J . Phys. Chem. 1986, 90, 383-389 surements the authors hoped to be able to recognize structures in the spectrum and identify areas of longer residence time of the complex on the ground-state surface from the intensity maxima. In addition to laser-excited transition states, emission from chemically excited intermediate complexes has also been observed. Polyani et a1.80,8’were able to detect fluorescence from the reaction F Na2 FNaNa** N a F Na*(32P) far from the N a D line. This “wing” fluorescence originates from Na* in the vicinity of N a F (and hence also, for example, from FNaNa+*). The higher intensity of the long wavelength “wing” is due to the fact that the upper and lower potential surfaces tend to lie closer together when passing through the intermediate configurations relative to the energy product difference. Although considerable efforts were made, no further structures or resonances were found in these experiments. The actual information on the transition states involved is thus very limited. Structured transition state spectra were observed by Kinsey et aLsz after laser photodissociation of ozone at 266 nm. Ozone dissociates within femtoseconds from the upper electronic state of the excited ”Hartley” bands to ground-state O2and O(’D).83 The dissociation is approximately six orders of magnitude faster than the spontaneous emission; nevertheless, fluorescence could be observed. From the frequency shift of the fluorescence lines relative to the excitation wavelength, one obtains vibrational
+
-
-
+
(80) P. Arrowsmith, P. E. Bartoszek, S. H. P. Bly, T. Carrington, P. E. Charters and J. C. Polyani, J. Chem. Phys., 73, 5895 (1980). (81) P. Arrowsmith, S. H. P. Bly, P. E. Charters, and J. C. Polyani, J. Chem. Phys., 79, 283 (1983). (82) D. G. Imre, J. L. Kinsey, R. W. Field, and D. H. Katayama, J. Phys. Chem., 86, 2564 (1982). (83) C.-K. Man and R. C. Estler, J. Chem. Phys., 75, 2779 (1981).
383
frequencies of the ozone in the electronic ground state up to 500 cm-’ below the dissociation limit. The relative line intensities characterize the upper potential surface and the time course of the photodissociation. In a more recent piece of work on the photolysis of CH31, Kinsey et aLs4 were able to show that the C-I bond has to be considerably extended before the CH, umbrella vibration is excited and the methyl group opens from the tetrahedral geometry in CH31 to the planar form in the CH3 product radical. The emission spectrum consists of a long progression in the C-I-stretching vibration v3 up to the 29th overtone. Toward longer wavelengths (later times), combination bands between u3 and the CH, umbrella vibration v2 appear, Clearly, the initial movement of the atoms occurs almost exactly along the C-Istretching vibration coordinate, whereas the CH3group opens later to the planar form. At very long wavelengths (high vibrational energies), the bands broaden as a result of the increasing vibrational state density. This picture of the dissociation is in agreement with calculations on the potential surface of the electronically excited CH31. Such structured emission spectra of dissociating molecules appear to be extremely promising for obtaining direct and detailed information on transition states of bimolecular chemical reactions. Acknowledgment. The financial support of the Deutsche Forschungsgemeinschaft is gratefully acknowledged. Registry No. HO, 3352-57-6; H, 12385-13-6; C 0 2 , 124-38-9; H 2 0 , 7732-18-5. (84) D. G. Imre, J. L. Kinsey, A. Sinha, and J. Krenos, J. Phys. Chem., 88, 3956 (1984). ( 8 5 ) T. Dreier, private communication.
Thermodynamics and Kinetics of Carbon-Carbon Bond Formation and Heterolysis through Reactions of Carbocatlons with Carbanions in Solution Edward M. Arnett* and Kent Molter Department of Chemistry, Duke University, Durham, North Carolina 27706 (Received: July 22, 1985)
Rate data are presented for the heterolysis of carbon-carbon bonds and their formation through coordination of resonance-stabilized carbocations and carbanions in acetonitrile solution at 25 OC. These rates were determined by NMR line broadening and by the T jump technique (in a solution containing 0.48 M supporting electrolyte). Preliminary results are given for a “master equation” to predict some heterolysis energies in solution as a complement to Benson’s method for homolysis energies. The results provide for the first time an opportunity to compare the effects of structure variation on the kinetic and thermodynamic properties for such an ostensibly simple reaction in solution. All evidence so far accumulated indicates that these reactions are dominated by ion-solvation factors so that they have little bearing on the gas-phase heterolysis energies. Ionic strength effects and substituent variation suggest that charge development is about half developed at the transition state, but we argue that this cannot be translated simply into pictures of transition-state structure. The results provide a flagrant reversal of the frequently invoked “reactivity selectivity principle” since the most reactive cation is also most selective. On the basis of these results and many others which have appeared recently it may be appropriate to discard the reactivity selectivity principle as a useful principle for either prediction or interpretation.
