J . Phys. Chem. 1984, 88, 3115-3116
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Bond Dissociation Energies in Alcohols: Kinetic and Photochemical Evidence Regarding Ion Thermochemistry Christopher R. Moylan and John I. Braurnan* Department of Chemistry, Stanford University, Stanford, California 94305 (Received: March 29, 1984)
Gas-phase proton transfer from benzyl alcohol to fluoride ion is encounter controlled, whereas transfer from neopentyl alcohol to fluoride ion is slow. These observations suggest that the equilibrium gas-phase acidity of HF lies between those of the two alcohols, in contrast to the tabulated order. A small adjustment in the gas-phase acidities of alcohols not only makes the kinetics consistent but also resolves a long-standing discrepancy between RO-H bond energies measured by kinetic techniques and those determined by acidity/electron affinity data. It also accounts for the observed branching ratio in the IR laser photolysis of the neopentyl alcohol-fluoride ion complex.
Consistency is the hallmark of experimental science, and one of the best tests of reliability is the agreement of values obtained by different methods. In this Letter, we report some kinetic observations which we interpret in a way that resolves a longstanding discrepancy between two well-established methodologies and that also provides a simple explanation for our previously reported infrared multiple photon decomposition results. Bond dissociation energies’ in neutral molecules can be obtained by thermokinetic methods (generally involving abstraction or pyrolysis reactions) or via thermochemical cycles2 involving gas-phase acidities3 and electron affinitie~:~ AHO,,i,j(AH) = DHo(A-H)
- EA(A)
+ IP(H)
(1)
Both methodologies produce generally reliable results, and bond energies determined by the two techniques usually agree.’,5 One notable exception, however, is the set of RO-H bond energies in alcohols. Careful thermokinetic measurements on dialkyl peroxides6,’ and alkyl yield heats of formation of alkoxy radicals which, when combined with alcohol heats of formation, give average RO-H bond energies of 104 kcal/mol for a variety of alcohols. Electron photodetachment from alkoxide ions’*13 yields electron affinities which, when combined with gas-phase aciditie~,~ produce average RO-H bond energies of 102 kcal/mol. The differences are consistent between alkyl groups; changes in basicities of alkoxides correspond closely to changes in EA’S,implying that bond energies are constant. Although the uncertainties in gas-phase acidities and heats of deprotonation are assigned to be as large as 2 kcal/mol, it is clear that a systematic error exists either in the kinetic measurements, the electron affinity measurements, or the gas-phase acidity scale. This persistent discrepancy has been
(1) McMillen, D. F.; Golden, D. M. Annu. Rev. Phys. Chem. 1982, 33, 493-532. (2) Moylan, C. R.; Brauman, J. I. Annu. Rev.Phys. Chem. 1983, 34, 187-2 15. (3) Bartmess, J. E.; McIver, R. T. Jr., In “Gas Phase Ion Chemistry”; Bowers, M. T., Ed.; Academic Press: New York, 1979; Vol. 2, pp 87-121. (4)Janousek, B. K.; Brauman, J. I. In ref 3, pp 53-86. (5) Brauman, J. I. In “Frontiers of Free Radical Chemistry”; Pryor, W. A., Ed.; Academic Press: New York, 1980; pp 23-30. (6) Benson, S. W.; Shaw, R. Ado. Chem. Ser. 1968, No. 75, 288-94. (7) Batt, L.; McCulloch, R. D. Int. J. Chem. Kinet. 1976, 8, 491-500. (8) Batt, L.; Milne, R. T. Int. J . Chem. Kinet. 1977, 9, 549-65. (9) Batt, L.; Islam, T. S. A,; Rattray, G. N. Int. J . Chem. Kinet.1978.10, 931-43. (10) Reed, K. J.; Brauman, J. I. J . Am. Chem. SOC.1975, 97, 1625-6. (11) Janousek, B. K.; Zimmerman, A. H.; Reed, K. J.; Brauman, J. I. J. Am. Chem. SOC.1978, 100, 6142-8. (12) Engelking, P. C.; Ellison, G. B.; Lineberger, W. C. J . Chem. Phys. 1978, 69, 1826-32. (13) Ellison, G. B.; Engelking, P. C.; Lineberger, W. C. J . Phys. Chem. 1982,86, 4873-8.
