o c

compound Mn2(CO)lLl IFe(?5-C5H5)(C0)212 IMO(V’-CJ%)(CO)J~

AG’lu-M, kcal mol-’ 28 & 4 25 & 5 22 f 3

KU-M

2.4 i 4.0 x 10-19 7.1 x 10-17

It is possible that the cathodic wave observed at E, = -0.9 V is due to reduction of the long-lived CO-bridged complex that was reported to form following flash photolysis of [Fe($-C,H,)(CO)2]2.19920If so, access to this intermediate is also possible via oxidation of [F~(T$C,H,)(CO),]a t an electrode. The equilibration reactions are reversible; no products other than [Fe(.r15-CsH,)(CO)2(NCCH~)]+, [Fe(q5-C,Hs)(CO)2]-, or [ F e ( ~ ~ C , H , ) ( C O ) J were 2 detected. Despite the experimental complications, the El value determined by Tilset and Parker for the [Fe($-CsHs)(Cd)2]o/- couple is consistent with the reversible wave observed by rapid-scan cyclic voltammetry.I6

Stabilities of the Metal-Metal Bonds. Free Energies of Dissociatioa. When the results of the redox equilibration and fast-scan

experiments are combined, it is possible to calculate the free energy change associated with dissociation of the metal-metal bonds. The relationship between this quantity, AGO’M-M, and the formal potentials for the one- and two-electron couples is shown in eq 17. The values of AGO’M-M, the free energy change in 0.1 M TBAH/CH3CN calculated by this method, and the associated equilibrium constants, KM-M, are listed in Table 111. AGo’(Mn2(CO)lo)(eV) =

-2[E0’(Mnz(CO),,/2[Mn(CO),1-) -

EO’(Mn(CO)s/ [Mn(CO)sI-)I (17) From the temperature dependence of K,ASo’= 15 eu in 0.1 M TBAH/CH3CN for the equilibrium in eq 5 . When this value is combined with ASo’ for the TQ2+/+couple, a value of ASo‘ = -14 eu is derived for the Mn2(C0),,/2[Mn(CO),]- couple. In order to calculate the entropic change associated with dissociation of the metal-metal bond in Mn2(CO)lo,it would be necessary to measure the temperature dependence of the Mn(CO),O/- couple. It was not possible to perform these measurements with any degree of accuracy because of the large uncertainties in the rapid-scan experiments. An independent estimate of AS = 32 eu for the metal-metal bond equilibrium has been made by flash photolysis in cyclohexane.’” This may be a reasonable value for ASo’in CH3CN since both components in the reaction are neutral and Os should be relatively independent of solvent. If this value is adopted, AHO’M-M = 38 kcal mol-’ for Mn2(CO)lo a t 298 K. Since the first estimate of the M-M bond strength in Mn,(CO)ioby Cotton and Monchamp,2Iathere have been many attempts to measure this enthalpic change. Values ranging from 19 to 41 kcal mol-’ have been obtained by using a variety of techniques. The values derived from electron impact studies21d-g are now ~~~~

~~~~~

~

~~~

(20) (a) Tyler, D. R.; Schmidt, M. A.; Gray, H. B. J . Am. Chem. SOC. 1979,101,2753. (b) Tyler, D. R.; Schmidt, M. A.; Gray, H. B. J . Am. Chem. SOC.1983, 105, 6018. (c) Hooker, R. H.; Mahmoud, K. A,; Rest, A. J. J. Chem. SOC.,Chem. Commun. 1983, 1022. (d) Hepp, A. F.; Blaha, J. P.; Lewis, C.; Wrighton, M. S.Organometallics 1984,3, 174. (e) Moore, B. D.; Simpson, M. B.; Poliakoff, M.; Turner, J. J. J . Chem. Soc., Chem. Commun. 1984,972. (21) (a) Cotton, F. A,; Monchamp, R. R. J . Chem. SOC.1960,533. (bl Bidinosti, D. R.; McIntyre, N. S.J . Chem. Soc.,Chem. Commun. 1966,555. (c) Hopgood, D.; Pol, A. J. J . Chem. SOC.,Chem. Commun. 1966,831. (d) Svec, H. J.; Junk, G. H. J . Am. Chem. SOC.1967,89,2836. (e) Svec, H. J.; Junk, G.H. Inorg. Chem. 1967, 7, 1688. (f) Junk, G.H.; Svec, H. H. J. Chem. SOC.A 1970, 2102. (g) Bidinosti, D. R.; McIntyre, N. S. Can. J . Chem. 1970, 48, 493. (h) Hughey, J. L.; Anderson, C. P.; Meyer, T. J. J . Organomet. Chem. 1977, 125, C49. (i) Stevens, A. E. Ph.D. Dissertation, California Institute of Technology, 1981. (j)Connor, J. A.; Zafarani-Moattar,