Introduction Reaction rates, equilibrium constants, and heats of reaction are perhaps the most fundamental physical observables which manifest reactivity and chemical driving force. The desire of chemists to find relationships between these properties and molecular structure is perennial and has provided much of the motivation for theories of chemical kinetics, especially transition-state theory. One of Sidney Benson’s most valuable contributions has been the development of a predictive scheme for relating homolytic bond energies and free radical stabilities through through the large empirical data bases of heats, entropies and free energies of formation, bond lengths, and bond angles.’-4 In view of the 0022-3654/86/2090-0383$01.50/0
relatively small solvation energies of molecules and free radicals, homolytic bond energies calculated by Benson’s methods, by molecular mechanics, or by quantum mechanics should apply within 1 or 2 kcal/mol to reactions in solution. In fact, the terms solvation or solvation energy are not to be found in the table of contents or index of ref 1. (1) Benson, S. W. “ThermochemicalKinetics”, 2nd ed.;Wiley-Interscience: New York, 1976. (2) Benson, S. W.; Buss, J. H. J . Chem. Phys. 1957, 29, 546. (3) Benson, S. W. ‘The Foundations of Chemical Kinetics”; McGraw-Hill: New York, 1960. (4) Benson, S. W. J. Chim. Phys. 1979, 76, 191
0 1986 American Chemical Society
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The Journal of Physical Chemistry, Vol. 90, No. 3, 1986
In contrast, this paper will review our recent thermodynamic studies of the direct reactions between resonance-stabilized carbocations and carbanions to produce covalent products in solution. The heats of reaction of such coordination reactions yield the heterolytic bond energies for the covalent bonds formed merely by changing the sign (AH,,, = -AHcmrd).Because of the large solvation energies of ions, the heterolytic bond energies determined in this way probably differ by hundreds of kcal/mol from the gas-phase values, none of which are presently known. However, the condensed-phase measurements are of more direct relevance to many important processes and turn out to be surprisingly easy to predict from readily available ionization constants. We shall also present our first results for the rates of several coordination/heterolysisreactions whose thermodynamic properties have been p ~ b l i s h e d . ~These data permit for the first time a comparison of the kinetics and thermodynamics for the formation of a carbon-carbon bond through coordination of its component Lewis acid carbocation (electrophile) with the carbanion (nucleophile) in solution with no leaving group, except electrostatically bound solvent molecules. Cation-Anion Reactions. There is ample precedent for kinetic studies of the reactions of resonance-stabilized carbocations with inorganic aniom6-I6 In contrast our work5v’7-20has used carbanions as nucleophiles so that the structures of both the cation and anion may be varied widely in a systematic way in order to examine the effect of such structural changes on the energetics of carbon-carbon bonds in solution. After a preliminary survey of the rates and equilibria of many carbocation-carbanion comb i n a t i o n ~ , ~ ’the - ~ ~reactions of trimethyl- and triphenylcyclopropenium ions with a series of substituted phenylmalonitrile anions were chosen for study in d e ~ t h . ~ . ’ ~
Arnett and Molter
p-methoxyphenylmalononitrileanions confirmed N M R evidence that a true covalent product results from the reaction rather than an ion pair21or charge-transfer2*complex. Good linear correlations were established between the standard free energies of heterolysis (AGOhCl),the enthalpies of heterolysis ( M ’ h e t ) , and the oxidation potentials of the carbanions. Of particular importance were good h e a r correlations between AGOhet, or AHohet,and the pKa’s of the p-substituted phenylmalononitriles. This demonstrates that the Bronsted basicities of the corresponding carbanions toward the proton in solution were closely proportional to their Lewis basicities toward cyclopropenium carbocations. Thus, for these carbanions the effects of structural change on the energetics of single electron transfer, proton transfer, and carbocation transfer are proportional and nearly identical. This is the first example known to use where the relationship between these fundamental reaction processes has been compared by direct experiment! A Master Equation for Heterolytic Bond Energies in Solution? Having shown that there is a direct proportionality between w h e t and the pK, of the carbanion’s conjugate acid it is reasonable to ask whether a comparable measure of carbocation stability may be employed as a contributor to Affoh,t. Such data are available for the ionization of many triarylcarbinols which give resonance-stabilized carbocations in aqueous acid solution. R O H 2H+ F= R+ + H,O+ (2)
+
and are expressed as pKR+for eq 2. If the heterolytic bond energy for carbon-carbon bonds to form resonance-stabilized ions (as in eq 1) were dominated by the stabilities of the ions as measured by the pK,’s of the carbanions and the PKR+’Sof the carbocations, a very simple master equation for heterolysis in solution follows (3)
which reduces to linear form
For several such combinations equilibrium constants could be determined directly by UV-vis spectrophotometry. Enthalpies of reaction obtained by the temperature coefficient of K through the van’t Hoff equation agreed within experimental error with the heat of coordination (AH-) determined calorimetrically from mixing solutions of dissociated salts of the carbocations and carbanions in acetonitrile. Variation of the equilibrium constant in response to changing the solvent dielectric constant followed the Born equation closely in nonbasic, non-hydrogen bonding media. Complete X-ray crystal structure determinations of the products from trimethylcyclopropenium cation with p-nitro- and
(5) Arnett, E. M.; Troughton, E. B.; Molter, K. E. J . Am. Chem. Soc. 1984, 106, 6726.
( 6 ) Ritchie, C. D. J . Am. Chem. SOC.1983, 105, 7313. (7) Ritchie, C. D.; Kubisty, C.; Ting, G . Y . J . Am. Chem. SOC.1983, 105, 279. (8) Ritchie, C. D. J. Am. Chem. SOC.1983, 105, 3573. (9) Ritchie, C. D.; VanVerth, J. E.; Virtanen, P. 0. I. J . Am. Chem. SOC. 1982, 104, 3491. (10) Ritchie, C. D. Pure Appl. Chern. 1978, 50, 1281. (11) Ritchie, C. D. Acc. Chem. Res. 1972, 5 , 348. (12) Ritchie, C. D.; Coetzee, J. F. “Solute-Solvent Interactions”; Marcel Dekker: New York, 1976; Vol. 2. (13) Ritchie, C. D.; Skinner, G . A,; Badding, V. G. J . Am. Chem. SOC. 1967,89, 2063. (14) (a) Kessler, H.: Feigel, M. Acc. Chem. Res. 1982, 15, 2. (b) Kessler, H.; Feigel, M. Chem. Ber. 1978, 1 1 1 , 1659. (15) Leffek, K. T.; Kim, C. B. Can. J . Chem. 1975,53, 3408. ( 1 6 ) (a) Bunton, C. A.; Huang, S. K. J . Am. Chem. SOC.1974, 96, 515. (b) Sinev, V. V.; Nikolova, T. A . J . Org. Chem. USSR 1984, 85, 707. (17) Arnett, E. M.; Troughton, E. B. Tetrahedron Lett. 1983, 24, 3299. (18) Arnett, E. M.; Troughton, E. B.; McPhail, A. T.; Molter, K. E. J . Am. Chem. SOC.1983. 105. 6172. (19) Arnett, E. M.;’Chawla, B.; Amarnath, K.; Healy, M.; Molter, K. J . Am. Chem. SOC.1985, 107, 5288. (20) Arnett. E. M.; Molter, K. E.: Troughton, E. T. Acc. Chem. Res., in press.
