J . Phys. Chem. 1989, 93, 1362-1365
1362
Rates of Detritiation of Various Activated Carbon Acids in Water-Hydroxide and Alcohol-Alkoxide Media: A Possible Probe for Classification of Carbon Acids John R. Jones,*'+ Kevin T. Walkin,+ John P. Davey,* and Erwin Buncel*,* Department of Chemistry, University of Surrey, Guildford, England GU2 5XH, and Department of Chemistry, Queen's University, Kingston, Canada K7L 3N6 (Received: March 31, 1988)
The kinetics of detritiation of a range of carbon acids activated by different kinds of groups have been investigated in water-hydroxide, methanol-methoxide, and ethanol-ethoxide solutions. The results show that for those acids that can be characterized as "normal" in the Eigen sense (chloroform, phenylacetylene, and dimethyl sulfone) the rates are up to 10 times faster in water-hydroxide solutions than in methanol-methoxide, not as expected from the basicities of the respective solutions. However, for the acetophenones it is found that k T O ~>r kToH-and this is also the case for the cyanocarbon acids. The ratio kToEt-/kToMrvaries within much narrower limits and is always greater than unity. It is suggested that the rate ratio kToH-/kToMr may be used as an additional probe for the classification of carbon acids.
Proton-transfer processes from carbon occupy a central position in the study of chemical reactions and the development of kinetic theories. Prior to the avialability of accurate pK, values for weak carbon acids through equilibrium measurements in dimethyl sulfoxide' and cyclohexylamine,2 it was accepted practice to use kinetic acidities as a measure of their ease of ionization. Implicit in this treatment was the expectation that on moving to a more basic system the relative kinetic acidities would follow the same order, irrespective of the nature of the carbon acid. The present study was undertaken to test this hypothesis, and the results have shown that the expected behavior does not hold in all cases. It has been found that the kinetic behavior on changing the base/solvent system is in fact dependent upon the nature of the activating group, Le., on the type of carbon acid. Structural modification of carbon acids gives rise to a wide spectrum of acidities and reaction rates, from encounter-controlled to reactions having half-lives of many years.3 The degree to which carbon acids may follow the behavior of typical oxygen and nitrogen acids, which Eigen has termed normal b e h a ~ i o rhas , ~ been a challenging problem over the past 2 decades.5-8 The essential characteristics of normal proton transfer occurring in aqueous solution as summarized by Long6band more recently by othersS are as follows: (a) When the transfer process is in the thermodynamically favored direction, the rate of proton transfer will be diffusion controlled. (b) In the reverse direction, general base catalysis will be exhibited with a B r ~ n s t e dcoefficient of unity. (c) Deviations from Bransted linearity are expected in the region where the difference in pK, between substrate and base is zero. (d) The transition state of the reaction resembles an encounter complex of the acidic and basic reactants. (e) In such systems, primary hydrogen isotope effects for the proton-transfer process will be small. In recent years it has become increasingly apparent that the nature of the activating group in a carbon acid plays an important role in determining whether its behavior can be classified as normal. Cyanocarbon acids were thought to qualify as being normal: but there were indications that not all would do so, as will be considered in the Results and Discussion. The evidence for chlor0form,6~~ phenylacetylene,' and disulfonyl-activatedcarbon acid^^,'^ all favored the normal category. The activating group will largely determine whether the charge formed on the anion will be localized or not, and this is turn will govern the degree of interaction with the solvent, with the attendant possibility of internal return.* It follows, therefore, that a study of the rates of ionization of various activated carbon acids
* Address correspondence to these authors.
University of Surrey. 'Queen's University.
+
in different media, with different basicities and hydrogen-bonding properties, might shed some light on these factors. With this aim in view, the rates of detritiation of nine carbon acids of various functional types were studied in both water-hydroxide and methanol-methoxide solutions; all but one were also studied in ethanol-ethoxide solutions. The unexpected results obtained are reported herein, and their significance is discussed.
