J. Phys. Chem. 1882, 86, 321-322
6.6 A formed by hydrogen-bonded water molecule^.^ We attempted to observe Raman spectra of cyclopentane in aqueous solution. Even though the solubility of cyclopentane in water is about three times higher than that of cyclohexane, the perpendicular component spectrum of cyclopentane could not be observed because of its low Raman line intensity. The largest free diameter of a hydrate cage ever reported is 6.6 A as in the above cyclopentane hydrate; a hydrate cage with a larger diameter is not known at present. The largest diameter of a cyclohexane molecule is about 6.9 A as in benzene, and benzene is the largest molecule which
321
forms a gas hydrate. Recently, we succeeded in observing the Raman spectrum of benzene in aqueous solution and found as well that the rotational motion of benzene molecules in aqueous solution is hindered more than in other solvents.lo These facts suggest that cyclohexane or benzene molecules dissolved in water are packed more closely in iceberg structures similar to the hydrate cages than in other solvents, and that the close packing may hinder the rotational motion of solute molecules in aqueous solution. (10) K. Tanabe, Spectrochim. Acta, in press.
Heat of Formation of Hydrogen Isocyanide by Ion Cyclotron Double Resonance Spectroscopy Chln-Fong Pau and Warren J. Hehre’ Department of Chemistry, University of California, Iwine, California 92717 (Recelvd: November 3, 1981)
The threshold for deuteron abstraction from protonated DCN has been determined by pulsed ion cyclotron double resonance spectroscopy to be 14.8 i 2 kcal mol-’ higher in enthalpy than the corresponding threshold for proton abstraction (i.e., proton affinity of HCN). Ignoring small effects arising from differences in isotopic substitution, this difference is a direct measure of the relative thermochemical stabilities of hydrogen cyanide and hydrogen isocyanide.
The hydrogen isocyanide molecule has been the subject of considerable attention both from experimentalists and theorists alike. First observed in the laboratory as a product of photolysis of CH3N3in an argon matrix at 4 K,’ the molecule was later suggested and then confirmed2 to be present in interstellar space. Precise knowledge of the J = 1 0 rotational line from radioastronomy prompted attempts to detect and to further characterize the species by microwave spectroscopy3 and by high-resolution infrared spectrometrp in the laboratory. Although the equilibrium geometrical structure of HNC has now been firmly established by a combination of experimental methods: the thermochemical stability of the species, relative to its more stable isomer, hydrogen cyanide, is less certain. An early unsuccessful search for the J = 1 0 microwave absorption of HNC in a normal sample of HCN at room temperature led Brown and his co-workers6 to
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(1) (a) D. E. Magan and M. E. Jacox, J.Chem. Phys., 39,712 (1963); for later work see (b) D. E. Milligan and M. E. Jawx, ibid., 47,278 (1967). (2) (a) L. E. Snyder and D. Buhl, Bull. Am. Acad. Sci., 3,388 (1971); (b) Ann. N.Y.Acad. Sci., 194,17 (1972); (c) Bull. Am. Acad. Sci., 4,227 (1972); (d) B. Zuckerman, M. Morris, P. Palmer, and B. E. Turner, Astrophys. J. Lett., 173, L125 (1972); (e) M. Morris, B. Zuckerman, B. E. Turner, and P. Palmer, ibid., 192, L27 (1974); (f) L. E. Synder and J. M. Hollis, ibid., 204, L139 (1976); (g) R. L. Snell and H. A. Wootten, ibid., 216, L l l l (1977); (h) P. D. Godfrey, R. D. Brown, H. I. Gunn, G. L. Blackman, and J. W. V. Storey, Mon.Not. R. Astron. SOC.,186, (1977); (i) R. D. Brown, Nature (London),270,39 (1977); 6)R. L. Snell and A. Wootten, Astrophys. J., 228, 748 (1979). (3) (a) R. J. Saykally, P. G. Szanto, T. G. Anderson, and R. C. Woods, Astrophys. J.Lett., 204, L143 (1976); (b) G. L. Blackman, R. D. Brown, P. D. Godfrey, and H. I. Gunn, Nature (London),261,395 (1976); (c) R. D. Brown, P. D. Godfrey, J. W. V. Storey, and F. D. Clark, ibid., 262,672 (1976); (d) R. A. Creswell, E. F. Pearson, M. Winnewisser, and G. Winnewsser, 2.Naturwissenschaften A , 31,221 (1976); (e) E. F. Pearson, R. A. Creswell, M. Winnewisser, and G. Winnewisser, ibid., 31, 1394 (1976). (4) (a) C. A. Arrington and E. A. Ogryzlo, J. Chem. Phys., 63, 3670 (1976); (b) A. G. Maki and R. L. Sams,to be published. (5) R. A. Creswell and A. G. Robiette, Mol. Phys., 36, 869 (1978). 0022-3654/82/2086-0321$01.25/0
