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 TransitionState. 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
390 The Journal of Physical Chemistry, Vol. 90, No. 3, 1986 /-
.
-
2C2H5
’ Figure 1.
is to demonstrate that the underlying ideas are still as relevant today as they were at the time of writing of ”Thermochemical Kinetics”. Before embarking on the details, however, it is worth making a few general observations. Thermochemical kinetics should not of itself be regarded as a theory but rather as a loose knit collection of ideas, admittedly some of them theoretical. It occupies an area between, but overlapping with, several recognized branches of chemistry (not just physical). To use the modern idiom it may be regarded as being at the interface of “reactivity” with “structure”. Figure 1 shows the connection with specific branches of the subject. As an illustration of its use, thermochemical kinetics can take rate or equilibrium measurements and turn them into entropies and heats of formation. This will be of interest to ab initio calculators. Alternatively, potential energy surfaces from ab initio calculations can be turned into thermodynamic data and used to predict rates and equilibria in practical systems of importance. Or again, mechanisms in complex systems may be analyzed and a key elementary reaction step isolated leading to new information (say a bond dissociation energy) which will be of importance for the structure of a molecular fragment and therefore of interest to valency theory. Thermochemical kinetics tries to paint a broad-brush picture, establishing patterns, both in thermodynamic data (entropies, heats of formation) and in kinetic data ( A factors, activation energies), and, seeking useful generalizations, it tries to keep special cases to a minimum. It has been applied to an enormous range of chemical systems, molecular, free radical, ionic, both in solution and in the gas phase. Because of the complexities of solvent interactions it works best in the gas phase. Its strength and merit lies in this broadness since the implication is that patterns do exist. This forces questions such as the following: “What should I expect for that A factor?” “Is that heat of formation reasonable?” “Why is this reaction reversible and that seemingly similar one not so?” “Why is this particular intermediate involved in the mechanism?” ... By looking for patterns this enables “special cases” to be identified. They may turn out to arise from errors in measurement in which case thermochemical kinetics serves a useful critical function. However, one may legitimately ask “when should a special case arise?”. In this way thermochemical kinetics leads one to ask the right questions and helps one to frame and design experiments. One should enter into a caveat. It is important to recognize that thermochemical kinetics, by itself, does not offer explanations for chemical phenomena at the deepest level. We may illustrate this by an example. The pyrolysis of ethane consists of the steps CzH6 2CH3. (1) CH3. + C2H6 CH4 CZHS. (2) C,H4 + H. CzHy (3) H. + C2H6 Hz + C2Hy (4)
--
-+
+
+
Walsh n-C4HIo(or C2H4
+ C2H6)
(5)
Given the heats of formation of the species involved and given transition state (and RRKM) theory and certain generalizations about activation energy, one may reasonably describe the products and rates of this reactioh under a variety of temperature and pressure conditions. What one may not supply are the answers to the questions: “But why is AHfe(CH3) = 146 kJ mol-’?” “Why is the activation energy for step 2 actually 49 kJ mol-’?”. These are questions for theoreticians. A recognition of this is necessary to avoid attributing virtues to thermochemical kinetics to which it does not lay claim. Some years ago the famous ”biradical mechanism” for cyclopropane isomerization came under particular fire. We discussed this in some detail3 and pointed out that regardless of the depth of the potential well associated with the biradical on the potential energy surface (even, in extremis, if there was no well) the biradical idea retained its usefulness as a predictive tool for activation energy estimation. This illustrates how, although the underlying theoretical basis of an idea may change, thermochemical kinetics retains its usefulness. In what follows this usefulness is illustrated in a variety of examples.