Results and Discussion The results of the rates of detritiation of the various functional carbon acids are presented in Table I. The most significant finding to emerge from the study is the fact that for three types of carbon acids (chloroform, phenylacetylene, and dimethyl sulfone) the second-order detritiation rate constants in hydroxide-water are 8-10 times faster than in methoxide-methanol, whereas for the other carbon acids the latter medium is more effective in abstracting the triton by up to a factor of ca. 25. Our results can be rationalized by considering two categories of carbon acids: those that give rise to carbanions with essentially localized charge and those in which charge is substantially delocalized through resonance. It is found that the former are characterized by the reactivity order kToH-> kToMrwhile the latter follow the order kToMZ> kTow The two kinds of reactivity (1) Matthews, M. S.; Bares, J. E.; Bartmess, J. E.; Bordwell, F. G.; Cornforth, F. J.; Drucker, G. E.; Margolin, 2.;McCallum, R. J.; McCollum, J. G.; Vanier, N. R. J. A m . Chem. SOC.1975, 97, 7006. (2) Streitwieser, A., Jr.; Juaristi, E.; Nebenzahl, L. L. Comprehensiue Carbanion Chemistry; Buncel, E., Durst, T., Eds.;Elsevier: Amsterdam, 1980 Part A. (3) (a) Bell, R. P. The Proton in Chemistry; Cornell University Press: Ithaca, NY, 1973. (b) Jones, J. R. The Ionization of Carbon Acids; Academic: London, 1973. (c) Buncel, E. Carbanions. Mechanistic and Isotopic Aspects; Elsevier: Amsterdam, 1975. (d) Reutov, 0. A,; Beletskaya, I. P.; Butin, K. P. CH-Acids;Pergamon: Oxford, 1978. (e) Stewart, R. The Proton. Applications to Organic Chemistry; Academic: Orlando, FL, 1985. (4) Eigen, M. Angew. Chem. (In?.Ed. Engl.) 1964, 3, 1. (5) (a) Kresge, A. J. Acc. Chem. Res. 1975, 8, 354. (b) Jencks, W. P. Chem. Rev. 1985,85, 51 1. (c) Bednar, R. A.; Jencks, W. P. J . Am. Chem. Sor. 1985, 107,7117. (d) Hibbert, F. Adu. Phys. Org. Chem. 1986, 22, 113. (f) Albery, W. J. Ann. Rev. Phys. Chem. 1980, 31, 227. (6) (a) Walters, E. A.; Long, F. A. J . Am. Chem. Soc. 1969, 91, 3733. (b) Hibbert, F.; Long, F. A.; Walters, E. A. J. A m . Chem. SOC.1971, 93, 2829. (c) Hibbert, F.; Long, F. A. J. Am. Chem. Soc. 1972, 94, 2647. (d) Margolin, Z.; Long, F. A. J. A m . Chem. SOC.1973, 95, 2757. (7) (a) Kresge, A. J.; Lin, A. C. J. Chem. SOC.,Chem. Commun. 1973, 761. (b) Dahlberg, D. B.; Kresge, A. J.; Lin, A. C. J . Chem. Soc., Chem. Commun. 1976, 35. (c) Lin, A. C.; Chiang, Y.;Dahlberg, D. B.; Kresge, A. J. J. A m . Chem. SOC.1983, 105, 5380. (d) Kresge, A. J.; Lin, A. C. J. Am. Chem. SOC.1975, 97, 6257. (e) Kresge, A. J.; Powell, M. F. Int. J. Chem. Kinel. 1982, 14, 19. (8) (a) Koch, H. F. Acc. Chem. Res. 1984, 17, 137. (b) Koch, H. F. Comprehensive Carbanion Chemistry; Buncel, E., Durst, T., Ed.; Elsevier: Amsterdam, 1987; Part C. (9) Hibbert, F. J. Chem. Sor., Perkin Trans. 2 1973, 1289. (10) Bell, R. P.; Cox, B. G. J . Chem. SOC.B 1971, 652.
