Is benzene oxide homoaromatic? A microcalorimetric study

Article abstract of DOI:10.1021/ja010226h

Rate constants and heats of reaction for the aromatization of benzene oxide (1) and the acid-catalyzed aromatization of benzene hydrate (2) in highly aqueous solution giving phenol and benzene, respectively, have been measured by heat-flow microcalorimetry. The measured heat of reaction of benzene oxide, ΔH = -57.0 kcal mol^-1, is much larger than that of benzene hydrate, ΔH = -38.7 kcal mol^-1, despite an unusually low reactivity of benzene oxide, rate ratio 0.08. The measured enthalpies agree with those calculated using the B3LYP hybrid functional corrected with solvation energies derived from semiempirical AM1/SM2 calculations. Comparison with the measured enthalpies of the corresponding reactions of the structurally related 1,3-cyclohexadiene oxide (3) and 2-cyclohexenol (4) of ΔH = -24.9 kcal mol^-1 (includes a small calculated correction of - 1.2 kcal mol^-1) and ΔH ～ 0 kcal mol^-1, respectively, gives a smaller aromatization energy for the benzene oxide than for the benzene hydrate reaction (ΔΔΔH = 6.6 kcal mol^-l). This suggests that benzene oxide is unusually stabilized by a significant amount of homoaromatization as has been proposed previously (J. Am. Chem. Soc. 1993, 115, 5458). This unusual stability accounts for more than half of the ～10⁷ times lower than expected reactivity of benzene oxide toward acid-catalyzed isomerization. The rest is suggested to originate from an unusually high energy of the carbocation-forming transition state.

Full text of DOI:10.1021/ja010226h

VOLUME 123, NUMBER 42
OCTOBER 24, 2001
© Copyright 2001 by the
American Chemical Society
Is Benzene Oxide Homoaromatic? A Microcalorimetric Study†
Zhi Sheng Jia, Peter Brandt,‡ and Alf Thibblin*
Contribution from the Institute of Chemistry, UniVersity of Uppsala, Box 531, S-751 21 Uppsala, Sweden
ReceiVed January 25, 2001. ReVised Manuscript ReceiVed April 15, 2001
Abstract: Rate constants and heats of reaction for the aromatization of benzene oxide (1) and the acid-catalyzed
aromatization of benzene hydrate (2) in highly aqueous solution giving phenol and benzene, respectively,
have been measured by heat-flow microcalorimetry. The measured heat of reaction of benzene oxide, ∆H )
-57.0 kcal mol-1, is much larger than that of benzene hydrate, ∆H ) -38.7 kcal mol-1, despite an unusually
low reactivity of benzene oxide, rate ratio 0.08. The measured enthalpies agree with those calculated using the
B3LYP hybrid functional corrected with solvation energies derived from semiempirical AM1/SM2 calculations.
Comparison with the measured enthalpies of the corresponding reactions of the structurally related
1,3-cyclohexadiene oxide (3) and 2-cyclohexenol (4) of ∆H ) -24.9 kcal mol-1 (includes a small calculated
correction of -1.2 kcal mol-1) and ∆H ∼ 0 kcal mol-1, respectively, gives a smaller aromatization energy for
the benzene oxide than for the benzene hydrate reaction (∆∆∆H ) 6.6 kcal mol-1). This suggests that benzene
oxide is unusually stabilized by a significant amount of homoaromatization as has been proposed previously
(J. Am. Chem. Soc. 1993, 115, 5458). This unusual stability accounts for more than half of the ∼107 times
lower than expected reactivity of benzene oxide toward acid-catalyzed isomerization. The rest is suggested to
originate from an unusually high energy of the carbocation-forming transition state.
Introduction
hydrocarbons to phenols in biological systems has been proposed
to occur via arene oxides. This has been strongly supported by
the observation of naphthalene oxide.1 Aromatic compounds also
undergo metabolism through arene hydrates, but this route is
probably of less importance.2,4
It has been proposed by More O’Ferrall and co-workers that
arene oxides may be stabilized by homoaromaticity.5 The low
reactivity of benzene oxide suggests a stabilization of at least
10 kcal mol-1. This issue is of great interest in connection with
studies of liver enzyme metabolism of aromatic compounds and
their ability to provide dihydrodiols and glutathione adducts.3
The proposed homoaromaticity of arene oxides is of course also
of interest from a purely theoretical point of view in connection
with the general concept of aromaticity.6
Arene oxides and arene hydrates are reactive compounds of
considerable biochemical interest.1-4 The oxidation of aromatic
† Presented in part at the 15th IUPAC Conference on Physical Organic
Chemistry, Gothenburg, Sweden, 2000.