AHhel= a + b(pKa pKR+) (4) if b = c. This remarkably simple equation has been shown to apply19 with good precision over a range of nearly 50 kcal/mol to the heterolysis of over 20 carbon-carbon bonds to form resonance-stabilized carbocations and carbanions of widely varied structures despite the fact that pK,’s were measured in dimethyl s ~ l f o x i d ethe , ~ PKR+’S ~ ~ ~ ~in aqueous sulfuric a ~ i dand ~ AHhet ~ , ~ ~ in benzonitrile. This reflects the fact that solvent variation has very little effect on the relative stabilities of resonance-stabilized ions as has been demonstrated by comparisons of gas-phase ionizations with pKR+values in aqueous acidz7and pKa)s in Me2S0.2s We expect eq 4 to break down as the stabilities of the ions are decreased and steric hindrance is introduced close to the site of coordination, and are presently testing its limits. At the present writing it is the only expression known to use for predicting heterolytic bond energies in solution from easily accessible data and in that way complements in a small way Sidney Benson’s extensive scheme for predicting homolytic bond energies. Rate-Equilibrium Relationships. Against this thermodynamic background we may proceed to the question of kinetics for the coordination and heterolysis reactions. The rate data presented below provide for the first time an opportunity to compare thermodynamic and activation parameters for simple electrophile-nucleophile reactions in which the structures of both parties may be varied systematically and in which there are no leaving groups. (21) Okamoto, K.; Kitagawa, T.; Takeuchi, K.; Komatsu, K.; Takahashi, K. J . Chem. Soc., Chem. Commun. 1985, 173. (22) LeGoff, E.; LaCount, R. B. J. Am. Chem. SOC.1963, 85, 1354. (23) Bordwell, F. G.; Matthews, W. S.; Bares, J. E.; Bartmess, J. E.; Cornforth, F. J.; Drucker, G. E.; Margolin, 2.; McCallum, R. J.; McCollum, G. J.; Vanier, N. R. J . Am. Chem. Soc. 1975, 97, 7006. (24) Bordwell, F. G.; Branca, J. C.; Hughes, D. L.; Olmstead, W. N. J . Org. Chem. 1980,45, 3305. (25) Freedman, H. H. In ’Carbonium Ions”; Olah, G . A,, Schleyer, P. V. R., Eds.; Wiley-Interscience: New York, 1973; Vol. IV, Chapter 28. (26) Breslow, R. Pure Appl. Chem. 1974, 40, 493. (27) Taft, R. W.; Wolf, J. F.; Abboud, J. L. M. J . Org. Chem. 1977, 42, 3316. (28) Taft, R. W.; Bordwell, F. G., personal communication.
Kinetics of Carbon-Carbon Bond Formation Except for a few example^*^^^ of alkyl transfer reactions whose rates and equilibria can both be measured, organic chemists have had to settle for comparing rates of nucleophilic substitution reactions with the Bronsted protolysis constantsof the nucleophiles. Such extensions of the Bronsted equation have been examined by B o r d ~ e l l ~ and ~ - ~found * to extend over a very wide range of rates and pK,'s. However, the question remains: how do the rates for attack of carbocations on carbanions relate to the thermodynamics of such processes as the structures of both species are varied? The results presented below provide the data to examine this question for the first time using rates for equilibrated systems (eq 1) as determined by two independent methods: N M R line broadening and temperature jump perturbation and relaxati~n.~"~
Experimental Section NMR. All N M R experiments were performed on a JEOL FX-90Q equipped with a probe temperature controller. The probe temperature was checked before and after each low-temperature spectrum was taken by inserting a sample of acidified methanol and measuring the separation between the methyl and hydroxyl hydrogen signals. The temperature was determined from the peak separation by using the conversion equations of Van Geet.43,44 Ethylene glycol was used in an analogous way for temperatures above 30 OC. The difference in temperature before and after the experimental spectrum was always less than 0.5 OC and the average of the two determinations was used as the spectrum temperature. Samples consisted of 0.2-0.5 M covalent compound in acetonitrile-d3. At these concentrations peaks due to ions were not detected, and the analysis was performed by using the cyclopropene methyl signals of each compound as follows. Spectra were taken at successively lower temperatures until no further sharpening of the peaks was observed at which point the widths at half-height and the peak separation at no exchange ( k = 0) were recorded. These values were used as inputs for a computer program developed by Professor Donald Chesnut which calculated theoretical spectra for given exchange rates by the density matrix method.41 Experimental spectra were obtained a t various temperatures over a range of about 50 OC and the widths of the methyl peaks were compared to the theoretical spectra. In this way a rate constant was assigned to each experimental spectrum at a given temperature. Temperature Jump Kinetics. Temperature jump experiments were performed on a Durrum-Gibson stopped flow kinetics apparatus equipped with a Model D-150 temperature jump accessory. The sample cell had a volume of 200 pL and an optical path length of 2.0 cm. The sample solution was introduced into the cell from reservoir syringes immersed in water which was continuously circulated through the sample cell block and a coil in a constant temperature bath. Temperature control was within approximately 1.0 OC. The temperature jump was achieved by the discharge through (29) Arnett, E. M.; Reich, R. J. J. Am. Chem. SOC.1980, 202, 5892. (30) (a) Lewis, E. S.;Hu, D. D. J . Am. Chem. SOC.1984,106,3292. (b) Lewis, E. S.;Smith, M. J.; Christie, J. J. J. Org.Chem. 1983, 48, 2527. (c) Lewis, E. S.; Kukes, S.; Slater, C. D. J. Am. Chem. SOC.1980,102, 1619. (d) Lewis, E. S.;Kukes, S. J . Am. Chem. Soc. 1979,202,417. (e) Lewis, E. S.; Vanderpool, S. H. J . A m . Chem. SOC.1978, 100, 6421. ( f ) Lewis, E. S.; McLaughlin, M. L.; Douglas, T. A . J . Am. Chem. SOC.1985, 207, 6668. (31) Hine, J.; Weimar, R. D. J . Am. Chem. Sor. 1965,87, 3387. (32) Kurz, J. L.; El-Nasr, M. M. S. J. Am. Chem. SOC.1982,104, 5823. (33) Schowen, R. L.; Gray, C. H.; Coward, J. K.; Schowen, K. B. J . A m . Chem. SOC.1979, 101, 4351. (34) Brauman, J. I.; Dodd, J. A. J . A m . Chem. SOC.1984, 106, 5356. (35) (a) Albery, W. J.; Kreevoy, M. M. Adu. Phys. Org. Chem. 1979,16, 87. (b) Albery, W. J. Annu. Reu. Phys. Chem. 1980, 32, 227. (36) Bordwell, F. G., personal communication. (37) Bordwell, F. G.; Hughes, D . L. J. Am. Chem. SOC.1984,206,3234. (38) Bordwell, F. G.; Hughes, D. L. J . Org.Chem. 1980, 45, 3314. (39) Caldin, E. F. "Fast Reactions in Solutions"; Wiley: New York, 1964. (40) Caldin, E. F. Chem. Br. 1975, 11, 4. (41) Sandstrom, J. "Dynamic NMR Spectroscopy"; Academic Press: New York, 1982. (42) Finholt, J. E. J . Chem. Ed. 1968, 45, 395. (43) Van Geet, A. L. Anal. Chem. 1970, 42, 679. (44) Van Geet, A. L. Anal. Chem. 1968, 40, 2229.