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TABLE I: Proposed Heats of Deprotonation (AHoscid) for Aliphatic Alcohols Less Acidic Than HF AHoacid, kcal/mol present value revised value (*2.0)“
CHSOH EtOH n-PrOH
i-PrOH i-BuOH t-BuOH t-BuCHZOH a
379.2 376.1 374.7 374.1 373.4 373.3 371.8
(f0.6) ‘
381.4 378.3 376.9 376.3 375.6 375.5 374.0
Reference 3.
previously remarked upon,’,’ although no explanation for it has been advanced. Gas-phase acidities are known absolutely by experimental reference to independently known values. The anchor point for alcohols is H F via neopentyl alcoh01~J~ (eq 2, Keq = 2). Using this value, we were able to explain the branching ratio for the infrared multiple-photon-induced decomposition of (CH,)$CH,OHF- (the likely intermediate in eq 2) only with diffi~u1ty.l~ In order to understand the magnitude of this branching ratio, we have measured the rate constants of two proton transfer reactions:
-
+ (CH,),CCH,OF- + C ~ H S C H ~ O HHF + C ~ H S C H ~ O -
F + (CH,),CCH,OH
HF
(2)
(3) Reaction 2 has a rate constant of 2.9 X cm3 molecule-’ s-I, corresponding to a reaction efficiency of onlyI6 12% based on an cm3 molecule-’ s-’ ADO collision rate constantI7 of 2.50 X (in reasonable agreement with an earlier valuelo of 2.2 x 1O-Io). Reaction of F with benzyl alcohol (eq 3), however, is much faster, exceeding the estimated” collision rate constant of 2.54 X The slow rate of reaction 2 and the surprisingly large branching fraction for F in the IR photolysis of the proton bound complex are consistent with each other because the proton transfer reaction presumably proceeds through the complex. In each case, more fluoride and less alkoxide is produced upon decomposition of the complex than one would expect based on the available thermochemical data, taking into consideration that eq 2 is slightly endothermic, with a positive ASo. Proton transfers are among the fastest reactions known, and they tend to be unit efficient unless a significant thermochemical -+
(14) McIver, R. T., Jr.; Miller, J. S. J . Am. Chem. SOC.1974, 96, 4323-5. (15) Moylan, C. R.; Jasinski, J. M.; Brauman, J. I. Chem. Phys. Lett. 1983. 98. 1-5. (16) We have previously r e p ~ r t e d ’a~preliminary value of 8%. (17) Su, T.; Bowers, M. T. In “Gas Phase Ion Chemistry”; Bowers, M. T., Ed.; Academic Press: New York, 1979; Vol. 1, pp 83-118.
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J . Phys. Chem. 1984, 88, 3176-3179
barrier exists or unless they are not exothermic. Proton transfers between a particular ion and a series of neutrals generally exhibit efficiencies close to unity until the reactions become endothermic, at which point the reactions become very slow or do not proceed at all. Such kinetic observations have been used to determine relative acidities.'8-21 Our kinetic data strongly suggest that the equilibrium gas-phase acidity of HF lies between neopentyl alcohol and benzyl alcohol. In our previous discussion of photochemical branching fractions,l5 we noted that the slow kinetics of reaction 1 could be explained by a barrier, and the size of the branching fractions for the proton-bound complex by an energy difference between the two decomposition channels greater than that indicated by the thermochemistry; the latter could be caused by rotational excitation of the complex. A simpler explanation now suggests itself the thermochemistry is in error by about 2 kcal/mol, which explains not only the branching fraction and kinetics, but also the long(18) Bohme, D. K.; Mackay, G. I.; Tanner, s.D. J . Am. Chem. SOC.1979,
101, 3724-30.
(19) Tanner, S. D.; Mackay, G. I.; Bohme, D. K. Can. J. Chem. 1981,59, 1615-21. (20) Bohme, D. K.;Rakshit, A. B. Mackay, G. I. J. Am. Chem. SOC.1982, 104, 1100-1. (21) Meot-Ner, M. J . Am. Chem. SOC.1982, 104, 5-10.
standing difference between RO-H bond energies as determined by kinetic and spectroscopic methods. The gas-phase acidity scale has been developed from many equilibrium measurements and anchored by as many trustworthy absolute values as possible. It seems likely that the 2 kcal/mol discrepancy is due to the difficulty in studying reactions of HF; therefore, only the neopentyl alcohol-HF equilibrium constant is suspect. We propose that the present set of relative acidities of aliphatic alcohols be used in conjunction with the absolute acidity of methanol, as determined by kinetic7 and spectroscopic12 measurements, to form a new set of alcohol acidities. These values are listed in Table I. Methanol is chosen as the anchor point because its 0-H bond energy (104.0 0.2 kcal/mol) and alkoxide electron affinity (36.2 f 0.5 kcal/mol) are both known to excellent precision. The resulting heat of deprotonation is 381.4 f 0.6 kcal/mol. We feel that the proposed changes will produce values that better describe the behavior of these alcohols and will serve as a more reliable anchor for the lower end of the acidity scale. Thus, it is likely that other acidities which lie near H F will have to be revised as well.