0022-365418912093-1362%01.50/0 0 1989 American Chemical Societv
The Journal of Physical Chemistry, Vol. 93, No. 4, 1989 1363
Classification of Carbon Acids
TABLE I: Detritiation Rate Constants (kTg, M-’ s-’) for Various Carbon Acids in Water-Hydroxide, Methanol-Methoxide, and Ethanol-Ethoxide Solutions carbon acid chloroform phenylacetylene dimethyl sulfone
p-(dimethy1amino)acetophenone acetophenone p-nitroacetophenone benzyl cyanide p-nitrobenzyl cyanide 1,4-dicyanobut-2-ene
“ Reference 6d.
*Reference 7a.
T, OC
kTOH-
kTOMe-
kTOEt-
25.0 0.0 25.0 25.0 25.0 25.0 25.0
0.172 (0.165)“ 4.34 ( i 9 4 ) b 9.24 X lo4 6.7 X lo4 5.5 x 10-3 ( 5 . 4 ~i o - 3 y 3.38 X 0.215 0.017 3.24 0.191 (0.198)d 0.0124
0.0174 0.572 1.20 x 10-4 1.42 x 10-3 1.94 X 0.207 1.34 0.131 >20 3.68 0.321
0.265 1.41 3.93 x 10-4 7.71 x 10-3 0.137 2.73
0.0 0.0 25.0 0.0
kTOMe-/kTOH-
0.10 0.13 0.13 2.1 3.5 6.1 6.2 7.7 >6 19.3 26
2.16 8.7
kTO!3t-/kTOMe-
15.8 2.5 3.4 5.4 7.1 13.2 16.5 27
Reference 16. dReference 6b.
orders will be determined in part by base strengths and in part by the hydrogen-bonding capabilities of solvent toward the carbanions generated on proton abstraction. Facile hydrogen-bond formation is considered an essential feature of oxygen and nitrogen acids and is thus a necessary condition for normal carbon acids.4 A striking demonstration of these considerations is seen in the recent study of Bednar and Jencksh on proton transfer from HCN to bases of various types which showed that for ApK = 0 the reactivity was in the order 0 > N > S > C. Noting that this “parallels the decreasing electronegativity and hydrogen-bonding ability of these atoms”, it was reasoned that ”hydrogen bonding stabilizes the transition state, so that loss of this hydrogen-bonding ability increases the intrinsic barrier for proton transfer”.5c In the case of chloroform the %C13 anion will have a tetrahedral structure with the negative charge localized in an sp3 orbital, and no formal resonance structures may be drawn. Ionization of phenylacetylene will leave a localized electron pair in an sp hybrid orbital orthogonal to the a system of the carbon-carbon multiple bond. Again, in the ionization of dimethyl sulfone, charge is largely localized on carbon. Theoretical calculations have shown that there is no carbanion stabilization by (d-p)a interaction between the carbanion lone pair and sulfur 3d orbitals.” It is more likely that electrostatic effects are important in the stabilization of sulfonyl carbanions, with evidence that the tetrahedral structure of the sulfone is maintained. The other two types of carbon acids studied are the substituted acetophenones and cyanocarbons. In these systems the carbanions obtained on proton abstraction will be of the resonance-delocalized type, though some stabilization of charge by polar-field effects could also occur. The resonance delocalized carbanions are characterized by major structural and electronic reorganization. If one considers two carbon acids of the same pK,, one forming a carbanion where charge is localized and the other delocalized, it follows that hydrogen bonding between the anion and the hydroxylic solvent would be much more favored in the first case. The importance of this factor is seen further from consideration of the detailed mechanism of proton exchange. Following the work of Eigen,4 the detailed mechanism of proton exchange can be formulated as occurring via the following sequence of reactions:’*