The Journal of Physical Chemistry, Vol. 90, No. 3, 1986 385 the sample cell of a 1 pF capacitor charged to 2000 V. Joule heating occurred with a time constant of 64 MS. Since the temperature rise in the cell was calculated to be 5.8 OC from the known electrical energy of the charged capacitor and the heat capacity of acetonitrile, the circulating water was controlled at 19.2 OC so that the final reaction temperature was 25.0 f 1.0 O C . Sample solutions (10-4-10-5 M) of covalent compounds were prepared in acetonitrile containing 0.48 M tetra-n-butylammonium tetrafluoroborate as the conducting electrolyte. Equilibrium constants for heterolysis at 25.0 OC were measured for each solution prior to the kinetic experiment by measuring the UVvisible absorption due to the phenylmalonitrile anions with a Varian DMS- 100 ~pectrophotorneter.~ Absorbance change in the sample cell was monitored on a storage oscilloscope which was triggered at the same time as the capacitor discharge. The resulting curve was photographed and the data digitized on an Apple Graphics Tablet and plotted as -In ( ( A , - A , ) / ( A , - A , ) ) vs. time, the slope of which was the first-order rate constant, kOw. Good linearity and zero intercept in this plot was indicative of a first-order process. The forward and reverse rate constants, k,,, and khet, were then calculated according to Caldin40 and Finholt4* from the equations kobsd= kmr((ct)q (A-)cp) khet and Khet = khet/k,,. The accuracy of the temperature jump experiment was checked by measuring the rate for the well-characterized reaction between ferric and thiocyanate ions in aqueous solution. The measured rate constants agreed within experimental error with published values.45 Materials. Sources and Purification of Materials. Acetonitrile was stored over 3A molecular sieves, and then distilled from phosphorus pentoxide and stored under argon. Tetra-n-butylammonium tetrafluoroborate (Fisher Scientific) was recrystallized from ethanol-ether. Acetonitrile-d3 (Aldrich "100 atom % D") was used as received. Carbanion Precursors. The preparation of the p-substituted phenylmalonitriles was reported previ~usly.~ The m-substituted phenylmalonitriles were prepared by an improved procedure, reported here, using (m-chloropheny1)malonitrileas an example: 2.5 g (0.062 mol) of sodium hydride (60% dispersion in mineral oil, Aldrich) was rinsed with pentane, and then suspended in 50 mL of dry tetrahydrofuran under argon. Five grams (0.028 mol) of (3-chloropheny1)acetonitrile(Trans World Chemicals) and 3.4 mL (0.028 mol) of phenyl cyanate4' were dissolved in 50 mL of dry tetrahydrofuran and added dropwise to the stirred sodium hydride suspension. The reaction mixture was then poured carefully over ice and acidified with 6 N HCl. Ether was added, the layers separated, and the aqueous layer extracted with ether (2 X 100 mL). The combined organic layers were rotoevaporated to dryness to remove the tetrahydrofuran, and the resulting solid material taken up in ether. After extraction with saturated sodium bicarbonate (3 X 50 mL) and acidification of the extracts with concentrated HC1, a yellow solid formed which was filtered out and rinsed with water. Recrystallization from cyclohexane afforded light yellow crystals: mp 89.0-89.9 "C; 'H N M R (CDCI,) 6 7.44 (m, 4 H), 5.04 (s, 1 H). Anal. Calcd for C9H5N2Cl:C, 61.21; H, 2.85; N, 15.86; C1, 20.07. Found: C, 61.26; H, 3.02; N , 16.01; C1, 20.14. (m-Trifluoromethylpheny1)malonitrile:mp 89.0-89.6 OC; IH N M R (CDC13) 6 7.74 (m, 4 H), 5.14 (s, 1 H). Anal. Calcd for C1,H5N2F3: C, 57.15; H, 2.40; N, 13.33; F, 27.12. Found: C, 57.05; H, 2.50; N, 13.40; F, 27.15. (m-Cyanopheny1)malonitrile:mp 135.8-1 36.2 O C ; 'H N M R (CDC13)6 7.79 (m, 4 H), 5.15 (s, 1 H). Anal. Calcd for CIOHSN3: C,71.85;H, 3.01;N,25.14. Found: C,72.05;H,3.16;N,25.18. (m-Nitropheny1)malonitrile:mp 136.8-137.5 OC; 'H N M R (CDCI,) 6 8.42 (m, 2 H), 7.85 (m, 2 H), 5.20 (s, 1 H). Anal. Calcd for C9HSN3O2:C, 57.76; H, 2.69; N, 22.45. Found: C, 57.68; H, 2.82; N, 22.57.
+
+
(45) Goodall, D. M.; Harrison, P. W.; Hardy, M. J.; Kirk, C. J. J. Chem. Ed. 1972, 49, 675. (46) Ciabattoni, J.; Curci, R.; Lucchini, V.; Kocienski, P. J.; Evans, G. T. Tetrahedron Lett. 1972, 32, 3293. (47) Zweifel, G.; Murray, R. E. Synthesis 1980, 150.
386 The Journal of Physical Chemistry, Vol. 90, No. 3, I986 Carbanion Salts. The arylmalonitrile salts were prepared as reported previously.* Properties of previously unreported salts are given here: The potassium salt of (m-nitropheny1)malonitrile: 'H N M R (acetone-d,) 6 7.64 (m, 1 H), 7.18 (m, 3 H); UV-vis (CH,CN) , , ,A 286 nm ( e 1.90 X lo4), 321 nm ( e 1.70 X lo4), 495 nm ( 6 7.15 X 10'). The potassium salt of (m-cyanopheny1)malonitrile: 'H N M R (acetone-d,) 6 7.1 (m, 3 H), 6.8 (m, 1 H); UV-vis (CH,CN) A, 316 nm ( e 2.58 X lo4), 381 nm ( e 1.70 X lo2). The potassium salt of (m-trifluoromethylpheny1)malonitrile: 'H N M R (acetonitrile-d,) 6 7.07 (m, 3 H), 6.76 (m, 1 H); UV-vis 309 nm ( t 2.53 X lo4). (CH3CN), , ,A The potassium salt of (m-chloropheny1)malonitrile: IH N M R (acetonitrile-d,) 6 6.83 (m, 4 H); UV-vis (CH,CN), , ,A 306 nm 2.6s x 104). Preparation of Carbocation-Carbanion Products. All the substituted phenylmalonitrile-substituted cyclopropene covalent compounds were prepared as reported previo~sly.~ Properties of previously unreported compounds are given here: Trimethylcyclopropenyl(m-nitrophenyl)malonitrile: mp 98.8-99.6 "C (from benzene-hexane); ' H N M R (CDCI,) 6 7.5-8.4 (m, 4 H), 2.05 (s, 6 H), 1.25 (s, 3 H). Anal. Calcd for CI5Hl3N3O2:C, 67.40; H, 4.90; N, 15.72. Found: C, 67.44; H, 5.03; N, 15.70. Trimethylcyclopropenyl(m-cyanophenyl)malonitrile: mp 73.5-74.5 "C (from benzene-hexane); 'H N M R (CDCI,) 6 7.71 (m, 4 H ) 2.03 (s, 6 H), 1.27 (s, 3 H). Anal. Calcd for CI6Hl3N3: C, 77.71; H, 5.30; N, 16.99. Found: C, 77.55; H, 5.43; N , 16.85. Trimethylcyclopropenyl(m-trifluoromethylphenyl)malonitrile: mp 54.6-55.1 "C (from hexane); ' H N M R (CDCl,) 6 7.66 (m, 4 H), 2.01 (s, 6 H), 1.21 (s, 3 H). Anal. Calcd for C16H13NzF3: C, 66.20; H, 4.51; N , 9.65; F, 19.63. Found: C, 66.40; H, 4.57; N, 9.62; F, 19.86. Trimethylcyclopropenyl(m-chlorophenyl)malonitrile: mp 52.5-53.5 "C (from hexane); 'H N M R (CDCl,) 6 7.38 (m, 4 H), 2.02 (s, 6 H), 1.23 (s, 3 H). Anal. Calcd for CI5Hl3N2C1:C, 70.18; H, 5.10; N, 10.91; C1, 13.81. Found: C, 69.99; H, 5.33; N, 10.98; C1, 13.69. Triphenylcyclopropenyl(m-nitropheny1)malonitrile: mp 184.4-185.1 OC (from benzene-hexane); ' H N M R (CDC1,) 6 7.2-8.3 (m, 19 H). Anal. Calcd for C30H19N302: C, 79.46; H, 4.22; N, 9.27. Found: C, 79.64; H, 4.30; N, 9.22. Triphenylcyclopropenyl(m-cyanophenyl)malononitrile: mp 182.5-183.0 "C (from benzene-hexane); 'H N M R (CDC13) 6 7.2-7.8 (m, 19 H). Anal. Calcd for C3'HI9N3:C, 85.89; H, 4.42; N, 9.69. Found: C, 86.02; H, 4.70; N, 9.71.
