Enthalpies of hydrogenation and formation of enones. Resonance energies of 2-cyclopentenone and 2-cyclohexenone

Article abstract of DOI:10.1006/jcht.1998.0407

The enthalpies of hydrogenation of 2-cyclopentenone, 2-cyclohexenone, 2-methyl-3-isobutoxycyclohex-2-enone, and 5-methylhex-5-en-2-one have been measured and found to be -(97.1 ± 2.7) kJ·mol^-1, -(109.2 ± 2.9) kJ·mol^-1, -(106.1 ± 3.0) kJ·mol^-1, and -(113.2 ± 2.7) kJ·mol^-1, respectively. In combination with the enthalpies of formation of the corresponding alkanones, we obtained the enthalpies of formation Δ_fH^o_m (cyclopentenone) = -(100.3 ± 3.0) kJ·mol^-1 and Δ_fH^o_m (cyclohexenone) = -(121.9 ± 3.0) kJ·mol^-1. By using isodesmic reactions, the resonance energy of cyclic enones is estimated to be (13 ± 4) kJ·mol^-1.

Full text of DOI:10.1006/jcht.1998.0407

J. Chem. Thermodynamics 1998, 30, 1393᎐1400
Article No. ct980407
Enthalpies of hydrogenation and formation
of enones. Resonance energies of
2-cyclopentenone and 2-cyclohexenone
Donald W. Rogers,a Yeping Zhao,
Chemistry Department, Long Island Uni¨ersity, Brooklyn,
NY 11201, U.S.A.
Marit Traetteberg,
Department of Chemistry, Norwegian Uni¨ersity of Science and
Technology, N-7055 Trondheim, Norway
Martin Hulce,
Department of Chemistry, Creighton Uni¨ersity, Omaha,
NE 68178, U.S.A.
and Joel Liebman
Department of Chemistry and Biochemistry, Uni¨ersity of Maryland
Baltimore County, Baltimore, MD 21250, U.S.A.
The enthalpies of hydrogenation of 2-cyclopentenone, 2-cyclohexenone, 2-methyl-3-
isobutoxycyclohex-2-eyno1 ne, and 5-methylhex-5-eyn-12-one have been measuryed1 and found to be
.
.
.
Ž
.
Ž
.
Ž
.
Ž
y 97.1 " 2y.71 kJ mol , y 109.2 " 2.9 kJ mol , y 106.1 " 3.0 kJ mol , and y 113.2 "
.
.
2.7 kJ mol
, respectively. In combination with the enthalpies of formation of the
corresponding alkanonyes1, we obtained the enthalpies of formation ⌬f HTm cyclopentenone s
Ž
.
and ⌬f HTm cyclohexenone s y 121.9 " 3.0 kJ moly1. By using
.
.
Ž
.
Ž
.
Ž
.
y 100.3 " 3.0 kJ mol
Ž
isodesmicy1reactions, the resonance energy of cyclic enones is estimated to be 13 "
.
.
4 kJ mol
. ᭧ 1998 Academic Press
KEYWORDS: enthalpy; hydrogenation; enone; resonance energy; calorimetry
1. Introduction
As part of our recent study on the structure and energetics of enynones, dienones,
and related species,Ž1. we carried out measurements of the molar enthalpies of
hydrogenation ⌬hyd Hm at T s 298.15 K of 2-cyclopentenone, 2-cyclohexenone, and
a
Ž
.
To whom correspondence should be addressed E-mail: drogers@phoenix.liu.edu .
0021᎐9614r98r111393 q 08 $30.00r0
᭧ 1998 Academic Press

1394
D. W. Rogers et al.
TABLE 1. Molar enthalpies of hydrogenation ⌬hyd Hm at the temperature 298.15 K of four enones; r
Ž
.
Ž
.
Ž
.
Ž
.
denotes n cyclohexene rn enone or n 1-hexene rn enone
a
²
:
y⌬hyd Hm
y⌬hyd Hm
r
y1
y1
.
.
kJ mol
kJ mol
Ž
.