The three-step scheme constitutes (1) proton abstraction with formation of a hydrogen-bonded carbanion ion-paired species, (2) separation and exchange of the hydrogen-bonded species with solvent, and (3) collapse of the hydrogen-bonded ion pair to exchanged products. Thus, the ability of different carbanions to partake in hydrogen bonding will have direct consequence on reactivities in proton exchange. It follows also that comparison of exchange rates in water with those in an alcoholic solvent, which differ in their hydrogen-bonding capabilities, could differentiate between the two classes of carbanions. It should be pointed out that although “normal” acid behavior is associated with diffusion-controlled proton transfer, whether in practice a carbon acid behaves “normally” depends on two factors: the magnitude of the Marcus intrinsic barrierI3 (AG*o, Le., AG* at AGO = 0) and the thermodynamic imbalance between reactants and products (AGO). Thus any highly exergonic reaction could become diffusion controlled, but ”normal” behavior will be observed for a smaller thermodynamic change when the intrinsic barrier is small. On the other hand, even proton transfer between electronegative atoms can become “nonnormal” when AGO for the reaction tends to Solvation changes will affect intrinsic barriers and thermodynamics of processes in a complex way, especially in view of the current discussions as to whether solvation changes (and delocalization) occur synchronously with proton transfer or subsequently to it.I5 For the present purpose, when changes in solvent and base are considered, we adopt a somewhat simplified approach, as follows (see later, however). (12) The scheme for roton exchange as given in q 1-3 is simplified. In a more detailed schemef eq 1 would be replaced by two separate processes shown in R-T
+B
R-T...B
3
R-T*-B
(la)
R‘ ...TB+
(1b)
to signify that a hydrogen-bonded encounter complex is formed first and that the actual transfer step occurs within the hydrogen-bonded s p i e s . Similarly eq 3 would be replaced by R-...HB* a R-H...B R-Ha-B
2
R-H
+B
(34 (3b)
An alternative scheme*that highlights the importance of internal return within the hydrogen-bonded carbanion intermediate from subsequent processes, including proton exchange, is given by R-T
(1 1) (a) Wolfe, S. Acc. Chem. Res. 1972, 5 , 102. (b) Wolfe, S.; LaJohn, L. A,; Weaver, D:F. Tetrahedron Left. 1984, 25, 2863. (c) Streitwieser, A,, Jr.; Williams, J. E., Jr. J . Am. Chem. SOC.1975, 97, 191.
+ B e R--TB+
-
products
(4)
We note also that, customarily, solvation of the various species in eq 1-3 is not shown and that hydrogen bonding with the bulk solvent would have to be considered for a deeper level of understanding of these systems. (13) Marcus, R. A. J. Phys. Chem. 1968, 72, 891. (14) (a) Cox, M. M.; Jencks, W. P.J. Am. Chem. SOC.1981, 103, 580. (b) Jencks, W. P.; Brant, S. R.; Gandler, J. R.;Fendrich, G.; Nakamura, C. J. Am. Chem. SOC.1982, 104, 7045. (15) (a) Bernasconi, C. F. Tetrahedron 1985, 41, 3219. (b) Bernasconi, C. F. Ace. Chem. Res. 1987, 20, 301.
1364 The Journal of Physical Chemistry, Vol. 93, No. 4, 1989
Jones et al.