‡ Present address: Organic Chemistry, Department of Chemistry, Royal
Institute of Technology, SE-100 44 Stockholm, Sweden.
(1) Lehr, R. E.; Wood, A. W.; Levin, W.; Conney, A. H.; Jerina, D. M.
In Polycyclic Aromatic Hydrocarbon Carcinogenesis: Structure ActiVity
Relationships; Yang, S. K., Silverman, B. D., Eds.; CRC Press Inc.: Boca
Raton, FL, 1988; p 31.
(2) (a) Boyd, D. R.; McMordie, R. A. S.; Sharma, N. D.; Dalton, H.;
Williams, P.; Jenkins, R. O. J. Chem. Soc., Chem. Commun. 1989, 339. (b)
Boyd, D. R.; Hand, M. V.; Sharma, N. D.; Chima, J.; Dalton, H.; Sheldrake,
G. N. J. Chem. Soc., Chem. Commun. 1991, 1630. (c) Boyd, D. R.; Dorrity,
M. R. J.; Hand, M. V.; Malone, J. F.; Sharma, N. D.; Dalton, H.; Gray, D.
J.; Sheldrake, G. N. J. Am. Chem. Soc. 1991, 113, 666.
(3) Jerina, D. M.; Daly, J. W.; Wikop, B.; Zaltzman-Nirenberg, P.;
Udenfriend, S. Biochemistry 1970, 9, 147.
(4) Boyd, D. R.; Jerina, D. M. Oxiranes, Arene Oxides, Oxaziridines,
Thietanes, Thietes, Thiazeles. In Small Ring Heterocycles-Part III; Hassner,
A., Ed.; John Wiley & Sons: New York, 1985; p 197.
(5) (a) More O’Ferrall, R. A.; Rao, S. N. Croat. Chem. Acta 1992, 65,
593. (b) Rao, S. N.; More O’Ferrall, R. A.; Kelly, S. C.; Boyd, D. R.;
Agarwal, R. J. Am. Chem. Soc. 1993, 115, 5458.
(6) Krygowski, T. M.; Cyran˜ski, M. K.; Czarnocki, Z.; Ha¨felinger, G.;
Katritzky, A. R. Tetrahedron 2000, 56, 1783.
10.1021/ja010226h CCC: $20.00 © 2001 American Chemical Society
Published on Web 10/02/2001

10148 J. Am. Chem. Soc., Vol. 123, No. 42, 2001
Jia et al.
Table 1. Rate Constants and Heats of Reaction for the Reactions
of 1-4 at 25 °C
Chart 1
106kobs
,
kH,
P0,
∆H,
∆∆H,
solventa
s-1
M-1 s-1 µW
kcal mol-1 kcal mol-1
Benzene Oxide (1)
glyc-waterb
glyc-waterc
535 ( 15 ∼22
169
-57.0 ( 1.9
-18.3
589 ( 15
Benzene Hydrate (2)
Scheme 1
glyc-waterd
glyc-watere
glyc-waterf
338
346
126
129
155 ( 2
58
1.7
43.5 -38.7 ( 0.8
55.7 -39.9 ( 1.5
MeCN-waterg 246 ( 4
1,3-Cyclohexadiene Oxide (3)
MeCN-waterh 173 ( 4
13
-23.7 ( 0.6
-24.9j
2-Cyclohexenol (4)
∼10-3
wateri
∼0
a The buffer concentrations and pH values: before mixing with the
organic cosolvent. b 25 vol % glycerol in water, 0.1 M phosphate buffer
pH 7.00. c 25 vol % glycerol in water, 0.1 M phosphate buffer pH 5.57.
d 25 vol % glycerol in water, 0.1 M acetate buffer pH 5.57; kinetics
were studied by UV spectrophotometry. e 25 vol % glycerol in water,
0.1 M phosphate buffer pH 5.57. f 50 vol % glycerol in water, 0.1 M
acetate buffer pH 5.57. g 50 vol % acetonitrile in water, 0.1 M acetate
buffer pH 3.85. h 50 vol % acetonitrile in water, 0.1 M phosphate buffer
pH 5.99. i Reference12. j After correction for dehydration (see text and
Scheme 2).
We now report results on the aromatization reactions of ben-
zene oxide (1) and benzene hydrate (2) under acidic conditions.