Cyclopentenone mass fraction, 0.994
0.8122
0.9015
98.7 " 1.3
97.5 " 1.7
98.1 " 1.7
97.1 " 2.7
b
Ž
.
.
.
.
Cyclohexenonec mass fraction, 0.996
Ž
.
0.8448
0.8545
108.3 " 1.1
112.2 " 2.7
110.2 " 1.9
109.2 " 2.9
b
b
b
Ž
Ž
.
2-Methyl-3-isobutoxycyclohex-2-enone mass fraction, 0.99
0.8710
0.8872
109.2 " 0.8
106.7 " 2.9
107.1 " 2.0
106.1 " 3.0
Ž
Ž
.
5-Methylhex-5-en-2-one mass fraction, 0.995
0.8197
0.8320
0.8180
113.0 " 2.1
115.5 " 1.7
114.6 " 1.7
114.2 " 1.7
113.2 " 2.7
Ž
a Mean value of nine measurements of ⌬hyd Hm . b Mean value corrected for ⌬sol Hm at T s 298 K.
c Mixed solvent mole fraction 0.616 of tetrahydrofuran and n-hexane see text .
Ž
.
2-methyl-3-isobutoxycyclohex-2-enone, members of a class of compounds not well
represented in the thermochemical literature.Ž2. Solvent effects on ⌬hyd Hm were
also examined, including comparison of a solution-phase ⌬hyd Hm of 5-methylhex-5-
en-2-one with a gas-phase study of an alkane of similar size, 2,2,4-trimethylpent-
1-ene.
2. Experimental
The compounds 2-cyclopenten-1-one, 2-cyclohexen-1-one, cyclopentanone, and
Ž
.
cyclohexanone are commercially available Aldrich, Acros . The enones were
fractionally distilled and stored as refrigerated samples of mass 1.0 g sealed in glass
Ž
.
under argon. Sample purity table 1 was determined by g.l.c. The cycloalkanones
Ž
.
were very pure mass fraction f 0.998 as obtained. Synthetic methods for obtaining
3-ethynyl-2-cycloalkenones have been given previously.Ž3.
The calorimeter and method have been described elsewhere.Ž4,5. Briefly, portions
Ž
.
of a dilute reactant unsaturate solution were injected into a reaction flask
containing 18.0 cm3 of the same solvent Aldrich, mass fraction ) 0.99 as the
Ž
.
Ž
.
Ž
.
reactant solution and 300 mg of 0.05 mass fraction of Pd, C catalyst Aldrich , and
stirred magnetically. Hydrogen at a pressure of 0.1 MPa above ambient was
conveyed into the flask through a hypodermic needle. The reaction flask contained
a thermistor connected as one arm of a Wheatstone bridge circuit which produces
Ž
a voltage that is amplified and fed into an ArD converter Alpha Products, Darien,
.
CT 06820 , and thence to a microcomputer. The microcomputer is programmed to

⌬
hyd Hm and ⌬f Hm of enones
1395
treat the temperature against time curve obtained during hydrogenation by
Ž
.
suitable extrapolations of bridge output proportional to temperature taken before
and after hydrogenation, in a manner analogous to the extrapolations used in
conventional combustion calorimetry.Ž6. Each of the extrapolations, forward and
back, intersects with a vertical passing through the inflection point of the T against
t curve. The length of the line segment between these intersections is proportional
to ⌬T for hydrogenation.
Identical volumes of a dilute standard solution of an alkene of known ⌬hyd Hm ,
and a similar solution of an unknown were alternately injected into the calorimeter
under as nearly as possible identical conditions. The ratio of ⌬T values for
hydrogenation of each reaction pair led to ⌬hyd Hm at T s 298.15 K of the
unknown.
Ž
.
Depending on the molar mass, 60 to 80 mg of cycloalkenone was weighed to
y6
.
"2 10 g on a Sartorius M3P electronic microbalance and diluted with solvent to
a volume of 0.500 cm3. The standard was made up so as to be thermochemically
equivalent to the sample, i.e. it produced approximately the same amount of heat
upon hydrogenation as the sample solution. This amount was known for
cyclopentenone and cyclohexenone from preliminary experiments on commercial
Ž
.
product Aldrich of mass fractions 0.98 and 0.97, respectively. Approximate
Ž
.