TABLE II: Summnrv of the Exwrimental Conditions Used for the Tritintion of the Carbon Acids HTO, carbon acid (wt or vol) solvent, cm3 pL catalyst T, ‘C time, h isolation procedure 24 added to 5 cm3 H,SO, (0.1 M) and organic layer dried over chloroform (5 cm3) dioxane (1 cm3) IO NaOH ( 1 pellet) room anhydr Na,SOh 48 added to 10 cm3 distilled water, ketone removed, dried over acetophenone ( 1 cm’) dioxane (1 cm3) 10 NaOH (1 pellet) room anhydr Na2S04 16 ketone precipitated with IO cm3 H,O and extracted into 5 NaOH (1 pellet) room p-nitroacetophenone (0.35 g) dioxane (1 cm3) CHC13; organic layer separated; dried (Na2S04) 24 ketone precipitated with 10 cm’ H 2 0 and extracted into NaOH ( 1 pellet) room p-(dimethylamino)dioxane (1 cm’) 5 CHC13; organic layer separated; dried (Na,SO,) acetophenone (0.3 g) 48 added to 20 cm3 H2S04 (2 M) and organic layer dried benzyl cyanide (0.5 cm’) ethyl acetate IO Na2C03 (0.5 g) room (Na2S04),washed with H 2 0 and extracted into CHCI, ( 1 cm’) (10 cm’); organic layer separated; dried; removed with N2 gas 100 added to 10 cm’ H 2 0 and IO cm3 CHCI,; organic layer dimethyl sulfone (50 mg) dioxane (0.3 cm’) IO NaOH (1 pellet) 85 separated; dried; removed with N, gas 100 added to 10 cm3 H 2 0 and IO cm3 CHCI,; organic layer 1,4-dicyanobut-2-ene (0.2 g) ethyl acetate IO Na2C03(0.5 g) 85 separated; dried; removed with N2 gas (1 cm’) 18 added to IO cm3 H 2 0 , organic layer separated, dried over phenylacetylene (1 cm’) dioxane ( 1 cm’) IO NaOH (1 pellet) room anhydr Na2S04
For those reactions with small activation energies, the rate will be controlled by a diffusion or rotation process, and the difference in the ratio kToMe-/kToHrepresents differences in solvation and related factors on the diffusion process. On the other hand, for ratio reflects the activation-controkd reactions, the kToMe-/kToHgreater basic strength of methoxide ion in The available evidence on phenylacetylene, chloroform, and dimethyl sulfone points to these being of the “normal” carbon acid type. Phenylacetylene is a typically weak carbon acid (pKa 20), exhibiting general base catalysis with a Brernsted /? coefficient of 0.99 f 0.05 determined from a series of primary amine bases.7c A low primary isotope effect of 1.34 has been reported by Long et a].“ from comparative studies in D 2 0 and H 2 0 (see, however, ref 5d), whereas K r e ~ g observed e~~ no primary isotope effect in initial rate experiment^,^^ with hydroxide ion and I-methylimidazole as bases. These studies indicate that proton transfer is not rate determining and that internal return is significant. Chloroform (pK, 24) similarly shows general base catalysis with p = 1.12 f 0.05 and a small primary hydrogen isotope effect of 1.48.7d A positive entropy of activation of 15 eua for detritiation by hydroxide ion is probably due to desolvation of OH- with consequent release of H 2 0 molecules of solvation in the transition state of reaction. Dimethyl sulfone has not been subjected to detailed recent investigation, although its pKa has been estimated as -24.16b However, extensive studies have been carried out on disulfonylactivated substrate^,^^'^ and for these a low primary hydrogen isotope effect, k H / k D 2, and a Brernsted (3 value of 1.1 f 0.1 indicate that their behavior follows that of normal carbon acids. It seems reasonable to assume that dimethyl sulfone will similarly have the properties of a normal carbon acid. Thus the available evidence concerning phenylacetylene, chloroform, and dimethyl sulfone indicates that in terms of the Eigen mechanism for proton exchange (eq 1-3),’* step 1 is a preequilibrium and step 2 is rate determining. This is consistent with small or nonexistent primary hydrogen isotope effects and Br~nsted(3 values near unity, as well as the relatively small energies of activation. As shown by Kre~ge,’~ considerations based on Marcus rate theoryI3 and the calculation of rate profiles for chloroform and phenylacetylene are also in accord with these carbon acids falling under the normal category (see also ref 17). The ketones acetophenone, p-(dimethylamino)acetophenone, and p-nitroacetophenone are representative of carbon acids exhibiting slow proton transfer. Large primary hydrogen isotope effects have been obbserved in hydroxide-ion-catalyzed proton