Examples 1. The Stability of Silenes (Silaolefins). In 1972 we addressed the question of the existence of the silenes, made more fascinating by statements in textbooks to the effect that pT-pr bonding did not occur involving elements outside of the first row of the periodic table. This work was stimulated by the pioneering experimental studies of Flowers and Gusel’nikov5 on the pyrolysis of 1,l-dimethylsiletane (dimethylsilacyclobutane):
I
-’D
=== Me,Si=CH,
+
CH ,,
(6, -6)
For this process one may write the relationship
which is a typical thermochemical relationship arising from simple consideration of the bonds broken and the bonds made during the dissociation process. This relationship then contains the unknown D,(Si=C), the 7~ bond strength in the product silene which is one index of its stability.6 D,(Si=C) may be evaluated provided all of the other quantities in eq A are known or may be reliably estimated. At the time of our publication4 we took such figures as were available and obtained D, = 138 f 22 kJ mol-’. The fairly wide error limits arose mainly out of uncertainty in m 6 , 4 which rested on experimental parameters. Other quantities, however, were also subject to some uncertainty, viz. D,(Si-C) and E (ring strain). Fortunately D,(C-C) and D,(C=C) were reasonably well-known.’,’ Thus apart from the experimental uncertainty itself this points to the need to have reliable values for Si-C (single) bond strengths and to know ring compound strain energies. This system, therefore, illustrates how thermochemical kinetics gives the signpost to further research. (3) Frey, H. M., Walsh, R. “Gas Phase Kinetics and Energy Transfer”, Donovan, R. J., Ashmore, P. G., Ed.; The Chemical Society: London, 1978; Specialist Periodical Reports, Vol. 3, p 1. (4) Walsh, R. J . Organometal. Chem. 1972,38, 245. (5) Flowers, M. C., Gusel’nikov, L. E. J . Chem. SOC.B 1968, 419, 1396. (6) Bond dissociation energies are of course related to heats of formation through simple cycles. Each term may be thought of therefore as an equivalent statement about the stability of a species. However, bond dissociation energies, in this case D,(Si=C), are usually much more easily grasped and therefore more directly illuminating. (7) In recent years C-H bond dissociation energies have been upwardly adjusted by ca. 8 kJ mol-’ from those originally proposed by Bens0n.I This require modification of D,(C-C) and D,(C=C) as well. Some of this is foreshadowed in an article by Doering.* The figures used here are based on AH$(C2H5,298 K) = 116 kJ mol-’.
Thermochemical Kinetics
The Journal of Physical Chemistry, Vol. 90, No. 3, 1986 391
It is not my purpose here to review our bond dissociation energy measurements in organosilanes. These have been well documer~ted.~A new and more reliable value for D,(Si-C) has been ~ b t a i n e d .The ~ ring strain energy, however, has not been reliably established perhaps until recently. A value was originally assumed based on the analogous thietane ring strain.4 An alternative value has been suggested from an analysis of the "CATCH Tables"Io by O'Neal and Ring." The "CATCH Tables" give a rather scattered fit to thermodynamic additivity schemes and recently a new value is available from a fluorine bomb calorimetric determinationL2of A I P f (1,l-dimethylsiletane) = -108.7 f 6.0 kJ mol-'. From our own additivity schemeI3 this gives a ring strain energy of 95 (f8) kJ mol-', somewhat higher than earlier values. Better values for D,(Si-C) and E (ring strain) are therefore now available. As far as A P 6 , . . 6 is concerned, however, a problem remains. The procedure for its evaluation is a little involved. It is obtained indirectly via
where E6 and E+ are activation energies. E6 is measured (262 kJ m01-I)~but E+ has to be inferred from data on the competition between step -6 and step 7, viz.
(7)
E.+ - '/2E7 is given5as 60 f 17 kJ mol-' and E, as zero by rather crude flow e ~ p e r i m e n tin ' ~ which k7 was found to be 106.6dm3 mol-' s-I, independent of temperature. These activation energies are unfortunately associated with A factors which are not mutually consistent. However, within a wide margin of adjustment the data may be reconciled. This is shown in section (i) of the Appendi~.'~ What is really indicated is the need for further experimentation to obtain a reliable measure of the rate constant k7 and its associated Arrhenius parameters. Since this is the major pathway for silene removal in many systems (not only this one) it is a key reaction. Here is another signpost. The best reconciliation of the data in our view yields Ed = 84 kJ mol-' but it has to be acknowledged that, in view of the assumptions involved, an uncertainty of f 2 0 kJ mol-' remains. With all the elements now assembled eq A may be evaluated (kJ mol-') 182 = 362
+ 351 - D,(Si=C)
- 264 - 95
i.e. D,(Si=C)
= 172 ( f 2 0 ) kJ mol-'