*
Acknowledgment. We are grateful to the National Science Foundatior, for support of this work and for a graduate fellowship (C.R.M.). We also thank the donors of the Petroleum Research Fund, administered by the American Chemical Society.
New Phases and Chemical Reactions in Solid CO under Pressure Allen 1. Katz: David Schiferl, and Robert L. Mills* University of California, Los Alamos National Laboratory, Los Alamos, New Mexico 87545 (Received: April 9, 1984)
The Raman scattering of solid carbon monoxide was studied in a diamond cell from 15 to 297 K at pressures from 1.0 to 5.8 GPa. At low temperature a transition occurs near 3.4 GPa from the known a-phase (space group P&3) to a new e-phase (structure unknown), rather than to the predicted y-phase (P42/mnm).A transformation from /3-CO (space group P63/mmc) into a new &phase (Pm3n?) was found near 5.2 GPa at room temperature. Above about 4.6 GPa and 80 K, CO reacts photochemically when irradiated with visible laser light. The photoreactivity may be associated with the formation of a yellow polymer, which can be recovered at zero pressure.
Introduction It is well-known that the condensed phases of carbon monoxide and nitrogen exhibit striking similarities. The molecules CO and N2are isoelectronic, have the same size and weight, and condense into liquids with almost identical molar volumes. Liquid CO and N, freeze at triple points of 68.15 and 63.14 K, respectively, forming a disordered hexagonal solid, the 0-phase, which on further cooling transforms into a molecularly ordered cubic structure, the a-phase. The @-a crystallographic transition occurs at 61.6 K for C O and 35.6 K for N,, and is driven primarily by electric quadrupole-quadrupole interactions. The structure of P-N2at low pressure has been determined',, by single-crystal X-ray diffraction to have space group P63/mmc, with the molecules tilted from the hexagonal c axis by about 56' and either precessing or statically disordered. The nature of the disorder in hexagonal 0-CO at low pressure is ~ n k n o w n . ~Recently, however, Cromer et al.435showed by X-ray measurements in diamond cells that both CO and N, freeze at room temperature near 2.5 GPa into P-solids with almost the same molecular disorder and molar volume. The molecults in these high-pressure crystals are statically disordered and tipped from the hexagonal c axis by 54O for N, and 49O for CO. 'Associated Western Universities appointee during 1984. Present address: Arizona State University, Tempe, AZ.
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In the a-phase, it is now generally accepted that the homonuclear molecules of N2 occupy centrosymmetric positions6 in space group Pa3, while heteronuclear CO, lacking inversion symmetry, packs7 as P213, with the molecules only very slightly shifted* from the Pa3 arrangement. In a-CO the molecules exhibit a large amount of end-for-end di~order,~ which becomes effectively frozen inlo well above 5 K, where transition to the ordered state would be expected." For this reason, the Raman spectrum of a-CO shows marked differences7 from that of a - N z . (1) Streib, W. E.; Jordan, T. H.; Lipscomb, W. N. J . Chem. Phys. 1962, 37, 2962. (2) Jordan, T. H.; Smith, H. W.; Streib, W. E.; Lipscomb, W. N. J. Chem. Phys. 1964, 41, 756. ( 3 ) Vegard, L. Z . Phys. 1934,88,235. (4) Cromer, D. T.; Schiferl, D.; LeSar, R.; Mills, R. L. Acta Crystallogr., Sect. C 1983, 39, 1146. ( 5 ) Schiferl, D.; Cromer, D. T.; Ryan, R. R.; Larson, A. C.; LeSar, R.; Mills, R. L. Acta Crystallogr., Sect. C 1983, 39, 1151. (6) Venables, J. A.; English, C. A. Acta Crystallogr., Sect. E 1974,30, 929. (7) Anderson, A.; Sun, T. S.; Donkersloot, M. C. A. Can. J . Phys. 1970, 48, 2255. (8) Hall, B. 0.; James, H. M. Phys. Reu. E 1976, 13, 3590. (9) Clayton, J. 0.; Giauque, W. F. J . Am. Chem. SOC.1932, 54, 2610. (10) Walton, J.; Brookeman, J.; Rigamonti, A. Phys. Rev. E 1983, 28, 4050. (11) Melhuish, M. W.; Scott, R. L. J . Phys. Chem. 1964, 68, 2301.
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