Results Rate measurements were originally attempted by stopped flow methods39but it was soon obvious that rates of heterolysis and coordination were both too fast to follow by that technique. However, it had been noted during product analysis that the 'H N M R peaks for the methyl groups attached to the cyclopropene ring were abnormally broad suggesting that exchange was occurring at a rate measurable on the "NMR time scale". These rates were determined in acetonitrile-d3 at 25.0 OC and are listed in the right-hand column of Table I. Since we were not clear what process was determining these rates, the temperature jump perturbation-relaxation method was attempted using 0.48 M tetrabutylammonium tetrafluoroborate (Bu4N+BF4-)as a supporting electrolyte. The initial temperature before discharge of the condenser was set to 19.2 "C so that the final equilibrium would be reestablished at 25.0 "C. Comparison of the two right-hand columns of Table I suggest strongly that the process whose rate is being measured by the N M R method is heterolysis. However, it should be noted that the heterolysis rates by T jump are all 5-8 times as fast as the N M R rates. We know of no published standards for either method which can be used to test either of these techniques under conditions directly comparable to those of this study. However, for the
Arnett and Molter TABLE I: Kinetics of Coordination and Heterolysis by Temperature Jump and NMR Line Broadening at 25 O C in Acetonitrile with and without Supporting Electrolyte k,,:
M-I s-* khct.(l5-I R = -CH,
X = p-NOz p-CN m-NOz m-CN m-CF, m-C1 p-CI -H p-CHI
7.21 f 0.7 X lo4 1.71 f 0.20 X lo5 2.72 f 0.34 X lo5 6.18 f 0.37 X lo5 7.86 f 1.20 X lo5 1.77 f 0.14 X lo6 6.07 f 1.14 X lo6 1.27 f 0.18 X lo7
X = p-NO, p-CN m-NOz m-CN
1.63 f 0.08 1.33 f 0.08 7.80 f 0.91 1.14 f 0.14
5.23 f 0.56 3.36 f 0.39 2.28 f 0.28 1.67 f 0.10 9.43 f 1.44 6.11 f 0.49 2.97 f 0.56 1.65 f 0.24
560 f 80 lo2 190 f 20 10, lo2 10' 33 f 4 X 10' 27 f 4 X 10' 14 f 2 X 10' 3.8 f 0.5 2.2 f 0.2 X 10' X X X X
R = Phenyl lo6 2.04 f 0.10 X lo7 9.95 f 0.60 X lo7 7.74 f 0.90 X lo8 6.51 f 0.81 X
X X X X
khct,bs-'
lo3 10, 10, 10,
a Results from temperature jump kinetics; Experiments performed under conditions of high ionic strength (0.48 M Bu4N+BFcin acetonitrile at 25 OC. bResults from NMR line broadening in acetonitriled , at 25 "C.
reaction of ferric and thiocyanate ions in aqueous acid, our observed rate constant by T jump agrees within experimental error with that calculated from the rate constants measured by Goodall using the same technique.45 Again for the N M R study, our computer program calculated rate constants from the line shapes reported by C i a b a t t ~ n for i ~ ~the tri-tert-butylcyclopropenylazide rearrangement in carbon tetrachloride, which agree within 5% with Ciabattoni's rate constants. Thus, for reactions studied under different conditions we have reasonable agreement with published results for both of our techniques but there are no available data for similar reactions in acetonitrile under our conditions. Ionic Strength Effects. In order to perform the T jump experiment it is necessary to have a reasonable concentration of supporting electrolyte to heat the solution by conductance rather than have the capacitor discharge by arcing. The observed rate curve is the resultant of the opposed heterolysis and coordination rates, and an independent measurement of the equilibrium constant in the salt solution is required to resolve the rate curve. Table I11 lists the salt effect on the free energies of heterolysis: the in Table 11. As difference between AGoheP'tand AGohetMeCN might be expected, the presence of added salt strongly favors heterolysis by the ionic strength effect. However, the concentration of added salt was far above the previously determined limit of applicability of the Debye-Hiickel limiting law.5 Thus, the values reported in Table I1 are derived from the observed of AGohC~aI' equilibrium constants, K , rather than those corrected with calculated activity coefficients, KO . 5 We were unable to make satisfactory direct tests of the ionic strength effect on N M R line broadening because the peaks from added salt overlapped too much of the methyl region of the trimethylcyclopropene system. There remains a good possibility that the kinetic process measured by N M R in pure acetonitrile (AG'h,tMeCN) is not identical with that measured by T jump in the 0.48 M Bu,N+BF4- solution. If that is the case, the following discussion of the ionic strength effects on AGhetoand AH,,,' is based on a false premise and is invalid. Discussion The rate data presented here shall be examined through comparing the response of AGlhet to two types of perturbations (structure and medium) with the corresponding thermodynamic response as manifested by the standard free energy of heterolysis (AGOhet). The opportunity to make such a direct comparison of kinetic and thermodynamic properties for such a simple bondforming and bond-breaking reaction between ions in solution where the structures of both parties can be varied systematically is unprecedented.