This result shows a considerable increase in value from our original result, without much diminution of uncertainty. It reinforces the view originally expressed that these species are quite stable with respect to unimolecular decomposition albeit very reactive toward bimolecular reactions with other molecules. Since our first publication a number of other estimates of D,(Si=C) both experimental and theoretical have been rnade,l6 the majority being not in any great disagreement with this estimate. Most notable (8) Doering, W. von E. Proc. Natl. Acad. Sci. U.S.A. 1981, 78, 5279. (9) Walsh, R. Acc. Chem. Res. 1981, 14, 246. (IO) Pedley, J. B.; Iseard, B. S. "CATCH Tables";University of Sussex: 1973. Pedley, J. B.; Rylance, J. "N.P.L.Computer Analysed Thermochemical Data"; University of Sussex: 1977. (11) ONeal, H. E.; Ring, M. A. J. Organornetal. Chem. 1981,213,419. (12) Steele, W. C., unpublished results (private communication). (13) Doncaster, A. M.; Walsh, R. J. Chem. SOC.,Faraday Trans. I , in press. (14) Gusel'nikov, L. E.; Konobeyevsky, K. S.; Vdovin, V. M.; Nametkin, N . S. Dokl. Akad. Nauk. SSSR 1977, 235, 1086. (15) This forms a part of a more substantial review of thermochemistry of silicon-containingcompounds: Walsh, R., to be submitted for publication. (16) Some of these are mentioned in our re vie^.^
of all, room temperature stable members17 of the silene family have been prepared by the device of blocking access to the T bond with large groups. 2. The Relative Stabilities of Silenes and Silylenes. It was discovered by Gordon in 197818 by means of a theoretical calculation that the divalent silicon species, silylenes, should lie close in energy to their isomeric silene counterparts, viz. the energy (or enthalpy) change for the reaction
-
Me2%: MeSiH=CH, (8) was close to zero. This process as with reaction 6 is susceptible to a simple thermochemical analysis, viz.
+
A P 8 = D,(C-H) - D,(Si-H) - D,(Si=C) DSSE (B) where the quantities represent once again the energies of the bonds broken and made. To evaluate eq B three of the quantities are fairly assessible, viz. D,(C-H), D,(Si-H), and D,(Si=C). The new quantity here, DSSE refers to divalent state stabilization energy. It has long been recognized that the divalent state species among lower group IV (group 14)38members of the periodic table (sometimes called the "heavy carbenes") are stabilized. We have discussed this point in detail9 and do not repeat that here. The question that arises is what value to attribute to DSSE (and the subsidiary question: is it dependent on substituents?). Using an operational definition of the difference between the first and second dissociation energies, we derived a value of 109 kJ mol-' for the DSSE in SiH; and, assuming a constant energy increment on Me for H substitution, a value of 113 kJ mol-' for DSSE in Me2Si.9 Since this involves a number of assumptions we looked again to thermochemical kinetics to offer another approach. A method with some promise may be based on the disilane decomposition studies by Davidson and Matthews,19 an example of which is the reaction Me3SiSiMe2H F! Me3SiH Me2%: (9, -9) for which log (k9/s-') = 12.93 - 198 kJ mol-'/RT In 10 was obtained. The examination of such a system is almost a classical exercise in thermochemical kinetics. This is outlined here but full ASe is details are given in section (ii) of the A p p e n d i ~ . ' First ~ obtained from known thermodynamic data, additivity estimates, and structural analogues. Then, via Zpe, it is adjusted to the mean temperature of experimental investigation. This gives the ratio of forward and reverse A factors, viz. A9/A-9 = 105.3mol dm-3. From the measured A9 therefore A-g = 107.6dm3 mol-' s-I. We then examine a whole range of disilane decompositions and, where available, relative insertion datal5 and conclude largely on the basis of low selectivity of silylene insertion reactions that A+ is too low by as much as a factor of Again there is a signpost here to the real need for absolute rate measurements for silylene insertions. Fortunately one has very recently been measured albeit in cyclopentane solution and only at one temperature. Griller and colleaguesZoobtained k = 2 X lo6 dm3 mol-' s-I for the insertion process Me2%: + Et3SiH Et3SiSiMe2H (10)
+
-
-
With an estimated A of 109.2dm3 mol-' s-' this implies E 16 kJ mol-'. This provides the model value for E-g. E, is adjusted for the underestimation of A9 (by the same factor of and A P 9 , + = E9 - E-9 is obtained. This is then corrected to room temperature and yields A P 9 , + = 207 kJ mol-'. From APfvalues for Me3SiH (reasonably w e l l - k n ~ w n )and ~ ~ ' Me3SiSiMe2H ~ (requiring estimation and pointing out the serious need for such data on disilanes in general) we finally obtain a value for APf(MezSi:). From other disilane systems similar values emerge and an average gives APf(Me2Si:) = 92 f 8 kJ mol-'. This corresponds to a DSSE of 134 f 12 kJ mol-', somewhat higher than previously (17) Brook, A. G.;Abdesaken, F.; Gutekunst, B.; Gutekunst, A,; Kallury, R. K. J . Chem. SOC.,Chem. Commun. 1981, 191. (18) Gordon, M. S. Chem. Phys. Lett. 1978, 54, 9. (19) Davidson, I. M. T.; Matthews, J. I. J . Chem. SOC.,Faraday Trans. 1 1976, 72, 1403.