M. T.; Bickerton, J.; El Saied, N. I.; Suradi, S.;Carson, R.; AI Takhin, G.; Skinner, H. A. Organometallics 1982,1, 1166. (k) Coville, N. J.; Stolzenberg, A. M.; Muetterties, E. L. J . Am. Chem. Soc. 1983, 105, 2499. (I) Goodman, J. L.; Peters, K. S.;Vaida, V. Organometallics 1986,5,815. (m) Simkes, J. A. M.; Schultz, J. C.; Beauchamp, J. L. Organometallics 1985. 4, 1238.

believed to be unreasonably A value derived from thermochemical data has large uncertainties.21j Our value is within experimental error of the original estimate. It is also consistent with a value obtained by time-resolved, photoacoustic calorimetry in hexanes.2i’ A more recent gas-phase estimate gave 41 kcal mol-’, with an uncertainty of 9 kcal mol-1.21m If the near constancy in AHO’M-M implied by the gas-phase and solution experiments is correct, the Mn2(CO) Io/ 2 [Mn( CO) ,] equilibrium is nearly medium independent and the solvation enthalpies for Mn2(CO),oand Mn(CO), must be the same to within a few kilocalories per mole. On the basis of the values of AGO’M-M listed in Table 111, the metal-metal bonds in these compounds are weak compared to typical chemical bonds. In the cis and trans CO-bridged isomers of [Fe(q5-C,H5)(C0),], there is no direct metal-metal and AGO‘M-M is a measure of the strength of the electronic coupling that exists through the CO bridges. On the basis of the results of NMR and infrared measurements, a nonbridged isomer also exists. Since the nonbridged isomer is in measurable equilibrium with the cis and trans isomers, AGO’M-M for [Fe(q5C,H5)(C0)2]2also provides a reasonable estimate of the metal-metal bond strength. It is in agreement with a previous determination of m h 1 - M ~ ~ ~ Reduction Potentials. Latimer Diagrams. Latimer diagrams, for couples that interrelate Mn,(CO)lo, [Fe(q5-C,H,)(CO)2]2, or [ M O ( ~ ~ - C , H , ) ( C O ) ,and ] ~ their adjacent oxidation states are shown in Scheme I. The potentials cited came from the following sources: (1) The two-electron couples came from the redox equilibration experiments. (2) The [Mn(C0),lo/- and [Fe(qSC5H,)(CO)2]o/- couples came from microelectrode cyclic voltammetry, and [MO(~~-C,H,)(CO),]~/from ref 16. (3) The one-electron oxidative couples [Mn(CO),(NCCH,)]+/Mn(CO), were calculated from AGO‘M-M and Eo’ values for the analogous two-electron couples Mn2(C0),,/2[Mn(CO),(NCCH3)]+. (4) The one-electron couples [Mn2(CO)lo]+/o, [ Fe(q5-C5H5)] ~ + /taken ~ as E a - 0.03 (CO)2]2+/o,or [ M O ( ~ ~ - C , H , ) ( C O ) ~were V and [MO(?~-C,H,)(CO),],O/-or [Fe(qW,H,)(CO),]$- as E; 0.03 V. The justification for using peak potentials to calculate E l l 2 values was based on the properties of the waves. They point to diffusion-controlled, heterogeneous electron transfers at the electrode followed by rapid, irreversible chemical steps. This approximation is known to be valid for the couples [Fe(q5c,H~)(cO)212+/O,[Fe(r15-C,H5)(C0)21,0/-, and [Mn2(CO)lol+/o where E I 1 2values are known by direct measurement in noncoordinating so1vents.l4J5 The approximation is not valid for reduction of M I I ~ ( C O because ) ~ ~ of the evidence from the wave shape for slow heterogeneous electron transfer. Chemical shift factors were not included in the calculations of E l l 2for the one-electron dimer couples. The potentials cited are only estimates of the thermodynamic potentials. The values cited for the couples [Mn(CO),(NCCH,>]+/Mnare consistent (CO), and 2[Mn(CO),(NCCH3)]+/Mnz(CO) with estimates made earlier on the basis of a flash photolysis procedure.24 In that study it was estimated that Eo < -0.8 V for [Mn(CO),(NCCH,)]’/Mn(CO), and Eo < -0.2 V for 2-