Kinetics of Carbon-Carbon Bond Formation
The Journal of Physical Chemistry, Vol. 90, No. 3, 1986 381
TABLE II: Standard Free Energies of Heterolysis and Free Energies of Activation by Two Independent Methods for Heterolysis at 25 Acetonitrile
X = p-NO2
p-CN m-N02
m-CN m-CF3 m-C1 p-c1
-H P-CH3
2.92 f 0.09 3.69 f 0.07 4.20 f 0.10 4.86 f 0.1 1 5.35 f 0.15 6.08 f 0.18 7.24 f 0.21 8.02 f 0.24
-4.32 -3.55 -3.04 -2.38 -1.89 -1.16 0.0 +0.78
13.74 f 0.06 14.00 f 0.07 14.23 f 0.07 14.41 f 0.04 14.75 f 0.09 15.01 f 0.06 15.43 f 0.11 15.78 f 0.09
3.96 f 0.05 5.63 f 0.10 6.82 f 0.30 7.15 f 0.31
-3.19 -1.52 -0.33 -0.0
12.93 f 0.03 13.36 f 0.04 13.50 f 0.07 13.61 f 0.08
2.76 f O.lOd 4.37 f 0.05" 5.63 f 0.02 5.90 f 0.04 6.55 f 0.04 7.29 f 0.08 7.87 f 0.31d 9.44 f 0.15" 10.32 f 0.20"
-6.68 -5.07 -3.81 -3.54 -2.89 -2.15 -1.57 0.0 -0.88
5.39 f 0.03d 7.00 f 0.16d 8.44 f 0.07 8.78 f 0.03
-3.39 -1.78 -0.34 0.0
O C
in
13.7 f 2 14.31 f 0.05 15.40 f 0.05 15.50 f 0.07 15.90 f 0.07 16.64 f 0.08 17.00 f 0.08
R = Phenyl X =p-NO,
p-CN m-NO2 m-CN
" Measured spectrophotometricallyunder conditions used in the temperature jump kinetics experiment, is., 0.48 M Bu4N+BF4-in acetonitrile at 25 "C. bMeasured spectrophotometricallyin acetonitrile at 25 "C. cMeasured in acetonitrile-d3at 25 "C. dReference 5. TABLE 111: Effect on Standard Free Energy of Heterolysis and Rate of Heterolysis for Substituted Phenylmalononitrile Trimethylcyclopropenes from Addition of Swamping Salt to Acetonitrile at 25 OC (in kcal/mol)
substituent
p-CN m-CF3 m-CI p-CI -H PCH3
salt effect on free energy of activation! heterolysis," 6,a~tAGo~ct6aa&G*,,ct -1.45 -1.69 -1.94 -1.79 -2.20 -2.30
-0.57 -0.99 -0.75 -0.89 -1.21 -1.22
6sa,tAG*hct/ 6aaItAG"hct 0.39 0.59 0.39 0.50 0.55 0.53 av 0.49
"&,~,AGohcl = (AGoheFit- A G o h c t M ~ Nfrom ) values in Table 11. b6saltAG*hct= (AG*hctPlt- AG*hetMSCN)from values in Table 11.
Structural Changes on AGlhetand AGOhet. As listed in Table 11, standard free energies of heterolysis in acetonitrile at 25 OC, AGohelMeCN,are given for 13 compounds formed from the reaction of trimethyl- or triphenylcyclopropenium ions with a series of mand p-substituted phenylmalonitrile anions. Corresponding data for free energies of heterolysis in a 0.48 M solution of Bu4N+BFc in acetonitrile, AGohet=lt,and rate data in both media, AG*hetMeCN and AGlhetPlt,are also listed for all of the systems which were accessible by our techniques. There is a practical limit to such rate-equilibrium studies over a wide range as either the forward or reverse reaction becomes diffusion controlled. Since the heterolysis reaction is in equilibrium with the bond-forming coordination of the carbocation with the carbanions, the standard free energies of reaction or activation for coordination are immediately available by changing the sign of AGOhctand adding AH*hetto AGO,,. Figure 1 displays the end product of this research: rateequilibrium comparisons of the heterolysis and coordination for two series of compounds derived from trimethyl- and triphenylcyclopropenium ions. The slopes of the lines, using GrunwaldLeffler notation for the structural variation operator 6R4*are as follows: (1) 4 . 7 9 6 = 6RAG*/6RAGo for triphenylcyclopropenium coordination with arylmalonitrile anions; (2) -0.596 = 6RAG*/ 6,AGO for trimethylcyclopropenium coordination and for the corresponding heterolysis plots; (3) 0.204 = GRAG*/SRAG0for triphenylcyclopropenyl arylmalonitriles; (4) 0.404 = 6RAG*/ &AGO for trimethyl cyclopropenylarylmalonitriles. Relation to the Reactivity Selectivity Principle. Of immediate interest is the fact that the lower two lines for the coordination reactions violate the so-called reactivity selectivity principle (RSP)-the faster reactions (lower AG*) of the triphenylcyclo(48) Lefflar, J. E.; Grunwald, E. 'Rates and Equilibria of Organic Reactions"; Wiley: New York, 1963.
6.0
1
2.0
I
I
4.0
AGO
I
I
6.0
I
I
I
8.0
kcsl mol-'
Figure 1. Correlations of free energies of activation and reaction for
coordination of cyclopropeniumcations with arylmalononitrilecarbanions and the reverse heterolysis, eq 1. Data determined in 0.48 N Bu4BF4in acetonitrile at 25 "C. 0 = Trimethylcyclopropenium, coordination;0 = trimethylcyclopropenium, heterolysis; = triphenylcyclopropenium, coordination; 0 = triphenylcyclopropenium, heterolysis. propenium ion are more selective (greater rate variation) as the structure of the arylmalonitrile anion is varied. Of necessity, this means that for the corresponding heterolysis reactions the slower trimethylcyclopropenylarylmalonitrilesare more responsive to the stabilities of the leaving group and, in this sense, do follow the RSP. The value of the RSP as a predictive tool has come under increasing fire. Several ~ r i t e r s including ~ ~ - ~ ~ ourselves29 have concluded that its failures are so numerous as to render it worthless as a general means for forecasting whether the more reactive reagent in a competitive situation will be more discriminating. (49) Johnson, C. D. Tetrahedron 1980, 36, 3461. (50) Jencks, W. P.; Young, P. R. J . Am. Chem. Soc. 1979, 101, 3288. (51) Bordwell, F. G.; Hughes, D. L. J . Org. Chem. 1980, 45, 3314. (52) Pross,A., personal communication. (53) Giese, B. Angew. Chem., Int. Ed. Engl. 1977, 16, 125.