( 2 0 ) Nazran, A. S.; Hawari, J. A.; Griller, D.; Alnaimi, I. S.;Weber, W. P. J . Am. Chem. SOC.1984, 106, 7267.
Walsh
392 The Journal of Physical Chemistry, Vol. 90, No. 3, 1986 proposed, thus suggesting that methyl groups in fact increase DSSE relative to hydrogen (this is perhaps not surprising as carbon is more electronegative than hydrogen and it is known that halogen substitution increases DSSE even more substantially2'). With the data now available and assembled we substitute into eq €3 A P 8 / k J mol-' = 415 - 378 - 172 134 = -1 (f24)
+
The high uncertainty arises not so much from DSSE as from D,(Si=C) with its intrinsic uncertainty as discussed in example 1 (and the additional assumption of no effect of H for Me substitution). This result corresponds, then, to a very small energy difference (which could, within error limits, be zero) between these two isomeric species. More elaborate ab initio calculationsz2 confirm this and Gordon's original findings. In a sense since theory reached this result first, it may be objected that thermochemical estimates are irrelevant. But this argument overlooks two important points. The first is that the thermochemical method (eq B) involves a trivial calculation (although with a far from trivial evaluation of input energies) and is essentially much simpler than ab initio calculations. The second is that it is important to calibrate against other methods or experiments. One such experiment is discussed in the next section. 3. Some Comments on the Stabilities of Silicon-Containing Zons. We were stimulated to look into the thermochemistry of silicon-containingions by some striking anomalies which have been recently brought up from studies of ionic systems. One of these concerned a claim that the important Si-Si bond dissociation energy, D(Me3Si-SiMe3), had the value 265 kJ mol-'. This result was obtained by Szepes and Baer23using the PEPICO technique to study the dissociation of the Me6Si2+ion and (by RRKM modeling) to deduce the dissociation energy (equivalent to the photoionization threshold for Me3Si+appearance). While there is no criticism of the technique itself, in order to use the results to establish thermochemistry it is important to assess the uncertainties of both the method and the ancillary information required. This is especially so since the "accepted" valuez4for D(Me,Si-SiMe3) is 337 kJ mol-'. The latter value is deduced from a direct kinetic study of the reaction
-
Me3SiSiMe3
2Me3Si
(1 1)
for which Davidson and Howard obtained log
(/C,~/S-*)
= (17.2 f 0.3) - (337 f 4 kJ mol-')/RT In 10
The A factor is reasonable for such a process and the only uncertainty beyond experimental scatter involves extrapolation from the temperature of study (ca. lo00 K) to room temperature. This should add no more than f4 kJ mol-' leading to a maximum uncertainty of f 8 kJ mol-I. Szepes and Baer's valuez3 is based on the relationship
TABLE I: Experimental Studies" of the Heat of Formationb of the Trimethylsilyl Cation (298 K)
AHle(Me3Si+)/ kJ mol-'
-----
reaction
Me4Si + e- Me$+ + CH3 + 2eMe3SiH + e- Me&+ + H.+ 2eMe3SiCI + e- Me3Si++ CI. + 2eMe@ + hw Me3Si++ CH3 + eMe3Si++ NH3 Me2Si=CH2 + NH4* Me4Si + hw Me3%++ CH3 + eMe3SiBr + hv Me3Si++ Bi + eMe3SiI + hw Me&+ + I. + e-
61 1
632 642 590 625 f 20 619
628 629
ref 25 25 25 26 27 23 23 23