Ž
.
thermochemical equivalence required n standard rn enone of 0.8 to 0.9, as seen
in table 1. Portions of standard and sample 0.0200 cm were injected into the
calorimeter by using a g.l.c. microsyringe fitted with a Cheney adaptor for maximum
reproducibility.
3
Ž
.
Ž
.
Two thermochemical standards Wiley Organics; mass fraction 0.999 were
used in this work. Results in table 1 were calculated from ⌬hyd Hm values at T s
y1 Ž7a,b.
.
Ž
.
Ž
or 1-hexene, y 126.6 "
298.15 K of cyclohexene, y 118.8 " 0.8 kJ mol
,
y1 Ž7c.
.
.
0.8 kJ mol
.
The value selected for cyclohexene is the arithmetic mean of
Ž7b.
y1
.
Ž
.
⌬
hyd Hm s y 119.0 " 0.3 kJ mol
in the gas phase at T s 355 K, which would
be expected to be slightly more negative than the result at T s 298.15 K, and
y1
.
Ž
.
⌬
hyd Hm sy 118.6 " 0.2 kJ mol
measured at T s 298.15 K and corrected for
solvent effects.Ž7a. These results have been discussed recently by Steele et al.Ž7d. An
early result of ours, ⌬hyd Hm sy 112.5 " 2.5 kJ moly1, should be discarded.Ž8.
Solvents used were n-propanol, or tetrahydrofuran, or a mixed solvent system of
.
Ž
.
Ž
.
tetrahydrofuran in n-hexane mol fraction tetrahydrofuran, 0.616 . Enthalpies of
solution of the reaction products were measured by the same procedure as above.
The apparatus was not designed to measure small negative temperature changes
but the results showed that the heat effect superimposed on ⌬hyd Hm at T s
y1
298.15 K was negative and not greater than 2 kJ moly1, i.e. 1 " 2 kJ mol
.
.
.
Ž
.
Accordingly, the results in table 1 have been adjusted to less negative values by
1 kJ moly1, and the experimental uncertainties of the enthalpies and the resonance
.
energies derived from them have been increased to include a ⌬sol Hm contribution
to the uncertainty of the final results.
Ž5a.
Ž
.
The calorimeter was operated at ambient temperature isoperibol operation .
The temperature difference between isoperibol operation and isothermal operation
at T s 298.15 K was within 3 K. Straightforward thermodynamic calculations show

1396
D. W. Rogers et al.
that the effect of isoperibol operation at temperatures that were slightly different
from T s 298.15 K does not influence the measured ⌬hyd Hm by a significant
amount. We believe that the measured values in table 1 are the same as the
⌬
hyd Hm at T s 298.15 K to within a tolerance that is a factor of 10 smaller than
the stated experimental uncertainties.
3. Results and discussion
The uncertainties in table 1 are standard deviations from the mean for nine pairs
Ž
.
of hydrogenations, alternating sample with standard 18 hydrogenations in all .
They express approximately 95 per cent confidence limits. Entries in parentheses
y1
.
have been corrected by 1.0 kJ mol
and their uncertainty estimate has been
increased as described above due to uncertainty in the measurement of ⌬sol Hm at
T s 298.15 K.
Thermograms were normal for the compounds in table 1, and reaction times
were less than the 10 s instrument system response time. Examination by g.l.c.
Ž
.
60 m wide bore capillary, SE 30, Supelco of the reaction product showed no
evidence of incomplete or competing reactions.
There was a small enthalpy effect superimposed on the reaction enthalpy for
the first few hydrogenations. We ascribe it to adsorption of the polar enone on the
catalyst or its carbon support. The initial results were more negative than the
arithmetic mean of the entire set by 2 to 6 kJ moly1. The exothermic effect
rapidly diminished, and subsequent results held steady at the values given in
table 1. Accordingly, 15 pairs of experimental measurements were made, the last
nine of which are reported as individual entries in table 1. We have encountered
this effect before and we handled it in a similar way.Ž9. In a mixed solvent system,
.