These reactions have previously been concluded to be subject
to specific acid catalysis with rate-limiting formation of an
R-hydroxycyclohexadienyl carbocation intermediate and a cy-
clohexadienyl carbocation intermediate, respectively.5,7 Surpris-
ingly, benzene hydrate was found to be 12 times more reactive
(after statistical correction for the two carbon-oxygen bonds
in 1) than benzene oxide despite acid-catalyzed cleavage of
carbon-oxygen bonds in simple epoxides occurring 106-107
times faster than in the structurally related hydrates.5 This is
illustrated in Chart 1 which gives the previously reported second-
order rate constants for the acid-catalyzed reactions measured
in aqueous solution.5
Scheme 2
The aim of our study was to test the hypothesis that the
unusually low reactivity of benzene oxide compared to benzene
hydrate is due to homoaromatic stabilization of the benzene
oxide reactant. The experimental results obtained by heat-flow
microcalorimetry and the results of calculation do support this
proposal. However, transition-state effects are also of impor-
tance.
M-1 s-1 at 25 °C. The uncatalyzed path has a rate constant of
k0 ∼ 0.53 × 10-3 s-1. The rate constants reported in Table 1
were measured at relatively high pH values, and the main part
of the reaction of benzene oxide is concluded to occur through
the uncatalyzed pathway. Microcalorimetric kinetic studies at
lower pH were not possible due to the large time constant of
the microcalorimeter.
Results
The dehydration reaction of 2 yields benzene as the sole
product (Scheme 1). The results, which were obtained by the
same methods as described above and by UV spectrophotom-
etry, are reported in Table 1. Due to difficulties in measuring
the benzene product concentration accurately in 25 vol %
glycerol, a larger fraction of the organic cosolvent was used. A
rate constant of 190 M-1 s-1 has previously been reported for
the reaction in water at the same temperature.5b
The reactions produce large amounts of heat and are easy to
follow by heat-flow microcalorimetry. The reaction heats for
the reactions of 1 and 2 were measured as -57.0 and -38.7
kcal mol-1, respectively. The details of the heat-flow micro-
calorimetric measurements are given in the Experimental
Section.
The substrates benzene oxide (1) and benzene hydrate (2)
were prepared from 1,4-cyclohexadiene following the procedures
of Gillard et al.8 and Stavascik et al.,9 respectively. 1,3-Cyclo-
hexadiene oxide (3) was prepared from 1,3-cyclohexadiene.10
The cleavage of 1 in 25 vol % glycerol in water at 25 °C
yields phenol as the sole product (Scheme 1). The reaction was
followed by a sampling high-performance liquid chromatogra-
phy procedure and by heat-flow microcalorimetry. The results
are presented in Table 1. The kinetics of the reaction has been
studied previously in aqueous 1 M KCl at 30 °C by Kasperek
and Bruice.7 They found that an uncatalyzed path (k0 ) 1.4 ×
10-3 s-1) accompanies the specific acid-catalyzed path (kH
)
32 M-1 s-1) according to
The cleavage reaction of 3 gives a mixture of diols with the
trans 1,2-diol as the main product (Scheme 2). The rate constant
and the enthalpy were measured in 50 vol % aqueous acetonitrile
(Table 1); the reaction was too fast in 25 vol % glycerol in
water for measurement by the microcalorimeter. The rate
constants and product composition of diols have been studied
previously in aqueous solution at 25 °C by Whalen; an
uncatalyzed pathway was found to compete with a specific acid-
catalyzed reaction.11 To obtain the reaction heat for the
production of the monohydroxy substituted compound (Scheme
kobs ) k0 + kH[H+]
(1)
A crude estimate of the second-order rate constant, based
upon the observed rate constants given in Table 1, is kH ∼ 22
(7) Kasperek, G. J.; Bruice, T. C. J. Am. Chem. Soc. 1972, 94, 198.
(8) Gillard, J. R.; Newlands, M. J.; Bridson, J. N.; Burnell, J. Can. J.
Chem. 1991, 69, 1337.
(9) Stavascik, J.; Rickborn, B. J. Am. Chem. Soc. 1971, 93, 3046.
(10) Ramesh, K.; Wolfe, M. S.; Lee, Y.; Velde, D. V.; Borchard, R. T.
J. Org. Chem. 1992, 57, 5861.

Is Benzene Oxide Homoaromatic?