abstraction of a number of acetophenones.18 In terms of the Eigen mechanism, step 1 will be rate determining in these systems. The situation with respect to the cyanocarbon acids appears to be somewhat uncertain. In the case of malononitrile (pK, = 11) and tert-butylmalononitrile (pK, = 13) the evidence6bstrongly pointed to normal behavior in the Eigen sense. For the latter, a Brernsted (3 value of 0.8 has been r e ~ 0 r t e d . l ~As well, for 21) general base catalysis with p 1,4-dicyano-2-butene (pK, values of 0.94 and 0.98 for phenoxide and secondary amine systems was observed, and the reverse rate was nearly diffusion controlled.6b However, a significant primary hydrogen isotope effect of 3.9 and an unexpectedly large negative entropy of activatiodb suggest that proton transfer occurs in the rate-determining step. p-Nitrobenzyl cyanide is characterized by a Brernsted /? value of 0.6,“ implying slow proton transfer. The cyanocarbon acids as well as the ketones are characterized by larger energies of activation. According to the results in Table 11, the koMe-/koH-values for benzyl cyanide, p-nitrobenzyl cyanide, and 1,4-dicyan0-2-butene are all significantly greater than unity, as is the case for the substituted acetophenones, but contrasting with chloroform, phenylacetylene, and dimethyl sulfone, which were considered as “normal” carbon acids. Thus, according to our analysis, the cyanocarbons in general, and 1,4-dicyan0-2-butene in particular, should not be considered “normal” in the Eigen sense. The classification of cyanocarbons together with acetophenones appears quite logical on the basis that both give rise to resonance-stabilized carbanions. The data for the kToE,-/kToMeratio (Table I) show considerably less variation than for the kToMe-/kToHratio, with all values being greater than unity. This implies that the mechanistic details are very similar for both these alcohol systems. In conclusion, the above studies have shown that even for two such similar solvent systems as hydroxide-water and methoxide-methanol, the detritiation rate ratios for various carbon acids are very sensitive to the nature of the activating group. Differences in hydrogen bonding may be partly responsible, but more quantitative information will be required since the relative basicity of O H - / H 2 0 and MeO-/MeOH could depend on the charge type of the reaction partner involved and on other properties that would affect differences in solvation. One approach to this problem could be from systematic studies of transfer energies of ground and transition states, which has been found to be of value in other systems.20 A second approach would be to focus on intrinsic barriers.13 The latter are expected to be
(16) (a) Bowden, K. Chem. Reo. 1966,68, 119. (b) Pearson, R. G.; Dillon, R. L. J . Am. Chem. SOC.1953, 75, 2439. (17) (a) Lewis, E. S.;Shen, C. C.; More O’Ferrall, R. A. J . Chem. Soc., Perkin Trans. 2 1981, 1084. (b) Kreevoy, M. M.; Konasewich, D. E. Ado. Chem. Phys. 1971, 21, 243. (c) Albery, W . J.; Kreevoy, M. M . Adu. Phys. Org. Chem. 1980, 31, 227. (d) Murdoch, J. R. J . Am. Chem. SOC.1980, 102, 71. Ibid. 1983, 105. 2159. (e) Hupe, D. J.; Wu, D. J . Am. Chem. SOC.1977, 99, 1653.
1967, 63, 1 1 1 .
-
-
-
-
(18) Jones, J. R.; Marks, R. E.; Subba Rao, S . C. Trans. Faraday
Soc
(19) Pratt, R. F.; Bruice, T. C. J . Org. Chem. 1972, 37, 3563. (20) (a) Jaiswal, D. K.; Jones, J. R.; Fuchs, R. J . Chem. SOC.,Perkin Trans. 2 1976, 102. (b) Fuchs, R.; Jones, J. R. Anal. Calorim. 1977, 4 , 22. (c) Jones, J. R.; Fuchs, R. Can. J . Chem. 1977, 55, 99. (d) Buncel, E.; Symons, E. A. J . Am. Chem. SOC.1976, 98, 656. (e) Buncel, E.; Wilson, H. Acc. Chem. Res. 1979, 12, 42. ( 0 Buncel, E.; Wilson, H. Ada Phys. Org. Chem. 1977, 14. 133
1365
J. Phys. Chem. 1989, 93, 1365-1368 greater for the proton-transfer processes leading to resonancestabilized carbanions owing to the greater structural and solvational reorganization accompanying the formation of such anions as compared to charge-localized specie^.'^^,^^ There is at present only limited information available concerning the influence of solvent on intrinsic barriers, and appropriate measurements (e.g., log k vs pK,, Bronsted type plots) in different solvents (e.g., aqueous vs methanolic) will be required before this factor can be experimentally assessed.