Ž
.
Ž
.
tetrahydrofuran q n-hexane , the exothermic effect was not observed.
Enthalpies of solution of the reaction products, cyclopentanone and
cyclohexanone, were also measured for the mixed solvent system, and ⌬sol Hm was
estimated as - 2.0 " 2.0 kJ moly1. We have already mentioned that the
calorimeter was not designed to measure such small enthalpy effects, which
accounts for the large relative uncertainty. The influence of solvent on the
measured ⌬hyd Hm at T s 298.15 K was also assessed by comparing our value of
.
Ž
.
Ž .
hyd Hm 5-methylhex-5-en-2-one , in which one expects no influence of the keto
⌬
oxygen on the rather distant double bond, with the gas-phase result for a comparable
alkene, 2,2,4-trimethylpent-1-ene.
Resonance is
a theoretical construct usually thought of as describing
intramolecular electronic interactions within isolated molecules. Therefore,
resonance energy should properly be measured on samples in the ideal gas state.
This is not feasible for any real samples, though it may be well approximated under
some circumstances. More to the point, hydrogenation thermochemistry has not
been carried out on real gases for many years owing to the low volatility of most
samples of interest. In this work, in which a dilute solution of, for example,
Ž
.
alkenone in n-propanol f 0.1 mass fraction is hydrogenated to a very dilute
y4
.
Ž
.
mass fraction , we argue that the solution-phase
solution of alkanone f 1 10

⌬
hyd Hm and ⌬f Hm of enones
1397
y1
.
and gas-phase ⌬hyd Hm differ by less than the f 2 kJ mol error limits assigned
y1
.
to ⌬f Hm at T s 298.15 K, and the 4 kJ mol assigned to the resonance energy.
Following Fuchs and Peacock,Ž10. we note that, for a fixed amount of neat
Ž
.
Ž .
reactant, the enthalpy of hydrogenation in the gas phase ⌬hyd Hm g is given by:
⌬
hyd Hm g s ⌬hyd Hm sln q ⌬trans Hm reactant y ⌬trans Hm product ,
1
Ž . Ž .
Ž
.
Ž
.
Ž
.
Ž
.
where ⌬hyd Hm sln is the enthalpy of hydrogenation in solution, and ⌬trans Hm is
the enthalpy of transfer of solute from the solution to the gas phase. Hydrogen is
rapidly replaced from the gas phase during hydrogenation of the sample, both in
solution and on the catalyst surface, so that the overall solution and
desorption᎐adsorption effect for H 2 is negligible. Each enthalpy of transfer is the
algebraic sum of ⌬vap Hm and ⌬sol Hm , the enthalpies of vaporization and solution
of the reactant and product, respectively. Now,
⌬
hyd Hm g s ⌬hyd Hm sln q ⌬vap Hm product y ⌬vap Hm reactant y
Ž .
Ž
.
Ž
.
Ž
.
⌬
sol Hm product q ⌬sol Hm reactant .
2
Ž .
Ž
.
Ž
.
We do not know ⌬vap Hm for all of the reactants and products in table 1, but we
draw an analogy between unsaturation in ketones and unsaturation in alcohols.
y1
Polak and BensonŽ11. found that ⌬vap Hm prop-2-en-1-ol s 47.3 " 1.3 kJ mol
.
Ž
.
Ž
.
y1
.
Ä
Ž .
4
Ž
.
and ⌬vap Hm propanol l s 47.5 " 0.04 kJ mol
at T s 298.15 K. Wadso obs-
¨
y1 Ž12.
.
Ä
Ž .
4
Ž
.
erved the value ⌬vap Hm propanol l to be 47.3 " 0.1 kJ mol
.
Comparison of
the saturated and unsaturated alcohols indicates that the double bond has little
influence on ⌬vap Hm , a conclusion supported by our previous and more general
work.Ž13.