+

[Mn(Co),(NCCH3)l+/Mn,(CO)

Implications for Reactivity. The thermodynamic data in the Latimer diagrams reveal much about the reactivity of the metal-metal bonds toward electron transfer. Reactivity via the Monomers. On the basis of the equilibrium constant for dissociation of Mn,(CO) (Table 111) and assuming that K = lo9 M-I s-I for recombination of Mn(CO),, the half-time (22) (a) Bullitt, J. G.; Cotton, F. A.; Marks, T. J. J . Am. Chem. SOC.1970, 92,7155; Inorg. Chem. 1972, 11, 671. (b) McArdle, P.; Manning, A. R. J . Chem. Soc. A 1970, 2120. (c) Abrahamson, H. B.; Palazotto, M. C.; Reichel, C. L.; Wrighton, M. S.J. Am. Chem. SOC.1979, 101, 4123. (d) Gransow, 0.A.; Burke, A. R.;Vernon, W. D. J . Am. Chem. SOC.1976,98, 5817. (e) Mitsehler, A.; Rees,B.; Lehmann. M. S. J . Am. Chem. SOC.1978,100,3340. (f) Jemmis, E. D.; Pinhas, A. R.; Hoffman, R.J . Am. Chem. SOC.1980,102, 2576. (9) Benard, M. Inorg. Chem. 1979, 18, 2782. (23) Cutler, A. R.; Rosenblum, M. J. Organomet. Chem. 1976, 120,97. (b) Barrett, P. F.; Sun, K. K. W. Can. J . Chem. 1970, 48, 3300. (24) Meyer, T. J.; Caspar, J. V. Chem. Reo. 1985, 187.

Redox Couples Involving Metal-Metal Bonds

J. Am. Chem. SOC., Vol. 114, No. 10, 1992 3791

for dissociation at 298 K is - 5 X lo5 h. The half-times for the Fe and Mo dimers are less but still appreciable. This means that any mechanism that relies on dissociation prior to electron transfer of the metal-metal bond is necessarily very slow at room temperature. At higher temperatures, dissociation could be more important. By using W ' M - M 38 kcal mol-' for Mn2(CO)lo,Kdissincreases from 2.4 X lo-*' at 295 K to 3.8 X lo-'' at 353 K. For a solution initially 2 X M in Mn2(CO)',, this corresponds to an increase M in the equilibrium concentration of Mn(CO), from 2 X M at 353 K. This concentration level at 295 K to 2.7 X is near that required for direct observation by NMR given the paramagnetic character of the monomers.25 At higher temperatures, loss of CO is known to occur and provides a basis for the reactivity of Mn2(CO)lo toward both oxidation and substitution,21k,2628 From the potentials for the couples [Mn(CO)s(NCCH3)]+/ Mn(COS), etc., the monomers are strong reducing agents. They are accessible ph~tochemically,'~*'~-~~*~~ and their strong reducing character is shown by reactions in which they abstract halogen atoms from halocarbon solvent^.^^.^^ They have both enhanced oxidizing and reducing strengths compared to the parent metal-metal bonds.30 This leads to an instability toward disproportionation in CH3CN for both Mn(CO), and Mo(q5-C5H5)(CO), while Fe(qS-CsH5)(CO)2is stable. 2Mn(C0)5

