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The observed results can now be explained in terms of inductive effects and the bond elongation. Although the a-carbon in tert-butyllithium is potentially more capable of transferring electron density t o the lithium than the a-carbon in ethyl- and sec-butyllithium based on inductive effects, it is in actuality less capable because of the drastically increased Li-C bond distance. Ethyl- and sec-butyllithium, therefore, absorb at a lower energy than tert-butyllithium and our rationale of predicting the order of transitions within a specific aggregate unit seems justifiable. Thus the alkyllithium aggregates have been shown t o have electronic transitions observable in the ultraviolet region which may be generally associated with a charge transfer from the alkyl group to the lithium framework. Both the experimentally observed spectra and the transitions calculated using C N D 0 / 2 procedures support these conclusions and also support the relative energies of the transitions observed. Based on these conclusions, one may further predict that the photochemical reactions of alkyllithiums will occur primarily at high energy and involve formation of radicals by the charge transfer process. This con-
tention is supported by Glaze6 who postulates an interaggregate disproportionation mechanism, e.g., (2), (CzHki)B* +.C~Hs(CzHsLi)s 4-Li
+ .CzH4Li(CzH4Li)4 .Cd-14Li(CzHaLi)4+C1H4 + (CzH6Li)d -I-Li
* C Z H ~ ( C Z H+ ~ L ~ )CZHB ~
(2)
(3)
(4)
(3), and (4) t o account for the photodecomposition products of ethyllithium. The (C2H5Li)4may in turn equilibrate to form more hexamer or eliminate LiH and ethylene or both. This “not so free” radical-type intermediate may be common to all photolytic reactions of alkyllithiums at high energy. Acknowledgments. The authors would like to express their thanks to the Formica Co., Ethyl Corporation, and Wayne State University for aid in the form of fellowships (to R. J. H.) and to the Research Corporation and the Greater Cincinnati Foundation for direct financial assistance. The authors also wish to thank Dr. Darrell Ebbing for helpful discussions. The National Science Foundation, Grant GP-8323 and 19299, also provided partial support for this work.
Preparation and Characterization of an Inorganic Analog for the Active Site of the Reduced 2Fe-S” Iron-Sulfur Proteins’ Kenneth A. Rubinson and Graham Palmer* Contribution from the Department of Chemistry, University of Michigan, Ann Arbor, Michigan 48104, and the Biophysics Research Division, Institute of Science and Technology, University of Michigan, Ann Arbor, Michigan 4810.5. Received May 31, 1972 Abstract: The reaction of [(CF,),C,S,] (ligand) with Fe(COIs in the presence of H2Sproduces a black crystalline solid, I (apparent space group p4,32 or p4132), which, when dissolved in hydrocarbon solvents, exhibits many of the physical properties of the reduced two-iron iron-sulfur proteins. A frozen solution of I in hexane exhibits an epr spectrum which upon addition of isopropyl disulfide narrows and has an axial g tensor (gL = 1.9880, gll = 2.0098, from computer simulation) similar to that observed in reduced adrenodoxin and putidaredoxin. The integrated intensity of the epr corresponds to one unpaired electron per two Fe. The magnetic susceptibility measured between 20 and 300 “K exhibits non-Curie behavior which can be described using the antiferromagnetic interaction Hamiltonian X = -US1.& with J = - 9 5 cm-1 and SI - Sz= 1. An endor measurement gave values for the I9F hyperfine constant of A , , ’ = 1.25 MHz and A,‘ = 0.78 MHz, from which it was concluded that the unpaired electron resides almost exclusively on the Fe-S center. There is a hydrogen in the molecule which acts as a counterion allowing the complex to be soluble in nonpolar solvents while having the irons in different formal oxidation states. The instability of the complex to Lewis bases parallels the chemistry of hydrogen polysulfides. Spectral properties of the complex were determined from 55,000 to 200 cm-1. Changes were observed following addition of isopropyl disulfide in 19F nmr, Mossbauer, and epr spectra but not in the optical measurements. The near-infrared electronic and far-infrared vibrational spectra of { Fe[SzCz(CFb)& ] 2, { CO[S,C2(CF3)2]z ] 2, and [(CF3)GSzJ~ and the Raman spectra of I and [(CF3)2C2S2]2 are reported. A convenient method for obtaining Raman spectra of highly colored solutions is described.