Enthalpies of solution are largely determined by the size of the cavity in the
solvent system made by intrusion of solute moleculesŽ14. which, in turn, correlates
directly with the size, shape, and functionality of the solute.Ž13. For hydrogenation
of the two simple cycloalkenones in table 1, product and reactant are almost
exactly the same size and have a similar electron distribution, differing only in the
presence or absence of the double bond. In the light of these observations, we
Ž .
believe that the last four terms in equation 2 nearly cancel pairwise, leaving
Ž .
Ž
.
⌬
hyd Hm g f ⌬hyd Hm sln within the limits of the other sources of experimental
error. It should be noted that the actual process occurring along with hydrogenation
in the calorimeter is an enthalpy change of the system due to dilution of a f 0.15
mass fraction alkenone solution to a very dilute alkanone solution. This thermal
effect can be assumed to be very much smaller than the difference between
Ž
.
Ž
.
⌬
sol Hm reactant and ⌬sol Hm product , hence the arguments in this section follow
a fortiori.
A factor that has not yet been accounted for is the possibility of association
between the solvent and solute, for example, the polar OH group of n-propanol
and the double bond in the reactant. Such an association would be broken during
hydrogenation, and the enthalpic effect would depend upon the strength of the
reactant ᎐solvent association. This enthalpy would, if present, appear as a difference

1398
D. W. Rogers et al.
Ž
.
Ž
.
in ⌬sol Hm reactant and ⌬sol Hm product in the solvent chosen as the calorimeter
fluid. At present, we do not have an accurate direct measure of this strength, but
Roth and LennartzŽ15. have measured ⌬sol Hm for cyclopentene, cyclopentane,
cyclohexene, and cyclohexane and found them to differ by slightly more than
y1
y1
.
.
1 kJ mol in methanol and by less than 1 kJ mol in cyclohexane.
In order to judge the systematic error brought about by any combination of the
Ž
.
factors described above, we measured ⌬hyd Hm 5-methylhex-5-en-2-one , in which
the keto oxygen and the ␣-methyl double bond are well separated from one
another and should have no mutual intramolecular energetic influence. We
Ž
.
compare our measurement with the gas-phase ⌬hyd Hm 2,2,4-trimethylpent-1-ene
obtained by Kistiakowsky.Ž7b. The results are ⌬hyd Hm s y 113.2 " 1.7 kJ mol
y1
.
Ž
.
and ⌬hyd Hm s y 113.8 " 0.4 kJ moly1, respectively, indicating no significant
difference and no significant solvent error. For completeness, we note that the
.
Ž
.
Ž
temperatures at which these experiments were carried out were different T s
.
298.15 K and T s 355.15 K, respectively , that ⌬hyd Hm of a C7 ␣-methyl terminal
alkene has not been measured in the gaseous phase, and that ⌬hyd Hm g for the C6
Ž .
y1
.
and C5 ␣-methylalkenes at T s 355.15 K are larger, y117 kJ mol compared with
y119 kJ moly1, than the compound we chose as the most apposite for comparison.
.
The standard molar enthalpies of formation in the gas phase at T s
Ž16.
298.15 K, ⌬f HTm cyclopentanone s y 197.4 " 1.3 kJ mol
and ⌬f HTm cyclo-
Ž
y1
.
Ž
.
Ž
.
Ž16.
y1
.
.
Ž
.
Ž
.
hexanone s y 231.1 " 0.8 kJ mol
lead to ⌬f Hm cyclopentenone s
y1
y1
.
.
Ž
.
Ž
.
Ž
.
y 100.3 " 3.0 kJ mol and ⌬f Hm cyclohexenone s y 121.9 " 3.0 kJ mol at
T s 298.15 K, the two main target compounds of this study. Uncertainties in ⌬f HTm
are given as the root-mean-square of the experimental uncertainties. We lack the
necessary ⌬f HTm values of the hydrogenation products for calculating ⌬f HTm of the
remaining two compounds in table 1 at T s 298.15 K.
4. Resonance energies of cylopentenone and cyclohexenone
We are in a position to evaluate the resonance energies of these cyclic enones.
Following from reference 2, the enthalpy of the following isodesmic reaction:
cycloalkenone q cycloalkane s cycloalkene q cycloalkanone,
3
Ž .
may be equated with the resonance energy of the enone. Inspection shows that this
enthalpy equals the difference in enthalpies of hydrogenation of the enone and
the parent olefin. For cyclopentenone, taking the enthalpy of hydrogenation of
y1
.