+ CH3CN

-

[Mn(CO),(NCCH,)]+

+ [Mn(CO)s]-

(18)

AGO' = -19.0 kcal mol-'

+

-

ZMo(Cp)(CO)3 CH3CN [Mo(Cp)(C0)3(NCCH3)1+ + [Mo(Cp)(CO)31- (19) AGO'

+

= -9.7 kcal mol-'

2Fe(Cp)(CO), CH3CN [Fe(Cp)(Co),(NCCH,)I+

+ [Fe(Cp)(CO)zl-

e-[

[Mo(Cp)(CO),Br]-

+ [Mo(Cp)(CO)3BrI-

-

[Mo(C~)(C0)31-+ MO(CP)(CO)$~ (22)

Reactivity of the Jhers. All three metal-metal bonds are stable with respect to disproportionation, with the free energy of disproportionation being +11.3 kcal mol-' for [ M o ( ~ ~ - C ~ H ~ ) ( C O ) ~ ] ~ , and +28.5 kcal mol-' for [Fe(qs-CsHs)(CO)2]2. The reverse reactions are used to prepare the metal-metal

+

Mnz(CO)lo CH3CN

-

[Mn(CO)J

+ [Mn(CO),(NCCH,)]+

AGO' = +9.0 kcal mol-'

One-electron oxidation or reduction at solid electrodes is rapid, except for [Mn2(CO)lo]o/-.The two-electron processes occur via the one-electron intermediates as shown by the appearance of their couples in noncoordinating solvents. This explains the patterns of irreversible waves that appear in the cyclic voltammograms in Figure 1. For example, for Mn2(CO)lo,initial, one-electron oxidation to give [Mn2(CO)lo]+does not occur until E l l 2 = 1.52 V, yet El12= -0.30 for the two-electron couple. This creates an overvoltage of 1.82 V for the oxidation of Mn2(CO)loto [Mn(CO),(NCCH,)]+. Likewise, reduction of [Fe(qS-CsHs)(CO)2]2 to 2[Fe(q5-C,H5)(CO),l2- cannot occur until the potential for reduction to [Fe(q5-CsH,)(CO)2]z-is reached at -1.77 V. This causes an overvoltage of 0.33 eV. The same factors lead to an inhibition to electron transfer in solution as illustrated by the oxidation of [Fe(q5-C5Hs)(C0)2]2 by [ R ~ ( b p y ) ~ C l ~ The ] + . ~slow ~ step in this reaction is the initial electron transfer as shown by the rate law which is first order in both reactants.

-

cis-[Ru( bpy)C12]+ + [Fe(Cp)(CO)z]2

AGO' = +4.8 kcal mol-' (20)

The instability toward disproportionation is driven by the binding of the donor solvent CH3CN to the cations, and by the solvation of the ions that are produced. These calculations provide a rationale for the photochemically- or thermally-induced disproportionations of Mn2(CO)lo,[Mo(q5-CSHs)(CO),l2, or C O ~ ( C O ) ~ in polar solvents, or of Mo(q5-CSHS)(CO),in the presence of halide ions. 12,18,24,29 Both disproportionation and metal-metal bond formation are thermodynamically spontaneous and competitive reactions for the monomers when they are prepared thermally or photochemically. -28 kCaL"0l

Mo(Cp)(CO),

+ Br-

cis-R~(bpy)~C +l ~[Fe(Cp)(C0),lz+ (23)

AGO' = 3.7 kcal mol-'

AG'

Mo(Cp)(CO),

I

Mnz(CO)10

cis-[Ru(b~y)~Cl+ ~ l 2CH3CN + + [Fe(Cp)(C0)212+,,pia cis-Ru( bpy),C12

+ 2 [Fe( Cp) (CO),( NCCH,)]