T
he iron-sulfur proteins (ferredoxins) are an intriguing and biologically significant class of compounds with the characteristic property that, when reduced, (1) Acknowledgment is made for the receipt of Predoctoral Fellovrship 5-FOl-GM-37,338 from the National Institute of General Medical Sciences, USPHS from July 1969 through Oct 1971, and NIH Research Grant GM12176. The Raman spectrometer (Spex Ramalog system) was provided to the Chemistry Department, University of Michigan, with the help of funds from N S F Grant GP-8345. The X-ray diffractometer (Syntex P1) was made possible by NSF Grant 18277.
they exhibit an electron paramagnetic resonance (epr) spectrum with both g,, and gX,,,< 2. A subclass of this group of proteins which contain two atoms of both iron and inorganic sulfide (e.g., spinach ferredoxin) has been extensively characterized by chemical and physical methods and a probable structure deduced. (2) W . R . Dunham, G. Palmer, R. H. Sands, and A. Bearden, Biochem. Biophys. Acta, 253, 373 (1971).
Rubinson, Palmer
Analog for Reduced 2Fe-S* Iron-Sulfur Proteins
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Simple chemical analogs of these proteins have been sought for several years,8,4in the hope of clarifying both their physical and biological properties. From the known physical data we can list a set of properties that should be exhibited by any pertinent chemical analog of the reduced 2Fe-S* proteins. The analog should have (a) two irons and no more than six ligating sulfurs,5-’ (b) a temperature-dependent magnetic susceptibility characteristic of a system with spin ‘iZ ground state and mtiferromagnetic coupling constant J = - 100 to -200 ~ m - l , (c) ~ , ~an epr spectrum with g,, < g, and g,, g, < g, < g,, the correct “sense” of the anisotropy,61 lo- l 2 (d) electron nuclear double resonance (endor) and Mossbauer spectral parameters characteristic of one high-spin ferrous and one high-spin ferric ion with hyperfine fields at the j7Fe nucleus indicative of electrons localized on covalently bound iron, 1 3 , 1 4 and (e) electronic transitions in the visible and nearinfrared region similar to those of the proteins with comparable circular dichroism properties in the nearinfrared.lj In this paper we describe the preparation and characterization of a crystalline compound which exhibits many of these properties. Experimental Procedures Preparative Method for the Model Compound. To 100 ml of n-hexane were added 7.4 g of bis(trifluoromethy1)-l,Z-dithiolene, (CF3)?C2SZ, and 2.3 g of Fe(CO);, a mole ratio of 3 :1. The mixture was stirred in a closed flask, with a valve allowing pressure release,I6 for 15-24 hr at ambient temperature (-23”). During this period the solution turned magenta with some precipitated (C0)3Fe[(S2CZ(CF&]. Then with the flask opened to air and still at room temperature, the solution was saturated with Hi3 passing through the solution for 3-5 min.” The solution was then heated, and when it began t o boil, the H S delivery tube tip was raised to a position just below the surface of the liquid. The solvent was replenished as needed. The solution turned deep blue and then gray-green within about 20 min, and the characteristic visible spectrum of the model compounds was fully developed within 1 hr.I8 (The 830-nm peak had become maximized relative to those around 600 nm after that time.) The solvent was evaporated at about 35” to one-half the original volume over 1-2 days. Some precipitate formed. The liquid which contained the desired compound was decanted into -~ -__