Ž
.
cyclopentene as y 112.5 " 1.7 kJ mol from reference 17, we deduce a resonance
y1
.
Ž
.
Ž .
energy of 15.4 " 3.2 kJ mol . Likewise, for cyclohexenone, from references 7 a
and 7 b , we deduce a resonance energy of 9.6 " 3.0 kJ moly1. An arithmetic
.
Ž .
Ž
.
y1
.
Ž
.
mean of the resonance energy of 13 " 4 kJ mol
for cyclic enones is thus
suggested.
There are disappointingly few species with which comparisons can be made.Ž2.
Ž .
The first is with E -2-butenal for which the gas-phase enthalpy of hydrogenation
y1
of y 104.2 " 0.4 kJ mol
has been reported.Ž18. That of the corresponding
.
Ž
.

⌬
hyd Hm and ⌬f Hm of enones
1399
y1 Ž19.
.
Ž .
Ž
.
olefin, E -2-butene, has been measured as y 114.6 " 0.6 kJ mol
,
resulting
in a resonance energy of 10.2 " 0.6 kJ moly1. The enthalpies of combustion, and
.
Ž
.
Ž .
thus of formation, of E -2-cycloheptadecenone and cycloheptadecanone have
been determined,Ž20. and corresponding sublimation enthalpies either directly
measured,Ž21. or derived.Ž22. The resulting enthalpies of formation of the gas-
y1
and y 408.3 " 13.4 kJ moly1, can
Ž .
.
.
Ž
.
eous ketones, y 460.3 " 10.9 kJ mol
be numerically combined to derive an enthalpy of hydrogenation of the 2-
y1
.
Ž
.
cycloheptadecenone, y 52.0 " 17.3 kJ mol
. We lack thermochemical
Ž .
information about E -cycloheptadecene but, by analogy with other cycloalkenes
with more than seven carbon atoms, or long acyclic olefins with double bonds that
are not in the terminal position,Ž23. we believe that its hydrogenation enthalpy will
Ž
.
not be very different from that of these species, approximately y 110 to 120
kJ moly1. Used without reservation, we would conclude that 2 -cycloheptadecenone
.
has a resonance energy of some 50 to 60 kJ moly1. We are very doubtful of this
last value, but suspect no error except in the measurement of the enthalpy of
formation of either large-ring ketone.
.
Ž
.
D.W.R wishes to thank the trustees of Long Island University and the release time
committee for a grant of release time for research. Acknowledgment is made by
M.H. to the donors of the Petroleum Research Fund, administered by the American
Chemical Society, the Research Corporation, and the National Science Foundation.
We thank the Chemical Science and Technology Laboratory of the U.S. National
Institute of Standards and Technology for partial support of J.F.L.’s research.
REFERENCES
1. Traetteberg, M.; Liebman, J. F.; Hulce, M.; Bohn, A. A.; Rogers, D. W. J. Chem. Soc. Perkin 2,
1997, 1925᎐1932.
2. Liebman, J. F.; Pollack, R. M. The Chemistry of Functional Groups: the chemistry of enones. Patai, S.;
Rappoport, Z.: editors. Wiley: Chichester. 1989, Ch. 4.
Ž .
Ž .
b
3.
a
Cheng, M.; Hulce, M. J. Org. Chem. 1990, 55, 964᎐975;
Lee, S-H.; Hulce, M. Tetrahedron
Ž .
Lett. 1990, 31, 311᎐314;
c Capon, B. The Chemistry of the Functional Groups: enones. Patai, S.;
Rappoport, Z.: editors. Wiley: New York. 1989.
Ž .
4. a Fang, W.; Rogers, D. W. J. Org. Chem. 1992, 57, 2294᎐2297;
Ž .
b
Rogers, D. W.; Crooks, E. L.;
Ž .
c Rogers, D. W.; Dagdagan,
Dejroongruang, K. J. Chem. Thermodynamics 1987, 19, 1209᎐1215;
O. A.; Allinger, N. L. J. Amer. Chem. Soc. 1979, 101, 671.
Ž .