#This compilation includes only more recent work. Another study (1982)'O involving H- or CI- transfer is not listed since the results only set very broad limits (548-674 kJ mol-') for AHfe(Me,Si+). bEvaluated as described in the text. = -233.2 f 3.2 kJ mol-' finally confirms the data for the methylsilanes which we inferred earlier9s'3 (but see below). Examination of the table but without critical analysis of each experimental technique suggests AHfe(Me3Si+) = 610 f 20 kJ mol-'. Thus it must be recognized that there is a fairly high uncertainty associated with AHfe(Me3Si+). Part of this may be due to the difficulties associated with measurement of the true ionization onset but it is still clear that further (and better) determinations of this quantity are required. The third item of interest AHHfe(Me3SiSiMe3)is equally subject to uncertainties. Only the "CATCH Tables"Io supply a value, and this, for reasons mentioned earlier, must be suspect. This quantity has been discussed by ONeal and Ring'' and is assigned by them an uncertainty of f 2 5 kJ mol-'. Once again here is a pointer to the need for more experimental work. In fact Me3SiSiMe3 is another key compound for heat of formation studies, since upon it depends not only this argument but also other disilane heats of formation (via additivity rules) and furthermore much silylene and silyl radical thermochemistry. Returning to eq C we may combine uncertainties to give the most probable uncertainty in D(Me3Si-SiMe3) as given by (6/kJ (2 X 20)2 (25)2, i.e. 6 = f51 kJ mol-'. mol-')2 = (2 X This suggests that this result is considerably less certain than that of Davidson and HowardSz4There is, however, an alternate and simpler method of treating Szepes and Baer's result which removes a number of uncertainties, in the form of
+
+
Do(Me3Si-SiMe3) = D,(Me,Si-X)
+ AE, - AEo(X)
(D)
where AEo represent the appearance threshold for Me3%+,given in their paper as 941 f 10 kJ mol-' (9.77 f 0.1 eV). There are several published values for AHfe(Me3Si+) based on a variety of different methods, mainly threshold energy measurements for isomerization or hydride-transfer studies. We have collected a number of these in Table I. However, the listed values are not those of the original publications but based on our own assessed9 heats of formation of the neutral precursors to the ions. In spite of the checkered history of the thermochemistry of silicon compounds there is no reason to believe the values for the heats of formation9 of Me3SiX compounds (X = H, Me, C1, Br, I) are substantially in error. Steele's recent value** for AHfe(Me4Si)
where AE,,(X) refers to the appearance threshold for Me,Si+ from Me3SiX. Szepes and Baer23have measured A b ( X ) and we have calculated9 D298(Me3Si-X) from AHfe(Me3Si). If AEo is increased3' from the experimental value of 9.77 to 9.82 eV (948 kJ mol-') and appropriate temperature corrections are made then the data yield for Dzgp(Me3Si-SiMe3)values of 324, 316, and 31 5 kJ mol-', respectively. These values are uncertain to about f 1 2 kJ mol-', arising from the joint experimental uncertainties in Do(Me,Si-X) and AEo. Thus the discrepancy between Davidson and Howard's result and the new PEPICO results is considerably reduced by avoidance of uncertain thermochemical data. Whether ionic thermochemistry (of silicon compounds) should be regarded as of sufficient reliability to use to evaluate neutral thermochemistry is an interesting general question. Baer has argued that the crossover energy in an ion breakdown diagram provides a better basis for ionic heat of formation determination than the traditional appearance threshold for ions, with its attendant uncertainty in thermal energy of the ions. In this respect it is interesting to note the differences in AHp(Me,Si+) arising from the PEPICO experiments on Me4Si, Me3SiBr, and Me3SiI. If there are no other sources of error they imply small inconsistences
(21) Walsh, R. J . Chem. SOC.,Faraday Trans. I 1983,79,2233. (22) Schaeffer, H. F. Acc. Chem. Res. 1979,12, 288. (23) Szepes, L.; Baer, T. J . Am. Chem. SOC.1984,106, 273. (24) Davidson, I . M. T.; Howard, A. V. J . Chem. Soc., Faraday Trans. 1 1975,71,69.