+

(24)

AGO' = -28.3 kcal mol-'

On this basis, the reactivity order toward oxidation of the metal-metal bonds should follow the potentials of the one-electron couples, [Fe(v5-CsH~)(CO)212> [ M O ( ~ ~ - C ~ H S ) ( C> O)~I~ Mn2(CO),o. The one-electron intermediates are powerful oxidants and powerful reductants. They are highly unstable with respect to disproportionation in polar, coordinating solvents. 2[Mnz(CO),o]+

+ 2CH3CN

-+

Mn2(CO)lo

2[Mn(C0),(NCCH3)]+ (25)

AGO' = -83 kcal mol-'

Metal-metal bond formation is known to be rapid.I2 Disproportionation would be favored by prior binding of the solvent or of an added ligand followed by outer-sphere electron transfer."J8 (25)Wayland, B. B.; Coffin,V. L.; Farnos, M. D. Inorg. Chem. 1988,27, 2745. (26)(a) Haines, L.I. B.; Hopgood, D.; PM,A. J. J . Chem. SOC.A 1968, 421. (b) Fawcett, J. P.; POZ, A.; Sharma, K. R. J. Am. Chem. SOC.1976,98, 1401.

(27)(a) Schmidt, S.P.; Trogler, W. C.; Basolo, F. Inorg. Chem. 1982,21, 1698. (b) Stolzenberg, A. M.; Muetterties, E. L. J . Am. Chem. SOC.1983, 105, 822. (28)Yu,W.; Liang, X.; Freas, R. B. J . Phys. Chem. 1991, 95, 3600. (29)(a) Wrighton, M. S.;Bredesen, D. J. Organomer. Chem. 1973,50, C35. (b) Wrighton, M. A.; Ginley, D. S.J . Am. Chem. Soc. 1975,97,2065. (30)Hepp, A. F.; Wrighton, M. S.J . Am. Chem. SOC.1981,103,1258.

This instability is driven by the formation of a strong bond with the solvent in the ( d ~ or) ~(dT)6 cationic products and by the re-formation of the metal-metal bond. The one-electron cations are also unstable with regard to dissociation. The coordinating solvent plays an important role. In CH2ClZ,[Fe(q5-C5Hs)(CO)2]z+ is stable on the cyclic voltammetric time scale, but in CH3CN, it dissociates rapidly. 14b [Mn2(CO)lo]++ CH,CN

-

[Mn(CO)S(NCCH3)]++ Mn(CO)S

AGO' = -27 kcal mol-'

Where data are available, the one-electron-reduced anions are also unstable toward disproportionation; AGO' = -22.5 kcal mol-' for [Mo(q5-C5HS)(CO),l2-. (31) Hughey, J . L.; Meyer, T. J. Inorg. Chem. 1975,14, 947.

Pugh and Meyer

3792 J . Am. Chem. SOC.,Vol. 114, No. 10, 1992 2 [WCp)(CO)212-

-

[Fe(Cp)(C0)212 + 2[Fe(Cp) (CO),l-

AGO' = -14.3 kcal mol-l

The anion [Fe(qS-C,HS)(CO),]y is stable with respect to dissociation, but [MO(~~-C,H,)(CO),]~is not (AGO' = -0.27 kcal mol-I).

-

[Fe(q5-C~H~)(Co)~lz [Fe(q5-CsHd(C0)21- + Fe(7S-C~Hs)(Co)2 AGO' = +5.4 kcal mol-I The Adjacent OxidationSlates. The requirement of one-electron

oxidation or reduction of the ( d ~ or) ~(d7r)6 cations or the corresponding anions also causes the overvoltages that cccur in their vol tammograms.