Experimental Section Tritiation Procedure. All the carbon acids were commercially available and were purified either by distillation or recrystallization. Tritiation was achieved by base-catalyzed exchange, and the conditions used are summarized in Table 11. The products (specific activities were generally in the range 3-100 mCi mol-I) were taken up in a deuteriated solvent (CDC13 or CD3SOCD3), and the IH and jH N M R spectra recorded. In all cases specific tritiation at the proposed site was achieved. Kinetic Measurements. The detritiation studies in aqueous media were carried out as previously reported2’ with the difference (21) (a) Davey, J. P.; Jones, J. R.; Buncel, E. Can. J . Chem. 1986,64, 1246. (b) Buncel, E.; Norris, A. R.; Elvidge, J. A.; Jones, J. R.; Walkin, K. T. J . Chem. Res. Symp. 1980, 326.
that for phenylacetylene the buffer tris(hydroxymethy1)aminomethane was employed. Dry methanol was prepared by refluxing the alcohol in the presence of magnesium turnings and iodine for 3 h before collecting the fraction that distilled at 64.3 OC. Dry ethanol was prepared similarly, the fraction boiling between 77.2 and 77.6 OC being collected. Stock solutions of sodium methoxide and ethoxide were prepared by adding freshly cut ribboned sodium metal, which has been washed with xylene and dry ether, to the respective alcohols under nitrogen in a drybox. These solutions were diluted with freshly distilled alcohol and stored in desiccators. The detritiation procedure was essentially the same as for the aqueous conditions although on a smaller scale. Usually the base concentration was varied 5-fold; quoted second-order rate constants (kTB-)are the mean of 3-5 determinations.
Acknowledgment. We are grateful to NATO, NSERC, and SERC for supporting this work. Discussions with Professors F. Hibbert, S . Hoz, and A. J. Kresge are acknowledged, as well as helpful comments from referees. Registry No. OH-, 14280-30-9; MeO-, 3315-60-4; EtO-, 16331-64-9; chloroform, 67-66-3; phenylacetylene, 536-74-3; dimethyl sulfone, 677 1-0;p-(dimethylamino)acetophenone, 2124-3 1-4; acetophenone, 98-862; pnitroacetophenone,100-19-6; benzyl cyanide, 140-29-4;pnitrobenzyl cyanide, 555-21-5; 1,4-dicyanobut-2-ene,11 19-85-3;water, 7732-18-5; methanol, 67-56-1; ethanol, 64-17-5.
Steric Constraints on Energy Deposition in the OH Product of the Reaction of O(3P,) With (CH3)SSIH Chan Ryang Park and John R. Wiesenfeld* Department of Chemistry, Cornell University, Baker Laboratory, Ithaca, New York 14853- 1301 (Received: April 25, 1988; In Final Form: August 8, 1988)
-
Characterization of the diatomic product arising from O(3PJ)+ (CH3)3SiH OH(X211) + (CH3)3Siwas accomplished by using the laser photolysis/laser-inducedfluorescencetechnique previously used in the examination of the analogous reaction involving SiH4. The observed vibrational population inversion, P(u”= I)/P(u”=O) = 2.3 & 0.2, supports a mechanism involving a direct abstraction of the hydrogen by atomic oxygen. An impulsive model of rotational excitation reveals that rotation of the nascent OH in the transition state is significantly hindered by steric interaction with neighboring methyl groups during that abstraction.