(25) Potzinger, P.; Ritter, A,; Krause, J. Z . Naturforsch. 1975,30,347. (26) Murphy, M. K.; Beauchamp, J. L. J . Am. Chem. SOC.1977,99,2085. (27) Pietro, W.J.; Pollack, S . K.; Hehre, W. J. J . Am. Chem. SOC.1979, 101, 7126. (28) Steele, W. V. J . Chem. Thermodyn. 1983, I S , 595
D(Me,Si-SiMe,) = 2AEo - 2AHfe(Me,Si+)
+ AHfe(Me3SiSiMe3) (C)
Thermochemical Kinetics
The Journal of Physical Chemistry, Vol. 90, No. 3, 1986 393 u e 6 , 4 ( 7 0 0 K) = 177 f 8 J K-' mol-'
TABLE I1 comud
Se/ J K-' mol-'
Me3SiSiMe2H MesSiH Me2Si
416 334 286
Cpe/J K-' mol-' 300 500 800 165 95 74
256 149 101
344 203 132
This gives (via In A f / A , = A P / R - 1 - In (R'T)]) ref
therefore
11 1
A-6 = 108.56*0.5 M-1
in measured AHfe values (either AHfe(Me4Si) is too low, or AHy(Me3SiBr) and AHy(Me3SiI) are too high by a.9 kJ mol-'). Another example of an anomalous conclusion from ionic studies is provided by an ion cyclotron double resonance study of proton (and deuteron) transfer from Me2SiD+ to a variety of bases, carried out by Hehre and c o - w o r k e r ~ . ~The ~ key processes are -En+
B
+
(CH3),SiD+
-
CH3SiD=CH2
(CH3),SiD+
+
B
f
cH'
D'
-
log (k-6/M-'
+
BH+
CH3 -t BD+
In this case BH+ and/or BD+ formation would depend on a competition between four-center and three-center elimination and the outcome would depend on exit channel barriers and activated complex structures. Whatever the explanation, this is clearly an experimental result worth further investigation.
Appendix (i) Thermochemistry of 1,l-Dimethylsiletane Decomposition. Based upon the experimental data of ref 5: log (k6/S-') = (15.64 f 0.30) - (262 f 3)/6' S-I)
= 8.56 - 82/e'
(b) A , = 106.6M-' s-I (a s measured). If this is combined with the experimental values for k-6/k71/2,then log (k4/M-' s-') = 6.6 - 60/W This is now adjusted, at the mean experimental temperature (700 K), to fit the calculated value for A-6 yielding log (k-6/M-' s-') = 8.56 - 86/0' Thus, with either assumption, E4 comes out with a similar value and we taken E-6 = 84 kJ mol-'. In combination with E6 this yields Ap6,-6(700 K) = E6 - E-6 R T = 184 kJ mol-'. Correction to room temperature3* gives
+
Ap6,-6
= 182 (f20) kJ mol-'
From the values for AHfe( 1,l-dimethylsiletane)12 and AHfe(C2H4),I therefore (ii) Thermochemistry of Pentamethyldisilane Decomposition. The data in Table I1 for Se and Cpe may be obtained by use of group additivity1' and, in the case of Me2Si, the assumption that values for the known Me2S (dimethyl sulfide) may be used' (structural analogue principle). For the reaction Me3SiSiMe2HF= Me3SiH + Me2Si: (9, -9) the data in Table I1 yield ASe(298 K) = 144 J K-l mol-'
-
Acknowledgment. I thank Tom Baer for helpful discussion of the results of his experiments.
log (k-6/k7'/2/M-1/2
S-I)
AHfe(Me2Si=CH2) = 21 (f20) kJ mol-'
CH3 >i=CH,
(CH,),Si:
This, combined with the experimental measurements, suggests A 7 -- 10'0.5'2.6 M-' s-I which is in poor agreement with the experimental value (ref 14) of k7 = A , = 106.6M-' s-I ( T = 298-573 K). We proceed with either of two extreme assumptions. (a) A , = 10'0.5M-I s-' (as calculated above). The experimental result for A,, adjusted on the assumption of the correct magnitude for k7 at the maximum temperature, yields which combined with the experimental values for k4/k71/2gives
By the use of a variety of bases, B, it was found that the onset of BH+ formation occurred at 117 kJ mol-' lower in energy than that of BD+ formation. This was interpreted as meaning that product species differed in stability by this energy. As we have discussed in example 2 thermochemical estimates and ab initio calculations are in agreement that no such ground-state energy difference exists. It seems extremely unlikely that the conclusion drawn from this experiment can be correct. However, it does not follow that this appealingly simple experiment should, therefore, be dismissed. Clearly there is something requiring explanation here. Although such experiments apparently work in other cases maybe there is a special situation here. One possibility, entertained by the authors, is that an excited state of (CH3)2Si: (maybe a triplet) is produced upon de-deuteration. Another may be that the reaction does not occur on a direct surface but rather via an addition-elimination sequence involving an intermediate (similar in type to the group I11 (group 3)38and group V (group 15) donor-acceptor molecules), viz.