J. Am. Chem. Soc., Vol. 123, No. 42, 2001 10149
Scheme 3
stability of their carbocation-forming transition states; reactant
stability is of less importance.13,14 The mechanism of dehydra-
tion of 2 is shown in Scheme 3. The proposed mechanism of
acid-catalyzed epoxide ring opening is very similar; i.e., the
rate-limiting step involves cleavage of a carbon-oxygen bond
to provide a carbocation.5,7 The carbocation intermediate is
destabilized to some extent by the â-hydroxy substituent, but
this effect is generally small compared to the gain in energy
brought about by the release of strain from the cleavage of the
epoxide ring. This mechanism for 1 is shown in Scheme 3. The
carbocation 1+ may undergo a hydride shift (NIH shift) before
the final aromatization step takes place.7
Epoxide ring opening may, in principle, also occur through
an SN2 mechanism. However, it has been pointed out that this
is not reasonable for substrates producing stabilized carbocations
in the presence of nucleophiles not stronger than water.5 This
is discussed in more detail at the end of the Discussion. There
is also direct evidence of the cation character of the transition
state of epoxide ring-opening reactions.7,16-18 Accordingly, the
result of our calculations is that the reaction rates are closely
related to the stability of the carbocations (vide infra).
Unusually Low Reactivity of Benzene Oxide (1). The
kinetic data of the reactions of 1 and 2 in 25 vol % glycerol in
water are in accord with the results previously reported for the
reactions under slightly different conditions.5b,7 The acid-
catalyzed reaction of 2 is 12-fold faster than the corresponding
reaction of 1 (after statistical correction, Chart 1). This was
unexpected since carbon-oxygen bond breaking in an epoxide
generally occurs 106-107 times faster than in a structurally
related alcohol.5 It seems reasonable to compare these reactions
with the reactions of the structurally related compounds 3 and
4. The reactions of these substrates show an epoxide/hydrate
rate ratio as large as ∼107 (Chart 1).
Table 2. Calculated Enthalpies of the Reactions of 1-4 Using
B3LYP/6-311+G**//B3LYP/6-31G*
∆H,a
∆∆H,a
∆H,b
∆∆H,b
substrate
kcal mol-1
kcal mol-1
kcal mol-1
kcal mol-1
1
2
3
4
-48.0
-29.8
-17.6
+5.2
-18.2
-22.8
-49.8
-32.1
-19.7
+3.4
-17.7
-23.1
a Without solvation. b With solvation (see the Experimental Section).
Table 3. Calculated Enthalpies of the Carbocation Intermediates of
the Reactions of 1-4, Using B3LYP/6-311+G**//B3LYP/6-31G*
∆H,a
∆∆H,a
∆H,b
∆∆H,b
substrate kcal mol-1 kcal mol-1 kcal mol-1 kcal mol-1 ∆∆G‡
exp
1
2
3
4
-51.8
-48.4
-45.7
-34.0
-3.4
-1.2
0.8
4.2
-2.0
1.0
-9.6
-11.8
-12.0
16.2
a Without solvation. b With solvation (see the Experimental Section).
2), the calculated dehydration enthalpy (-1.2 kcal mol-1) was
More O’Ferrall and co-workers have thoroughly examined
the rate ratios of other relevant epoxide/hydrate derivatives.5
They found that other arenes show behavior similar to that of
benzene. Benzoannelation greatly slows down dehydration, but
the effect on epoxide reactivity is smaller. The result is that the
oxide/hydrate rate ratio increases with more unfavorable ben-
zoannelation and decreasing stability of the carbocationic
transition state. Thus, while koxide/khydrate of the benzene deriva-
tives is 0.08, naphthalene exhibits a ratio of 31 and 56
(depending on structure), and anthracene a ratio of 1430. The
nonaromatic isobutene derivatives show a ratio of 4.3 × 106.
Thus, the large epoxide/hydrate rate ratio of nonaromatic
compounds is not related to the cyclic nature of the reactants.
As discussed in detail previously,5 there is no reason that the
unusually small rate ratios of the arene derivatives are caused
by an unusually high reactivity of arene hydrates. Therefore,
the question arises: why do arene oxides show such a low
reactivity?
used to correct the measured value. The error in the calculated
dehydration enthalpy is estimated to be less than 1 kcal mol-1
.
A rate constant of carbocation formation from 2-cyclohexenol
(4) estimated by More O’Ferrall is given in Chart 1.5b The
derived rate constant at 25 °C is smaller (∼0.6 × 104 M-1 s-1
,
based upon the measured rate constants for reaction in 5.64 M
sulfuric acid at 20 and 30 °C),12 presumably because the
carbocation formation is not rate-limiting. A reaction enthalpy
of ∆H ∼ 0 has been reported.12
Calculations. The thermodynamics of the reactions depicted
in Schemes 2 and 3 was calculated using the Gaussian 98
program and the B3LYP hybrid functional along with a
continuum solvation method (see Computational Details for
more information). The results are given in Tables 2 and 3.