+

-

[Mn(CO)S(NCCH3)]+ eMn(CO),(NCCHJ [Mn(CO)S]-

Mn(CO),

=i

Mn(CO)S

+ CH3CN

+ e-

-

~[MO(CP)(CO),(NCCH,)I++ Mn2(CO)lo [Mo(CP)(CO),I~+ ~ I M ~ ( C O ) S ( N C C H ~ ) I + AGO' = -10.6 kcal mol-]

These results are qualitatively consistent with the results reported recently by Corraine and Atwood.Il A kinetic inhibition is expected in either case since the initial one-electron steps are nonspontaneous and the AGO' values represent minimum free energies of activation. [Fe(Cp)(CO)21-

-

+ [Mo(CP)(CO)~I~

Fe(Cp)(CO), AGO'

+ [Mo(Cp)(CO),lz-

= +3.4 kcal mol-'

-

[Mo(Cp)(CO),(NCCH,)I+ + Mn2(CO)l~ Mo(CP)(CO)~+ [Mn2(CO)lol+ AGO' = +46.0 kcal mol-'

Acknowledgments are made to the Office of Naval Research Overvoltages are +1.82 V for 2[Mn(CO)S(NCCH3)]+/Mn2for support of this research under Grant Number NOOO14-87(CO),, and +0.51 V for [MO(~~-C~H,)(CO),]~/~[MO(~~K-0430. C,H,)(CO),]-. For electron transfer in solution the order of reactivity as one-electron oxidants should fall in the order [MoRegistry NO.T Q', 67509-62-0; PQ2', 4685-14-7; DQ2', 2764-72-9; (tls-CsHs)(CO),(NCCH,)l+ > [Fe(?'-CsHs) ( C ~ ) Z ( N C C H ~ ) I + TBAH, 3109-63-5;Mn2(CO)lo, 10170-69-1;Mn(CO)5(NCCH3)t, > [Mn(CO),(NCCH,)]+ and as reductants [Fe(qs-CsH,)(CO)z]27674-37-9;[Fe(Cp)(C02)12,12154-95-9;[R~(bpy)~]', 56977-24-3; >> [MO(~~-C,H,)(CO)~]-, Mn(CO)S-. This assumes that reorg[Fe(Cp)(CO),(NCCH,)] ', 46134-92-3;[Mo(Cp)(CO),]2, 1 209 1-64-4; [Os(bpy)2C12]', 15702-72-4;[MO(Cp)(CO)3(NCCH,)]', 68868-63-3; anizational energies are comparable or negligible. [Mn(CO)J, 14971-26-7;[Fe(Cp)(C02)]-, 12107-09-4;[MO(Cp)The anion [Fe(qS-C,H,)(CO)2]-is capable of reducing either (CO),]-, 12126-18-0;[Os(bpy)(PPh3)2(CO)H]+, 841 17-36-2;[OSMn2(CO)10 or IMo(~s-c,H,~(co)312. (bpy)(PPh,)2(CO)H], 139759-59-4; [0~(5,5'-(COzNEt2)2-bpy)(PPh3)2(CO)H]', 139869-56-0;[Os(bpy~)(PPh3)2(CO)H]',89689-63-4;[OS2[Fe(Cp)(CO),I- + [Mo(Cp)(CO),I, (5,5'-(C02Et)2-bpy)(PPh3)2(CO)H]', 139759-65-2;[Mn(CO)5], [Fe(CP)(CO)zIz + 2 [ M o ( ~ 5 - ~ , H , ) ( c o ) , l 14971-26-7;[Mn2(CO)lo]t, 119886-99-6;[Mn2(CO),o]-, 64539-72-6;

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AGO' = -40.5 kcal mol-'

The cation [Mo(q5-CsH5)(CO),(NCCH3)]+ is capable of oxidizing either or [Fe(qS-CsH,)(CO)2]2.

91760-55-3;[Fe(Cp)[Fe(Cp)(CO),], 55009-40-0;[Fe(Cp)(C0)J2', (CO)2]2-,91864-52-7;[MO(Cp)(CO)J, 12079-69-5; [MO(Cp)(CO),]2', 139759-60-7;[MO(Cp)(C0)3]2-,139759-61-8;CH3CN,75-05-8;Fe, Mo,7439-98-7;Mn,7439-96-5. 7439-89-6;