conclude that the difference in the thermochemical stabilities of the two molecules was no less than 10.8 kcal mol-’. More recently Maki and Sams’ successfully recorded the infrared spectrum of HNC in “pure” HCN at 1000 K. Based on analysis of their intensity data, these authors estimated the relative energies of the two forms as 10.3 kcal mol-’. Work by Ellison and co-workers8suggests that HNC lies somewhere between 17.2 and 26.3 kcal mol-’ above HCN in enthalpy based on the fact that while reaction of CN- with HI leads both to HNC and HCN, proton abstraction either from HBr or from HC1 results only in the production of the more stable isomer. Both of these experimental determinations are in serious disagreement with all high level quantum mechanical calculations which have been performed on the HCN/HNC system to date. In particular, work by Pearson, Schaefer, and Walgrengusing the extensive configuration interaction expansion yielded an energy separation of 14.6 kcal mol-’; efforts by Redman, Purvis,and Bartlettg and by Krishnan and Poplegusing many-body (Moller-Plesset)perturbation theory through fourth-order suggested differences of 15 f 2 and 15.8 kcal mol-’, respectively. In view of the importance attached to the role of HNC as a active participant in the chemistry of interstellar (6) G. L. Blackman, R. D. Brown, P. D. Godfrey, and H. I. Gunn, Chem. Phys. Lett., 34, 241 (1975). (7) A. Maki and R. Sams, J . Chem. Phys., 75,4178 (1981). (8) M. M. Maricq, M. A. Smith, S. J. S. M. Simpson, and G. B. Ellison, J . Chem. Phys., 74, 6154 (1981). (9) (a) P. K. Pearson, H. F. Schaefer, 111, and U. Wahlgren, J. Chem. Phys., 62, 350 (1975); (b) L. T. Redman, G. D. Purvis, 111, and R. J. Bartlett. ibid.. 72, 986 (1980): IC) R. Krishnan and J. A. Pode. to be published. Other theoreticd work includes (d) P. K. Pearion,’ G. L. Blackman, H. F. Schaefer, 111, B. Roos, and U. Wahlgren, Astrophys. J. Lett., 184, L19 (1973); (e) P. Botschwina, E. Nachbaur, and B. M. Rode, Chem. Phys. Lett., 41,486 (1976); (0 P. R. Taylor, G. B. Bacskay, N. S. Hush, and A. C. Hurley, J. Chem. Phys., 69, 1971 (1978).
0 1982 American Chemical Society
322 The Journal of Physical Chemistry, Vol. 86, No. 3, 1982
space, and the uncertainty in its fundamental thermochemical stability, we felt it a suitable choice for investigation using pulsed ion cyclotron double resonance spectroscopy.lOJ1 The basis of the experimental method is to measure the relative free energies of deprotonation and of dedeuteration of protonated deuterium cyanide in the gas phase. This yields directly the relative thermochemical stabilities of HCN and HNC without reference to any other data. The predominant ion-molecule reactions which occur when a mixture of methyl bromide, deuterium cyanide, and some base B of known proton affinity (in approximate ratio torr) are introduced into 501:l and total pressure 3 X an ion cyclotron resonance spectrometer are :ri3E;r
e CH5Brk-
EH-
Electron impact on CH3Br eventually leads to a buildup of protonated compound by way of reaction of initially formed fragment ions with methyl bromide itself. This in turn reacts exothermically with the deuterium cyanide to yield DCNH+ and with B to yield BH+. If B is a sufficiently strong base it will be able to abstract the nitrogen bound proton in DCNH+, giving rise to an additional source of BH+. If stronger still it will be capable of deuteron abstraction (from carbon), leading to the formation of BD+, and concurrently to production of hydrogen isocyanide. Therefore, by using a series of abstracting bases of known and increasing strength, and by monitoring the onset of production of BD+, the enthalpy of (carbon) dedeuteration of DCNH+ may be determined. Relating to the known enthalpy for (nitrogen) deprotonation, and ignoring any small effects due to isotopic substitution, one finds that this corresponds precisely to the difference in the heats of formation of HCN and HNC. Propionaldehyde (enthalpy of proton transfer 14.4 kcal mol-' greater than that of hydrogen cyanide)12was the strongest base considered for which carbon dedeuteration was not ~~