The pKa values of the substrates were derived by a calibrated
method from local ionization energies (see Supporting Informa-
tion). The pKa values were found to be -2.77, -2.13, -2.42,
and -1.97, respectively.
Homoaromatic Stabilization of Arene Oxides? A hypoth-
esis presented by More O’Ferrall and co-workers is that 1 and
other arene oxides show unusual stability attributable to homo-
aromaticity.5 The homoaromatic stabilization effect is expected
to be largest for the smallest molecule 1. Epoxides of alkenes
do not have this possibility of stabilization.
Discussion
Carbocation-like Transition States. The mechanisms of
acid-catalyzed carbon-oxygen bond breaking of arene hydrates
have been the subject of some recent studies.5,12-15 It was
concluded that the high reactivity is mainly related to the
Homoaromaticity is generally considered to be unimportant
for neutral molecules, but there is a consensus about its impor-
(11) Ross, A. M.; Pohl, T. M.; Piazza, K.; Thomas, M., Fox, B.; Whalen,
D. L. J. Am. Chem. Soc. 1982, 104, 1658.
(12) Jensen, J. L.; Uaprasert, V. J. Org. Chem. 1976, 41, 649.
(13) Boyd, D. R.; McMordie, R. A. S.; Sharma, N. D.; More O’Ferrall,
R. A.; Kelly, S. C. J. Am. Chem. Soc. 1990, 112, 7822.
(16) Kasperek, G. J.; Bruice, T. C.; Yagi, H.; Jerina, D. M. J. Chem.
Soc., Chem. Commun. 1972, 784.
(17) Pritchard, J. G.; Siddiqui, I. A. J. Chem. Soc., Perkin Trans. 2 1973,
452.
(14) Pirinccioglu, N.; Thibblin, A. J. Am. Chem. Soc. 1998, 120, 6512.
(15) Jia, Z. S.; Thibblin, A. J. Chem. Soc., Perkin Trans. 2, 2001, 247.
(18) Lamaty, G.; Maleq, R.; Selve, C.; Sivade, A.; Wylde, J. J. Chem.
Soc., Perkin Trans. 2 1975, 1119.

10150 J. Am. Chem. Soc., Vol. 123, No. 42, 2001
Jia et al.
Chart 2
-24.9 kcal mol-1 from the measured enthalpy of -23.7 kcal
mol-1 to give diol products (Table 1) combined with the
calculated enthalpy for dehydration (kdeh in Scheme 2) of the
trans 1,2-diol of -1.2 kcal mol-1 (see Results).
Approximative energies of aromatization may be estimated
as the difference in reaction energies of the two reaction pairs,
i.e., using the reactions of 3 and 4 as references (Scheme 2).23
The measured/derived aromatization energies calculated in this
way are for benzene oxide, ∆∆H ) -57.0 - (-24.9) ) -32.1
kcal mol-1, and for benzene hydrate, ∆∆H ) -38.7 - 0 )
-38.7 kcal mol-1. Thus, benzene oxide seems to be particularly
stable, and the homoaromaticity could be estimated as ∆∆∆H
) -32.1 - (-38.7) ) 6.6 kcal mol-1. In the same way, using
the calculated energies (Table 2) yields in the gas phase for
benzene oxide, ∆∆H ) -48.0 - (-17.6) ) -30.4 kcal mol-1
,
and for benzene hydrate, ∆∆H ) -29.8 - (+5.2) ) -35.0
kcal mol-1, which corresponds to a homoaromatization energy
of ∆∆∆H ) -30.4 - (-35.0) ) 4.6 kcal mol-1 (5.4 in water).
However, this stabilization energy does not fully account for
the low reactivity of benzene oxide.
The differences in free enthalpy of formation in aqueous
solution have been calculated as ∆∆Hf0 ) -33.9 kcal mol-1
for 1 and 2 and -37.7 kcal mol-1 for 3 and 4. This also suggests
that 1 is unusually stable (by 3.8 kcal mol-1).