388 The Journal of Physical Chemistry, Vol. 90, No. 3, 1986
, I
.c
~
2.0
3.0
I
I
4.0
5.0
PKHA
Figure 2. Extended Bronsted plot of coordination rates for reaction of trimethylcyclopropenium cation with arylmalononitrile anions vs. the pKHA’sof the corresponding arylmalononitriles.
Naturally, the loss of its empirical validity has eroded the interpretive value of the RSP and finally its theoretical relevance to rate-equilibrium questions. Serious doubts about the RSP began in the early 1970’s when C. D. Ritchie and his students discovered that the relative reactivities of a series of carbocations with a series of nucleophiles were essentially constant even though the reactivities of the cations varied over ten powers of ten and of the nucleophile over six powers of ten.54-58 During the ensuing years, Ritchie’s equation which expresses the relative rate of reaction of any carbocation in his series solely in terms of the nucleophilicity constant Nf of the anion has survived and been steadily extended. Although none of the cation-anion combinations studied by us are comparable to Ritchie’s series they are obviously directly related. It is therefore amusing to see that our first entry into the field not only does not follow the RSP but reverses it. We join C. D. Johnson in doubting ”if such a widely held concept [as the RSP] has ever rested on slimmer or more dubious evidence”.49 Johnson’s review covers the extensive literature on the validity and significance of the RSP up to 1980. More recent tests and discussions of the matter have been published by B ~ r d w e l l , ~J ’e n c k ~and , ~ ~pros^.^^ One may hope that eventually the RSP will disappear as an article of faith from textbooks of organic chemistry. Relation to the Bronsted Equation. The earliest and most widely used rate-equilibrium correlation is the Bronsted equation59 which remains of seminal i m p o r t a n ~ e to ~ ~the~ understanding ~*~~~ of extrathermodynamic relationships. The original form of the Bronsted equation compared the rates and equilibrium constants for proton transfer to or from a given series of acids or bases. However, it has since been extended by comparing the reaction rates of various acids or bases as electrophiles or nucleophiles in many kinds of reactions with their pK,’s. Such comparisons have been made out of expediency rather than rigor because the rates could be measured, but not the equilibria, and the pK,’s were known. Figure 2 presents such an extended Bronsted plot of log khet(which represents the free energies of activation for heterolysis of the carbon-carbon bonds in the trimethylcyclopropenylarylmalononitrile series) in acetonitrile vs. the pK,’s of the same arylmalonitriles in dimethyl sulfoxide.36A good linear plot of slope -0.461 is obtained for comparison with the more rigorous correlation of log khetvs. log Khetwhose slope is 0.449 and whose Ritchie, C. D.; Virtanen, P. 0. I. J . Am. Chem. Soc. 1972,94, 4966. Ritchie, C. D.; Virtanen, P. 0. I. J . Am. Chem. SOC.1972, 94, 4963. Ritchie, C. D.; Fleischhauer, H. J. Am. Chem. SOC.1972,94, 3481. Ritchie, C. D.; Virtanen, P. 0. I. J . Am. Chem. SOC.1972, 94, 1589. Ritchie, C. D.; Wright, D. J. J . Am. Chem. Sor. 1971, 93, 6574. Ingold, C. K. “Structure and Mechanism in Organic Chemistry”, 2nd ed.; Cornell University Press: Ithaca, NY, 1969; p 425. ( 6 0 ) Hammond, G.S. J. Am. Chem. SOC.1955, 77, 334 (54) (55) (56) (57) (58) (59)
Arnett and Molter linearity is equivalent to that of the free energy plot in Figure 1. Obviously in this case, there is good correspondence between the familiar extended Bronsted treatment and the more relevant correlation of rates and equilibria for the heterolysis reaction itself. The Ionic Strength Effect on Acthetand AGOhe,. If it is true that rate constants derived from the N M R experiments in pure acetonitrile refer to exactly the same heterolytic process as that determined by the Tjump experiment, we can estimate the salt effect on the free energy of activation (6saltAG*kt)as the difference (AGlhetsalt- AG*he,MeCN).These values are listed in Table 111 for six substituted phenylmalononitrile trimethylcyclopropenes. Clearly, the general trend is in the same direction as for 6saltfiGohet but of only about half the magnitude. The salt effect on A G O h e t is the sum of the medium effect (salting out or salting in) on the neutral covalent reactant and the stabilizing ionic strength effect from the tetrabutylammonium tetrafluoroborate solution on the triethylcyclopropenium cation and the substituted phenylmalonitrile anions. The same factors apply to AGlhet. Since both processes have the same initial state, the ratio 6saitAG*/6saltAGoin Table 111 compares the effect of adding 0.48 M Bu4N+BF; to acetonitrile on the solvated transition state for heterolysis to its effect on the partial molar free energies of the ions. By this criterion, charge development and delocalization, both by resonance into the molecular frameworks and by Coulombic interaction with the solvent, are about halfway developed at the transition state. A notion of the tradeoff between internal and external charge distribution in the ions can be gained by comparing the columns in Table I1 for dRAGosaltand 6,AGoMeCNwhich describes the effect of substitution on the phenylmalononitrile moiety in the presence and absence of added salt to the equilibrium system using the unsubstituted system for reference. In every case, the substituent effect in pure acetonitrile is larger than it is in the salt solution. vs. AGohe?lt is linear with a slope of 1.14 A plot of AGohetMeCN and a correlation coefficient of 0.997. The 10-20% larger substituent effects in the absence of salt represents the greater demand that is placed on substituents to delocalize charge in the less polar medium. Relevance to Rate-Equilibrium Theories. Cation-anion reactions have been employed deliberately by Ritchie, and more recently by our group, as one of the simplest conceivable reactions for testing the basic principles of structure-reactivity relationships in solution. Granted that at least 90% of the commonly used processes in the organic laboratory involve electrophiles and nucleophiles in solution, what could be simpler than to study their direct reaction with no covalently bound leaving group? However, it should be clear from our results, and Ritchie’s, that the apparent simplicity of formally making and breaking single bonds by coordination/heterolysis is a mirage and that in solution the reaction is heavily dominated by solvation of the ions. For over half a century it has been recognized that solvation of ions and dipoles is the determining factor which favors polar reactions over radical reactions in solution.59 This must apply a fortiori to reactions which have been tailored to give measurable equilibria for heterolysis/coordination of carbon-carbon bonds in solution. Three facts provide overwhelming evidence that solvation of the resonance stabilized ions in eq 1 is the driving force for heterolysis: (1) In nonpolar solvents no heterolysis is observed. Transfer to a more polar solvent is all that is required to break the central bond. (2) In polar media, the degree of dissociation depends on the dielectric constant. An excellent Born plot of AGOhet vs. 1/ t is ~bserved.~ (3) Full crystal X-ray structures of p-nitro and p-methoxy trimethylcyclopropenylphenylmalononitrileshow no measurable difference in the length (and therefore strength) of the central C-C bond which is heterolyzed even though the pK,’s of the parent malononitriles differ by about 6 units and the AGhet’S by about 8 kcal/mol. These results completely support Ritchie’s conclusion6-’ that cation-anion reactions are controlled primarily by solvation.