-
s-I
log (k,/M-' s-')= 10.5 - 43/W
%(CH3),Si:
CH3 S ' 'i
A6/A-6 = 107,08'0.5M
11
= (3.3 f 1.2) - (60 f 17)/8'
at T = 667-724 K where M mol dm-3 and 0' = R T In lO/kJ mol-'. Evaluation of A Factors: Based on structural analogies3' (as in ref 4)
ACpe(298-650 K) = -2 J K-l mol-' hSe(650 K) = 143 J K-l mol-' This gives (via In ( A f / A , )= ASe/R - 1 - In (R'T)') A9/A-9 =
M
The experimental data of ref 19 gives log (k9/~-') = (12.93 f 0.31) - (198 f 4)/0' T = 620-700 K
Thus A9 and A9/A-9combine to give A- 9 -- 107.6 M-l
s-l
From the experimental measurements of
and a similar
~~
(29)Pau, C.F.;Pietro, W. J.; Hehre, W. J. J. Am. Chem.SOC.1983,105, 16. (30) Eyler, J. R.;Silverman, G., Battiste, M. A. Organometallics 1982, I , 477. (31) The analogous process is that for dimethylcyclobutane decomposition.4 The analogy will be reliable so long as the silicon-for-carbon replacement does not produce an imbalance of low-frequencyvibration modes. In the absence of detailed assignments this seems a reasonable assumption.
= 4.5 J K-l mol-' (again by structural analogy3'). (32)Using (33)John, P.; Purnell, J. H. J . Chem.SOC., Faraday Trans. I 1973,69, 1455. (34)Cox,B.; Furnell, J. H. J. Chem.Soc., Faraday Tram. I 1974,70,859. (35)Vanderweilen, A.J.; Ring, M. A,; ONeal, H. E. J . Am. Chem.SOC. 1975,97,993. (36)Dzarnoski, J.; Rickborn, S.F.; ONeal, H. E.; Ring, M. A. Organometollics 1982,1 , 1217.
J. Phys. Chem. 1986, 90, 394-397
394
-
analysis for each system it appears that a per-bond insertion A factor for Me2Si should be ca. 109.2(*1.0) M-' s-I (a direct, absolute value is not available). The presumption has to be that here the experimental Arrhenius parameters are subject to some distortion. Thus, to proceed we take
Correction via ACpe gives Ap9,+(298 K) = 207 kJ mol-' From AHfe(Me3SiH) = -163 kJ mol-' and Mfe(Me3SiSiMe2H) = -279 kJ mol-] (additivity e ~ t i m a t e ' ~we ) obtain
log ( k 9 / ~ - '=) 14.5 - 217/8'
AHfe(Me2Si) = 91 kJ mol-'
From the experimental data of ref 20 (see text), we estimate
From a similar analysis of other methylated d i ~ i l a n e decom'~ positions, values in the range 88-96 kJ mol-' emerge. Together with other uncertainties this suggests 92 f 8 kJ mol-'. By definition
log (k-g/M-'
S-I)
= 9.2 - 16/8'
At 650 K A p 9 , - 9 = E9 - E-9
+ R T = 206 kJ mol-'
(37) This small adjustment is accommodated within the model used to fit
the Me6Si2* ion breakdown data: T. Baer, private communication. (38) The group notation in parentheses is in accord with recent actions by IUPAC and ACS nomenclature committees. A and B notation is being eliminated because of wide confusion. Group I becomes groups 1 and 11, group I1 becomes groups 2 and 12, group I11 becomes groups 3 and 13, etc.