Carbocation Intermediates. When discussing factors af-
fecting the rate of a reaction, it is relevant to consider the rela-
tive stability of the intermediates. Thus, we have calculated, in
a way similar to that for the total reactions, the energies of
carbocation formation employing B3LYP in combination with
AM1/SM2 to account for solvation effects. As shown in Table
3, there is a good correlation between the experimentally de-
rived ∆∆G‡ and the calculated ∆∆H for the formation of
cations from 1 and 2 and for the formation of cations from 3
and 4, respectively. If we use the reactions of 3 and 4 as
references as we did for the calculation of homoaromatization
(vide supra), we obtain for 1 and 3: ∆∆H ) -1.2 - 4.2 )
-5.4 kcal mol-1 and for 2 and 4 ∆∆H ) 0.8 - 16.2 ) -15.4
kcal mol-1, i.e., a difference between the two pairs of 10.0 kcal
mol-1. This is in good agreement with the experimentally
derived difference in activation energy of ∼10.6 kcal mol-1
(Table 3).
tance in stabilizing carbocations. However, Kollman and co-
workers have suggested that 1 is stabilized by interaction of a
lone pair on the oxygen atom with the π-orbitals of the benzene
oxide ring.19 Thus, 1 might be considered a 6π-analogue of the
2π-homoaromatic 7-norbornenyl carbocation (5) (Chart 2). Also
the neutral norcaradiene (6) has been suggested to be homoaro-
matic with a stabilizing resonance energy of 5-8 kcal mol-1 20
.
This molecule does not have any lone pair electrons, and the
proposed stabilization should come from σ-π interactions.
A homoaromatic stabilization effect of at least 10 kcal mol-1
for benzene oxide is suggested by the about 107-fold lower than
expected reactivity. However, the microcalorimetric measure-
ments gave reaction heats of the reactions for 1 and 2 (Table 1)
of ∆H ) -57.0 and -38.7 kcal mol-1, respectively. Thus, there
is a large driving force in both reactions but a very much greater
one for the epoxide ring opening reaction (∆∆H ) -18.3 kcal
mol-1).
The enthalpy of formation, ∆Hf0(aq), for 1 and 2 can be
derived from the reaction heats and the enthalpies of formation
of the products (Scheme 1). The results are ∆Hf0(1,aq) ) 20.7
kcal mol-1 and ∆Hf0(2,aq) ) -17.4 kcal mol-1, respectively.21
Thus, the enthalpy difference ∆∆Hf0(aq) of the two reactants 1
and 2 is 38.1 kcal mol-1. This is a large difference, but, without
having appropriate reference compounds, it is not possible to
conclude whether benzene oxide is substantially stabilized or
not.
In addition to the activation barrier of the C-O bond breaking
step, the energy of the rate-limiting step is directly dependent
on the pKa of the hydronated substrate (Schemes 2 and 3). Thus,
the rate difference could in principle be due to a difference in
basicity. A careful analysis using a calibrated model based on
average local ionization energies suggests a difference of less
than 0.4 pKa-units for substrates 1 and 3, corresponding to 0.5
kcal mol-1, with benzene oxide being the more acidic species,
pKa ) -2.77 (the acidities of all four substrates are within 1
pKa unit; see Computational Details). Thus, an unfavorable
preequilibrium (K, Schemes 2 and 3) could only explain a minor
part of the low reactivity of benzene oxide. The fact that benzene
oxide is more difficult to hydronate than the other substrates is
in accord with its unusual stability (homoaromaticity), which
is partially lost in the hydronated species.
Calculation of thermodynamic reaction energies using the
B3LYP hybrid functional and the continuum solvation method
AM1/SM2 yielded energies somewhat smaller than those
measured,22 ∆H ) -49.8 and -32.1 kcal mol-1 for 1 and 2,
respectively (Table 2), i.e., a difference of ∆∆H ) -17.7 kcal
mol-1, which is very close to the measured value of -18.3 kcal
mol-1 (Table 1). The calculated reaction energy for the related
compound 3 is much smaller -19.7 kcal mol-1, and the reaction
of 4 is endothermic by 3.4 kcal mol-1 (Scheme 2). An
experimental value of ∆H ∼ 0 for the reaction in 5.64 M sulfuric
acid has been reported.12 The enthalpy of the reaction of 3 to
the alcohol product shown in Scheme 2 could be derived as
(19) Hayes, D. M.; Nelson, S. D.; Garland, W. A.; Kollman, P. A. J.
Am. Chem. Soc. 1980, 102, 1255.
(20) Roth, W. R.; Kla¨rner, F.-G.; Siepert, G.; Lennartz, H.-W. Chem.
Ber. 1992, 125, 217.
In conclusion, the results suggests that roughly half of the
unusually low reactivity of 1 is due to a transition-state effect
(∼4 kcal mol-1) which also seems to be reflected in a high
energy of the hydroxy-substituted carbocation. As discussed
above, an unusual stability (homoaromatization) of benzene
oxide accounts for the other half.