Introduction The reactivity of atomic oxygen plays a critical role in establishing the chemistry of such important systems as combustion flames and semiconductor processing Kinetic of elementary reactions involving O(3PJ) have been numerous and well-documented; comparable studies of their dynamics6-I3 are far less common. One of the earliest of the latter experimental investigations centered upon the important reaction of O(3PJ)with saturated hydrocarbons6 typified by o(3pJ) + ( C H ~ ) ~ C H o~(x21-1) ( c H ~ ) ~ c (1)
-
+
(1) Rand, R. J. J . Vac. Sci. Technol. 1979, 16, 420. (2) Boyer, P. K.; Roche, G. A,; Ritchie, W. H.; Collins, G. J. Appl. Phys. Lett. 1982, 40, 716. (3) Chen, J. Y.; Henderson, R. C.; Hall, J. T.; Peters, J. W. J . Electrochem. SOC.1984, 131, 2146. (4) Herron, J. T.; Huie. R. E. J . Phys. Chem. Re$ Data 1974, 2, 467. ( 5 ) Huie, R. E.; Herron, J. T. Prog. React. Kine?. 1975, 8, 1. (6) Andresen, P.; Luntz, A. C. J . Chem. Phys. 1980, 72, 5842. (7) Luntz, A. C.; Andresen, P. J . Chem. Phys. 1980, 72, 5851. (8) Kleinermanns, K.; Luntz, A. C. J . Chem. Phys. 1982, 77, 3533. (9) Kleinermanns, K.; Luntz, A. C. J . Chem. Phys. 1982, 77, 3774. (10) Kleinermanns, K.; Luntz, A. C. J . Chem. Phys. 1982, 77, 3537. ( 1 I ) Dutton, N. J.; Fletcher, I. W.; Whitehead, J. C. Mol. Phys. 1984, 52, 475. (12) Barry, N. J.; Fletcher, I. W.; Whitehead, J. C. J . Phys. Chem. 1986, 90. 491 1. (13) Dutton, N. J.; Fletcher, I. W.; Whitehead, J. C. J . Phys. Chem. 1985, 89, 569.
0022-3654/89/2093- 1365$01.50/0
The OH(X211) product energetics were characterized by laserinduced fluorescence (LIF). Where permitted by the reaction exothermicity, vibrational excitation of the diatomic results, an observation consistent with a mechanism in which O(’PJ) directly abstracts a hydrogen atom from the hydrocarbon substrate. The reaction of O(3PJ) with SiH, O(3PJ) + SiH4
-+
OH(X21-I)
+ SiH3
(2)
also results in a substantial inversion of the O H vibrational population distribution, with values of P(u”= l)/P(u”=O) = 3.4 f 0.4 (ref 14) and 4.2 f 0.6 (ref 15) having recently been reported. The rotational distribution of the OH(X21-I,u”=1) resulting from reaction 2 is thermal with Trot= 600 20 K, but that of the ground vibrational state displays a thermal distribution (Trot= 750 K) only below N” = 6. At higher energies to the limit of reaction exoergicity, the distribution inverts in a manner similar to that observed in OH resulting from the reaction of O(’D2) with some hydrogen-containing substrate^.'^-^' Both the strong vi-
*
(14) Park, C. R.; White, G. D.; Wiesenfeld, J. R. J . Phys. Chem. 1988, 92, 152. (15) Agrawalla, B. S.; Setser, D. W. J . Chem. Phys. 1987, 86, 5421. (16) Smith, G. K.;Butler, J. E.; Lin, M. C. Chem. Phys. Lett. 1979, 65, 115.
(17) Luntz, A. C.; Schinke, R.; Lester, Jr., W. A,; Gunthard, Hs. H. J . Chem. Phys. 1979, 70, 5908.
0 1989 American Chemical Society