DSSE(Me2Si) = D(Me3Si-Me) - D(Me,Si-Me) From published values of AHr(Me4Si),28AHfe(Me3Si),9and the above AHfe(Me2Si) DSSE(Me,Si) = 134 f 12 kJ mol-' Registry No. Me2Si=CH2, 61 153-00-2; Me2%:,6376-86-9; Me&+, 28927-31-3; Me3SiSiMe2H, 812-15-7; 1,l-dimethylsiletane, 2295-12-7.
Hexafluorocyclopropane and Octafluorocyclobutane: A Study of the Strain Energies Joel F. Liebman,* Department of Chemistry, University of Maryland Baltimore County, Catonsville, Maryland 21 228
William R. Dolbier, Jr.,* Department of Chemistry, University of Florida, Gainesville, Florida 3261 1
and Arthur Greenberg* Department of Chemical Engineering and Chemistry, New Jersey Institute of Technology, Newark, New Jersey 07102 (Received: July 22, 1985)
The strain energies of hexafluorocyclopropane and octafluorocyclobutaneare discussed in terms of Benson group increments. Unlike the near constancy of t h e heat of formation of the C(H),(C), group as found in a variety of acyclic species, that for its fluorinated counterpart is found to be. very dependent on its next-nearest-neighborenvironment and the groups C(F),(CF,), and C(F),(C), must be carefully distinguished. Regardless of how the group increments are chosen, the strain energy of hexafluorocyclopropaneis found to be remarkably high relative to the parent cyclopropane, while that of octafluorocyclobutane, likewise low relative to cyclobutane.
Introduction Geminal-difluoro substitution provides marked thermodynamic stabilizations in alkyl derivatives which are diminished or absent in substituted cyclopropanes and olefins.' For example, 1,ldifluorocyclopropane has been calculated to have about 12 kcal/mol more strain (relative to the acyclic 2,2-difluoropropane) than cyclopropane (relative to propane)? and this value is reflected in the unusually low barriers to thermal rearrangements of these c o m p o ~ n d s . ~It is a bit surprising that the e x t r a s t r a i n in 1 , l difluorocyclopropane is only about 2 kcal/mol greater than the 4.5-5 kcal/mol increment per F substituent proposed by O N e a l and Benson4 since the thermodynamic favoring of gem substitution' is a measure of nonadditivity. Hexafluorocyclopropane extrudes CF2 under fairly mild cond i t i o n ~and ~ this could arise from extra strain which might be about (1) Sheppard, W. A,; Sharts, C. M. "Organic Fluorine Chemistry"; Benjamin: New York, 1969; pp 20-40. (2) Greenberg, A.; Liebman, J. F.; Dolbier, W. R., Jr.; Medinger, K. S.; Skancke, A. Terrahedron 1983, 39, 1533-1538. (3) Dolbier, W. R., Jr. Acc. Chem. Res. 1981, 14, 195-201. (4) O'Neal, H. E.; Benson, S. W. J. Phys. Chem. 1968, 72, 1866-1887. (5) Birchall, M. J.; Fields, R.; Haszeldine, R. N.; McClean, R. J. J. Fluorine Chem. 1980, 15, 487.
0022-3654/86/2090-0394$01.50/0
27-30 kcal/mol (total strain 55-58 kcal/mol) if one simply employs the O'Neal-Benson increment. Bernett estimated a strain energy of 68.6 kcal/mol for this molecule.6 In an earlier paper,, we employed an experimental value for AHfo(g) for hexafluorocyclopropane,' derived from mass spectrometric rather than calorimetric data, along with a published C(F),(C), group increment (-104.9 kcal/mol)* in order to derive a total strain energy of 80.9 kcal/mol. This value was checked against a value derived from published thermochemical data9 for eq 1 (76.7 kcal/mol) F
x
F
-
3C3Fe- 3CzF6
(1)
F F
and the derived conclusion was that "non-next-nearest-neighbor" corrections are quite small. The purpose of the present paper is (6) Bernett, W. A. J. Org. Chem. 1969, 34, 1772-1776. (7) Bomse, D. S.; Berman, D. W.; Beauchamp, J. L. J. Am. Chem. Soc. 1981, 103, 3967-3971. (8) Dolbier, W. R., Jr.; Medinger, K. S. Terrahedron 1982,38,2411-2413. (9) Pedley, J. B.; Rylance, J. "Sussex-NPL Computer Analysed Thermo-
chemical Data: Organic and Organometallic Compounds"; University of Sussex: Sussex, England, 1977.
0 1986 American Chemical Society