(21) Based upon ∆Hf0(phenol,aq) ) -36.34 kcal mol-1, ∆Hf0(benzene,aq)
) 12.24 kcal mol-1, and ∆Hf0(water,aq) ) -68.32 kcal mol-1. These
values include the solvation energies for phenol and benzene given in the
following references: Nichols, N. F.; Wadso¨, I. J. Chem. Thermodyn. 1975,
7, 329. Gill, S. J.; Nichols, N. F.; Wadso¨, I. J. Chem. Thermodyn. 1975, 7,
175.
(22) Also other aromatization reactions show smaller than calculated
measured heats of reaction, presumably owing to an underestimation of
the solvation energies in the calculation methods.
(23) The reaction of 3 produces diols, mainly the trans 1,2- and trans
1,4-isomers.14

Is Benzene Oxide Homoaromatic?
J. Am. Chem. Soc., Vol. 123, No. 42, 2001 10151
were of HPLC quality and HPLC UV gradient quality, respectively.
All other chemicals used for the kinetic experiments were of reagent
grade and used without further purification.
Benzene oxide (1) was prepared by a three-step synthesis.8 First,
4,5-dibromocyclohexene was prepared from 1,4-cyclohexadiene. This
material was oxidized by m-chloroperoxybenzoic acid in chloroform,
followed by base-promoted elimination of the 4,5-dibromo-1,2-epoxi-
cyclohexane using potassium tert-butoxide in dry diethyl ether. The
NMR spectra of the product 2 were in accord with those previously
published.
An attempt to find direct evidence for an electronic/structural
difference in the transition states was carried out using gas-
phase calculations on transition states solvated by one or two
water molecules. The transition-state structure for benzene oxide
was easily found (Figure S5, Supporting Information). Unfor-
tunately, no complete set of structures could be achieved due
to barrierless C-O bond cleavage in several cases. Clearly, there
is need for more sophisticated solvation methods to tackle this
problem.
Benzene hydrate (2) was also prepared from 1,4-cyclohexadiene,
in a two-step process.9 Oxidation with m-chloroperoxybenzoic acid in
chloroform afforded 1,2-epoxi-4-cyclohexene. Subsequent cleavage of
the epoxide with methyllithium in dry diethyl ether gave a colorless
liquid. The purity was checked by NMR.
1,3-Cyclohexadiene oxide (3) was prepared from 1,3-cyclohexadiene
and m-chloroperoxybenzoic acid in a one-step process.10 Distillation
afforded pure product (NMR10 and HPLC).
Kinetics-UV Spectrophotometric Procedure. The reactions were
run in 3 mL standard quartz cells using the above-mentioned equipment.
Addition of a few microliters of a concentrated solution of the substrate
in acetonitrile to 2.5 mL of reaction solution gave an initial concentra-
tion of the substrate in the reaction flask of about 0.1 mM. The change
in absorbance was followed as a function of time and the pseudo-first-
order rate constant calculated by a nonlinear regression computer
program.
Kinetics-Microcalorimetric Procedure. This technique has the
advantage in that both the kinetic data and reaction heats are obtained
from the same kinetic experiment. The reactions were run in parallel
in two channels, both composed of a sample compartment and a
reference compartment. Glass vials (3 mL) were used as reaction and
reference vessels. All four vessels were filled at the same time with
2.5 mL of premixed buffer solution (glycerol-water or acetonitrile-
water). After this step, 20 µL of substrate in acetonitrile was added to
the two reaction vessels (final concentration ∼ 0.5 mM), while 20 µL
of pure acetonitrile was added to the reference vessels. The vials were
sealed with gastight PTFE septa and slowly introduced into the
compartments of the instrument for about 15 min of prethermostating.
They were then lowered further down into the detection chambers. After
a total time of 30-45 min it was possible to start recording the first-
order heat-flow decay. The reactions were followed for at least 10 half-
lives. The substrate concentrations in the reaction vials were obtained
indirectly by measuring the concentration of the product by HPLC,
i.e., phenol and benzene, respectively, after more than 10 half-lives.
However, the acetonitrile solution of 3 was prepared by weighing the
pure substrate.
Unusually High Reactivity of Nonaromatic Epoxides?
Concerted Ring-Opening? In principle, an alternative explana-
tion to the unusually low rate ratio of the benzene oxide reaction
compared with that of 1,3-cyclohexadiene might have been that
simple epoxides do not react through the carbocation but via
another faster reaction route. On the other hand, there is good
evidence that the acid-catalyzed ring opening of benzene oxide
goes through the carbocation; e.g., there is no build up of diol,
and there is no chloride ion effect reported for the reaction under
nonacidic conditions.