J. Phys. Chem. 1986, 90, 389-394 Furthermore, they can only be brought to a point of observable equilibrium by using ionic fragments that are highly stabilized by resonance and solvation. Thus the heterolysis energies we have reported reflect almost entirely the stabilities of the solvated ions and say very little about the strength of the central bond as might be inferred from its length. Put another way, the price that has been paid for being able to measure both the rates and equilibria for these formally simple reactions is to make them polymolecular through the involvement of solvent molecules whose role can be observed only indirectly. Furthermore, this dilemma must be general for rate-equilibrium studies of reactions where ions are created or neutralized in solution. Comparisons of kinetic and thermodynamic properties frequently employ Hammond's proposal for relating the structures and energies of transition states to those of nearby energy minima (reactants, products, or intermediates) along the reaction coordinate.60 This principle was framed originally with great care for application to reactions which are very exothermic or very endothermic so that the transition state must be close in energy to the initial or final state. Cases such as the present one whose free energies and enthalpies of reaction are close to zero with sizeable activation barriers are outside the predictions of Hammond's postulate and the transition state cannot be pictured by simple molecular models. A more formal treatment relating rates and equilibria through the Marcus equation6' has been applied with great success to electron-transfer, proton-transfer, and even to methyl-transfer reactions. Albery has proposed application of the Marcus equation to cation-anion reaction^^^,^^ but Ritchie has provided strong arguments against doing so7 primarily through the lack of appropriate "identity" reactions. M ~ r d o c h ~ " and ~ * Miller69 have (61) (a) Marcus, R. A. J. Chem. Phys. 1956,24,966. (b) Marcus, R. A . J. Phys. Chem. 1968, 72, 891. (62) Albery, W. J. Annu. Rev. Phys. Chem. 1980, 31, 227. (63) Albery, W. J. Pure. Appl. Chem. 1979, 51, 949. (64) Murdoch, J. R.; Donella, J. J. Am. Chem. SOC.1984, 106, 4724.
389
developed more general treatments to include systems that do not have identity reactions, but we fail to see their application to the present case. Position of the Transition State. The slopes of free energy plots such as those in Figure 1 are frequently interpreted as representing the position of the transition state on the potential surface of the reaction. We see little profit in such an exercise in the present case in view of the great complexity of the potential surface which must involve the motions of many solvent molecule^.^ More generally, we are skeptical of such interpretations on the pragmatic grounds that they are tautological in the sense that nothing more is gained by inferring the transition state on a complex hyperspace from a plot of AG* vs. AGO than is available from the linear free energy correlation itself which moreover has true empirical predictive value. The intellectual hazard here is that by restating the empirical result in sophisticated language, one may believe one has learned something of basic value. Five years ago we presented similar misgivings concerning the widespread use of free energy analyses for describing transitionstate s t r ~ c t u r e s .The ~ ~ present study of cation-anion reactions which are formally even simpler than the Menschutkin quaterour previous skepticism of common interpretive n i ~ a t i o nenforces ~~ practices which give an oversimplified view of the activation process in solution. Acknowledgment. We acknowledge with gratitude financial support for this research from The Gas Research Institute, and N S F grant CHE-8006202 and the assistance of D. Meinholtz, A. T. McPhail, and D. B. Chesnut. Acknowledgment is also made to the donors fo the Petroleum Research Fund, administered by the American Chemical Society, for financial support. (65) (66) (67) (68) (69)
Murdoch, J. R. J . Am. Chem. SOC.1983, 105, 2660. Murdoch, J. R. J . Am. Chem. SOC.1983, 105, 2667. Murdoch, J. R. J . Am. Chem. SOC.1983, 105, 2159. Murdoch, J. R. J . Phys. Chem. 1983,87, 1571. Miller, A. R. J . Am. Chem. SOC.1978, 100, 1984.
Thermochernlcal Kinetics: A Success Story Robin Walsh Department of Chemistry, University of Reading, Whiteknights, Reading, RG6 2AD U.K. (Received: July 22, 1985)
This paper presents an analysis of the nature of the subject of thermochemical kinetics, the term coined by Sidney Benson some 17 years ago. The subject is placed in the context of its relationship with various aspects of chemistry and its unifying themes are discussed. To assist this exercise a number of examples are presented from our own and others' recent work. These examples cover a discussion of the stability of silenes (silaolefins), silylenes, and silicon-containing cations. Critical analysis suggests the following heats of formation (AH$/kJ mol-'): Me2Si=CH2, 21 (h20); Me2& 92 (S); Me3Si+,610 (f20). These are consistent with a t-bond energy (Me,Si=CH2) of 172 f 20 kJ mol-' and a divalent state stabilization energy (Me,Si) of 134 h 12 kJ mol-'.
Introduction It is now 17 years since Sidney Benson's "Thermochemical Kinetics" was first published.' This book brought together and focussed many of the ideas that had guided gas kineticists up to that date. It has had and continues to exercise great influence. At the time R. M. Noyes2 wrote "Dr. Benson has attempted to show how to estimate the rate for any hypothetical reaction in(1) Benson, S. W. "Thermochemical Kinetics"; Wiley: New York, 1968, 1st ed.; 2nd ed., 1916. (2) Noyes, R. M. J . Am. Chem. SOC.1969, 91, 3110.
0022-3654/86/2090-0389$01.50/0
volving reasonably conventional compounds ... probably today nobody could do better". On the occasion of this meeting, with its symposium to honor Sidney Benson, the time seems appropriate to attempt some sort of modern appraisal of thermochemical kinetics. To take the subject matter of the book and chapter-by-chapter review modern developments and thought is of itself a daunting task and well beyond the scope of this article. In this paper I take the more modest approach of applying some of the central ideas to examples of chemical systems which we have encountered during our research, largely in more recent years. The purpose of this exercise 0 1986 American Chemical Society