(3) E.g., M. Clare, H. A. 0. Hill, C. E. Johnson, and R . Richards, Chem. Commun., 1376 (1970). (4) E.g., E. Bayer, H. Eckstein, H . Hagenmaier, D. Josef, J . Koch, P. Krauss, A . Roder, and P. Schretzmann, Eur. J . Biochem., 8, 33 (1969). (5) B. B. Buchanan and D. I. Arnon, Adcan. Enzj,mol. Relat. Subj. Biochem., 33, 119 (1970). (6) J . Fee and G . Palmer, Biochem. Biophys. Acta, 245, 175 (1971). (7) T. Kimura, Struct. Bonding (Berlin), 5, 1 (1968). (8) T. Kimura, A. Tasaki, and H. Watari, J . Biol. Chem., 245, 4450 (1971). (9) G . Palmer, W. R. Dunham, J. A. Fee, R . H. Sands, T. Iizuka, and T. Yonetani, Biochem. Biophys. Acta, 245, 201 (1971). (10) G . Palmer and R. H. Sands, J . Bid. Chem., 241, 253 (1966). ( 1 1 ) G . Palmer, H. Brintzinger, and R . W. Estabrook, Biochemistrjj, 6, 1658 (1967). ( 1 2) J. C. M. Tsibris, R. I. Tsai, I. C. Gunsalus, W. H. Orme-Johnson, R . E. Hansen, and H. Beinert, Proc. Nat. Acad. Sci. I/. S., 59, 959 (1968). (13) J . Fritz, R. Anderson, J. Fee, G. Palmer, R. H. Sands, J . C. M. Tsibris, I . C. Gunsalus, W. Orme-Johnson, and H. Beinert, Biochem. Bi0phj.s. Acta, 253, 110 (1971). (14) W. R. Dunham, A. J. Bearden, I. T. Salmeen, G. Palmer, R. H. Sands, W. H. Orme-Johnson, and H. Beinert, ibid., 253, 134 (1971). ( 1 5 ) K . A. Rubinson, Ph.D. Thesis, University of Michigan, Ann Arbor, Mich., 1972. (16) This is easily effected using a sealed hypodermic needle resting in a fire-polished glass tube with a liberal coat of petroleum jelly. (17) A tube with i.d. greater than 2 mm should be used to avoid
clogging. (18) If the solution becomes purple, [Fe(SrC~Fa)nl~, more ligand may be added ; the yield will be reduced, however. Also, too much reaction time and’or H9S results in an emerald green solution which reverts to gray-green during the prolonged solvent evaporation.
Journal of the American Chemical Society
94:24
another flask, and the solvent was allowed to evaporate to dryness at -35” (another 3-4 days) whereupon the desired crystals and a tarry residue remained. The crystals were washed with hexane and then pentane unt.il all the tarry substances were removed. Yields were up to 2.6 g, 45% by iron, but were highly variable. The compound was best recrystallized from hexane or cyclohexane by dissolution at ambient temperature, evaporation to dryness, and again washing off the tarry substances. The crystals formed were air-stable, shiny black octahedra or truncated octahedra. When recrystallized from CHZCl2,the crystals tended to be trigonal prisms. Compound I may be stored indefinitely at ambient temperature. Purity may be easily checked by comparing the ab= 1.0 is sorbance ratios of the 830- and 630-nm bands. A6z8/A830 good. The best value achieved was A628/A830 = 0.9. The crystals melt with decomposition at 225-259“ in open capillaries. The material is destroyed by solvents which are Lewis bases. Isopropyl disulfide (Aldrich Chemical Co.), Halocarbon oil (Halocarbon Products Co., Hackensack, N. J.), and bis(trifluor0methy1)butyne-2 (Pierce Chemical Co.) were used as received. Iron pentacarbonyl (Alfa Inorganics) was filtered when necessary. Normal pentane, hexane, heptane, and octane were “pure grade” (Phillips Petroleum Co.) and used as received. For spectra, “spectroquality” (Matheson) reagents were used. Carbon tetrachloride and trichloroethylene (J. T. Baker Reagent Grade) were distilled from over Linde Type 4a molecular sieves prior to use for the molecular weight determination and spectra. Sulfur, sublimed (nonferrous), and various sulfides, thiols, and disulfides were obtained from various sources in this chemistry department. Bis(trifluoromethyl)-1,2-dithiolene, (CF&CISP, was prepared by the method of Krespanlg in which the fluoromethylacetylene is passed over the surface of. or even through, sulfur refluxing at about 450”. Two full-length condensers in series were necessary to condense the liquid. The material was not further purified and contained as impurities dissolved elemental sulfur and polysulfides of the fluoroethylene. Elemental analyses were performed by Spang Microanalytical Laboratory, Ann Arbor, Mich. 48106. Melting points were determined on a Gallenkamp melting point apparatus in open capillaries. Densities were determined by the flotation methodz0 using an aqueous ZnBrn solution with the commercial detergent “Alconox” (Alconox, Inc., New York, N. Y . 