5. a Hemminger, W.; Hohne, G. Calorimetry, Fundamentals and Practice. Verlag Chemie: Weinheim.
Ž . Ž .
1984, pp. 83-84;
b Rogers, D. W. Amer. Lab. 1980 I , 18.
6. Sime, R. J. Physical Chemistry. Methods, techniques and experiments. Saunders: Philadelphia. 1990, p.
422.
Ž .
Ž .
b
7. a Roth, W. R.; Lennartz, H.-W. Chem. Ber. 1980, 113, 1806᎐1817;
T. L.; Kistiakowsky, G. B.; Vaughan, W. E. J. Amer. Chem. Soc. 1937, 59, 831᎐841;
Ž .
Dolliver, M. A.; Gresham,
Ž .
c
Rogers,
D. W. J. Phys. Chem. 1979, 83, 2430᎐2431;
d Steele, W. V.; Chirico, R. D.; Knipmeyer, S. T.;
Nguyen, A; Smith, N.K.; Tasker, I.R. J. Chem. Eng. Data 1996, 41, 1268᎐1284.
8. Rogers, D. W.; McLafferty, F. J. Tetrahedron 1971, 27, 3765᎐3775.
Ž .
Ž .
b Rogers, D. W.; Hoyte,
9.
a
Rogers, D. W.; Siddiqui, N. A. J. Phys. Chem. 1975, 79, 574᎐577;
Ž .
O. P. A.; Ho, R. K. J. Chem. Soc., Faraday I 1978, 74, 46᎐52; c Rogers, D. W.; Choudhury, D. N.
J. Chem Soc., Faraday I 1978, 74, 2868᎐2872.
10. Fuchs, R.; Peacock, L. A. J. Phys. Chem. 1979, 83, 1975᎐1978 and references therein.
11. Polak, J.; Benson, C. G. J. Chem Thermodynamics 1971, 3, 235᎐242.
12. Wadso. I. Acta Chim. Scand. 1966, 20, 544᎐552.

1400
D. W. Rogers et al.
13. Chikos, J. S.; Hesse, D. G.; Liebman, J. F.; Panshin, S. Y. J. Org. Chem. 1988, 53, 3424᎐3429.
14. Saluja, P. P. S.; Young, T. M.; Rodewald, R. F.; Fuchs, F. H.; Kohli, D.; Fuchs, R. J. Am. Chem.
Soc. 1977, 99, 2949᎐2953.
15. Roth, W. A.; Lennartz, H.-W. Chem. Ber. 1980, 113, 1806᎐1817.
16. Wilberg, K. B.; Crocker, L. S.; Morgan, K. M. J. Am. Chem. Soc. 1991, 113, 3447᎐3450.
17. Allinger, N. L.; Dodzuik, H.; Rogers, D. W.; Naik, S. N. Tetrahedron 1982, 38, 1593᎐1597.
18. Dolliver, M. A.; Gresham, T. L.; Kistiakowsky, G. B.; Smith, E. A.; Vaughan, W. E. J. Am. Chem.
Soc. 1938, 60, 440᎐450.
19. Kistiakowsky, G. B.; Ruhoff, J. R.; Smith, H. A.; Vaughan, W. E. J. Am. Chem. Soc. 1935, 57,
876᎐882.
20. Ruzicka, L.; Schlapfer, P. Hel¨. Chim. Acta 1933, 16, 162᎐168.
21. Wolf, K. L.; Weghofer, H. Z. Phys. Chem. 1938, B39, 194᎐208.
22. Cox, J. D.; Pilcher, G. Thermochemistry of Organic and Organometallic Compounds. Academic:
London. 1970.
23. Slayden, S. W.; Liebman, J. F. The Chemistry of Functional Groups Supplement A3: the chemistry of
double-bonded functional groups. Patai, S.: editor. Wiley: Chichester. 1997, Ch. 11.
(
)
Recei¨ed 25 March 1998; in final form 22 June 1998
0-720