Letters
observed. The weaker bases tested, methyl formate (14.1 kcal mol-', above HCN), acetonitrile (13.2 kcal mol-l), methanol (8.2 kcal mol-'), and hydrogen cyanide itself, did not lead to deuterium abstraction. Ethane thiol (enthalpy of proton transfer 15.2 kcal mol-' above hydrogen cyanide) was the weakest base considered which did result in dedeuteration as evidenced by the production of an ion of mass corresponding to the molecular formula CH3CH2SHD+. All stronger based examined, dimethyl ether (16.7 kcal mol-' above HCN), ethyl formate (17.8 kcal mol-'), and acetone (21.1 kcal mol-'), also resulted in deuterium abstraction from DCNH+. Deuterium incorporation due to reaction with fragment ions (i.e., those from CH3Br) is precluded by double resonance experiments.1°J3 Specifically the intensity of the ion of mass corresponding to BD+ (e.g., CH3CH2SHD+)was observed to decrease in response to ejection of DCNH+ from the system. For every base B, the total pressure of the reactant mixture was increased (by increasing the amount of CH3Br)until the observed double resonance spectrum was no longer altered. This indicates that the reactant ions have undergone sufficient collisions and are vibrationally relaxed. We conclude that the enthalpy of carbon dedeuteration from DCNH+ is between 14.4 and 15.2 kcal mol-' higher than that for deprotonation from nitrogen, and assign the actual threshold14to the average of these two values, 14.8 f 2 kcal mol-'. If we ignore any small efforts due to isotopic substitution, this corresponds to the difference in thermochemical stabilities between HCN and HNC. This value agrees quite closely with estimates obtained from the most complete quantum chemical calculations (14.6,9a15 f 2,9b 15.8 kcal mol-' 9 c ) , but is in serious disaccord with the recent work both of Maki and Sams7 and of Ellison and co-workers.8 Further efforts are required in order to clarify the situation. Acknowledgment. We thank Professor J. A. Pople (Carnegie-MellonUniversity), R. C. Woods (University of Wisconsin), and Dr. A. G. Maki (National Bureau of Standards) for useful discussions and for allowing us access to their work prior to publication. We are also grateful to an anonymous reviewer for bringing ref 8 to our attention.
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(10) (a) R. T. McIver, Jr., Reo. Sci. Instrum., 41,555 (1970); (b) J. D. Baldeschwieler and S. S. Woodgate, Acc. Chem. Res., 4,114 (1971); (c) R. T. McIver, Jr., and R. C. Dunbar, Int. J.Mass. Spectrum. Ion Phys., 7, 471 (1971); (d) R. T. McIver, Jr., Reu. Sci. Instrum., 49, 111 (1978). (11) For recent examples of the use of similar techniques in the elucidation of the thermochemical stabilties of short-lived neutral molecules, see (a) free radicals, D. J. DeFrees, R. T. McIver, Jr., and W. J. Hehre, J. Am. Chem. SOC.,102,3334 /1980); (b) enol of acetone, S. K. Pollack and W. J. Hehre, ibid., 99,4845 (1977); (c) methyleneimine, D. J. DeFreea and W. J. Hehre, J. Phys. Chem., 82, 391 (1972); (d) 1,l-dimethylsilaethylene, W. J. Pietro, S. K. Pollack, and W. J. Hehre, J. Am. Chem. Soc., 101, 7126 (1979); (e) o-benzyne, S. K. Pollack and W. J. Hehre, Tetrahedron Lett., 2483 (1980).
(12) Proton affinities from J. F. Wolf, R. H. Staley, I. Koppel, M. Taagepera, R. T. McIver, Jr., J. L. Beauchamp, and R. W. Taft, J. Am. Chem. SOC.,99,5417 (1977),with slight modifications to account for the higher ambient temperature of the ion cyclotron resonance Spectrometer than previously believed. (13) D. J. DeFrees, W. J. Hehre, R. T. McIver, Jr., and D. H. McDaniel, J. Phys. Chem., 83, 232 (1979). (14) The principle assumption made here is that thermoneutral or exothermic proton (deuteron) transfer proceases will be observed and that endothermic reactions will not. It is likely, however, that slightly endothermic reactions will occur to sufficient extent as to be detected. We suspect that the quoted 2 kcal mol-' error bound is large enough to account for any uncertainty in the established transfer threshold.