10003) used as a wetting agent. Using the same method one known compound, [CO(S2C4F&]*,was measured to have density 2.23 compared to a literature valuez1 of 2.28. The epr spectra were obtained on a Varian V-4500-10A spectrometer with a Varian K-band (35 GHz) bridge or a special X-band (9 GHz) microwave bridge.zz Calibration of the g values was obtained using pitch (g = 2.0029; Varian No. 904450-01) with the frequency determined in X-band by a Hewlett-Packard Model X532B frequency meter installed in the crystal bias arm of the bridge. The cryogenic temperature controller used down to 77°K is similar to that described by Hansen, et aI.Z3 For measurement down to lO”K, the apparatus described by FritzZ was used. The powder spectra line-shapes were fit with a program described by Salmeen. 25 Spin quantitation of a sample was done as previously described.** Correction was made for transition probability variation with g, 2 F here a factor of 1.14. No correction was made to account for addition of reagents such as isopropyl disulfide since in the quantitation experiments they diluted the solutions less than 4%. Such a correction would lessen the Fe/spin ratio. Endor spectra were obtained on an instrument described by Fritz.I3 I9Fcontinuous wave nmr spectra were obtained at ambient temperature on a Varian Model HA-100 spectrometer operating at 94.077 MHz locked on signal of the internal reference CFCh. The model compound sample used had been recrystallized and was dissolved in dichloromethane to saturation for the experiment. Since (19) C. G. Krespan, J . Amer. Chem. SOC.,83, 3434 (1961). (20) G. H. Stout and L. H. Jensen, “X-Ray Structure Determination,” Macmillan, New York, N. Y . , 1968. (21) J. H. Enemark and W. N. Lipscomb, Inorg. Chem., 4, 1729 (1965). (22) G . Palmer in “Methods in Enzymology,” Vol. X, R . W. Estabrook and M. E. Pullman, Ed., Academic Press, New York, N. Y . , 1967, p 602. (23) R. Hansen, T. T. I 0.01. Because the Mossbauer data show that the compound is significantly distorted, it seems that this transition is not magnetically allowed. For d-d transitions arising from the same term we expect all transitions except d,, + d,, and dZ2-y2 d22 to be magnetically allowed. Assuming basically, but not perfect octahedral symmetry around the iron, either the band is from the d,, + d,, transition or from an electric-dipole forbidden charge transfer transition. The transitions at 2.3 and 3.6 kK have half-widths that are one-half that of all the other bands as a result of smaller vibronic effects. This indicates that the transitions involve the iron d orbitals which are somewhat isolated from the bonding effects in contrast to the orbitals involved in charge transfer. Because of the low intensity, it is not possible to decide the detailed nature of these d-d transitions. Mossbauer. The Mossbauer spectrum of the crystalline material was a simple quadrupole doublet from 300 to 4.2"K. The value of the quadrupole splitting was 1.79 mmjsec at 300°K and increased slightly to 1.85 mmjsec at 4.2"K. The value of the isomer shift is 0.12 mm (us. Fe metal). The values do not change significantly when the solid is dissolved in toluene, in isopropyl disulfide, or in a mixture of the two solvents (data taken at 77 and 4.2"K). However, the linewidths of the quadrupole doublet increased by about 25% on going from the solid to a solution in toluene and by an additional 30% when the disulfide was added. The disulfide produced a small distortion of the quadrupole doublet which may be due to either hyperfine interaction or a second species of iron. In the spectrum the solid absorption decreased to zero between the absorption peaks indicating that there was no broad underlying peak. Since the full range of possible isomer shifts for 57Feis about 1.5 mmjsec, the --f
(59) W. Kuhn, Trans. Faraday SOC.,46, 293 (1950). (60) S. F. Mason, Proc. Chem. SOC.,137 (1962).
Figure 8. Far-infrared spectrum of I in Nujol (mineral oil) mull (top) and the Raman spectrum in CCl, solution (bottom). All the numbered Raman peaks were fully polarized.
two peaks cannot each be from one of two inequivalent We may ask what difference in Fe environment and oxidation states are permitted by the data. The quadrupole splitting parameter is related to the symmetry around the iron and also may contain some contributions from outside the coordination shell. The relationship between this indirect measure of distortion and actual spatial displacement is not understood. As seen in Figure 7 the absolute values of the quadrupole splitting parameter of the sulfur-coordinated compounds appear to fall into three separated ranges: < I , 1.5-1.9, and >2.3 mmjsec. The model compound falls into the middle group, the majority of which are distorted octahedrally coordinated compounds, but contains at least one compound with the irons coordinated by a distorted trigonal-bipyramidal set of sulfurs and, in addition, two full iron bonds (sample
J).63 The half-width at half-height for each component of the quadrupole pair is about 0.33 m m i ~ e cwhich ~ ~ is comparable with other spectra of the type.65,66Thus we should be able to resolve differences in quadrupole splitting of about 1 mmjsec and see some skewing of the line shape if the isomer shifts differ by more than 0.3 mmjsec. If both quantities differ, the variations should be observable below the limits set for each parameter changing independently. From the variation of parameters for the differences in "oxidation states" of isomorphic complexes,67we conclude that some skewing of the line shapes should be observable if the irons in the model compound are in different average oxidation states. However, it is unlikely that slight differences in (61) N. E. Erickson in Adcan. Chem. Ser., N o . 68,86 (1967). (62) E. Fluck, W. Kerler, and W. Neuwirth, Angew. Chem. In?. Ed. Engl., 2, 277 (1963). (63) I. Bernal, private communication. (64) The source half-width is about 0.025 mm'sec. (65) E. A Pasek and D. K. Straub, Inorg. Chem., 11, 259 (1972). (66) R. H. Herber, W. R. Kingston, and G. I gewith the anisotropy in the correct “sense.” It is the only analog for adrenodoxin and putidaredoxin which have essentially axial spectra. The chromophores for all appear to be located on an iron-sulfur site. The antiferromagnetic interaction is within the same range as the spinach protein, but is much smaller than that of adrenodoxin to which the g values seem t o be more closely related. In addition the visible and near-infrared spectra are in exceptional agreement. There is less correspondence in other aspects. In the model compound there are more than nine sulfurs coordinated to the two irons in contrast to six or seven believed for the proteins. The spin state and formal oxidation numbers of the irons probably differ from those of the proteins. Further, the circular dichroism properties and Mossbauer parameters of the model compound differ significantly from those of the proteins. The conclusion that must be reached is that even though the model compound exhibits many chemical and physical properties which match those of the proteins, fundamental differences do exist, and while the chemistry of this compound is interesting in its own right, the value as a model compound for the ironsulfur proteins is limited. Acknowledgments. For experimental assistance we thank the following: Professors B. J. Evans and E. Munck for obtaining the Mossbauer spectra; Jill Rawlings at Professor Harry Gray’s laboratory for measuring the solution magnetic susceptibility ; Professor Ian Dance at the University of Wisconsin for the solid state susceptibility measurements ; and Dr. William Eaton at the National Institutes of Health in Bethesda for measuring the near-infrared circular dichroism properties of the model compound. We thank Professors R. H. Sands and T. M. Dunn for helpful discussions.
1 November 29, 1972