The reaction of 3 has been found to be catalyzed by added
chloride ion.11 Only nucleophilic attack of the chloride ion on
the neutral substrate was discussed. However, the hydronated
epoxide, having an even better leaving group, should be even
more prone to react by an SN2 mechanism. Therefore, the large
reactivity of 3 and other simple epoxides in acidic solution may
arise from SN2 and SN2′ reactions of the hydronated substrate
with water and added nucleophiles. The SN2 transition state of
such a reaction is positively charged and consequently should
show characteristics similar to the transition state of the proposed
stepwise mechanism.
The reaction of 3 did not show any bimolecular reaction in
acidic solution. The small increase in observed rate at pH 5.28
with added chloride ion was attributed to a salt effect.11 This is
consistent with a very reactive protonated epoxide substrate
showing a very small discrimination between different nucleo-
philes. Reactions of SN2′ type are expected to show syn
stereochemistry. Thus, the large amount (ca. 35%) of trans 1,4-
diol product, which accompanies the main product trans 1,2-
diol, indicates that the reaction is stepwise. Therefore, we
conclude that the reaction of 3 also occurs through the
carbocation.
The microcalorimeter was statically calibrated after a kinetic
experiment using the reaction solutions. The cooling constant of the
instrument was found to be 140 s; i.e., no correction for this parameter
was necessary for calculation of the rate constants of the reaction heat
decay.24 Very good first-order rate constants (kobs) were measured. These
agree well with those obtained by the UV-spectrophotometric technique
(Table 1).
Experimental Section
General Procedures. NMR spectra were recorded for CDCl3 solu-
tions with a Varian XL 300 spectrometer. Chemical shifts are indi-
rectly referenced to tetramethylsilane (TMS) via the solvent signal
(chloroform-d1 7.26 and 77.0 ppm). The high-performance liquid
chromatography analyses were carried out with a Hewlett-Packard 1090
liquid chromatograph equipped with a diode-array detector on an
Inertsil 5 OSD-2 (3 × 100 mm) reversed-phase column. The chroma-
tography was performed isocratically using acetonitrile in water as
the mobile phase. The UV spectrophotometry was performed on a
Kontron Uvicon 930 spectrophotometer equipped with an automatic
cell changer kept at constant temperature using a thermostated water
bath (HETO 01 PT 623). The reaction solutions for the kinetic
experiments were prepared by mixing acetonitrile or glycerol with
water at room temperature, ca. 22 °C. The pH was measured before
and after reaction using a Radiometer PHM82 pH meter equipped with
an Ingold micro glass electrode. The pH values given are those
measured before mixing with the organic solvent. The microcalorimetric
experiments were carried out with a dual channel calorimeter (Ther-
mometric Thermal Activity Monitor 2277). The signals were recorded
on both a two-channel potentiometric recorder (LKB 2210) and on a
computer.
The extrapolated heat flow at time zero (P0) was used for calculation
of the reaction heat (∆H) according to
P0 ) ∆Hkobs
n
(4)
where n is the amount of substrate (mol) in the reaction vial.
The estimated errors are considered as maximum errors derived from
maximum systematic errors and random errors.
Computational Details. All geometry optimizations reported in this
work were conducted using the Gaussian 98 program25 and the B3LYP
hybrid functional26 together with the 6-31G* basis set. Energies were
subsequently determined using B3LYP in combination with the larger
basis set 6-311+G**. Zero point energies and thermal corrections to
enthalpy were determined at the B3LYP/6-31G* level of theory. To
ensure that the B3LYP results were of sufficient quality, geometry
optimizations of benzene oxide and phenol were performed at MP2/
Materials. Diethyl ether and tetrahydrofuran were distilled under
nitrogen from sodium and benzophenone. Methanol and acetonitrile
(24) Thibblin, A. Chem. Scr. 1983, 22, 70.

10152 J. Am. Chem. Soc., Vol. 123, No. 42, 2001
Jia et al.
6-31G* with energies at MP2/6-311+G**. This showed no significant
changes in structure or thermodynamics.
The pKa values of the substrates 1-4 were derived from average
local ionization energies calculated using the program Hardsurf 9530
with the HF/6-31G* wave function and the B3LYP/6-31G* geometries.
The method was calibrated using the experimentally measured pKa
values of methanol, ethanol, dimethyl ether, and diethyl ether (see
Supporting Information for details).31
Free energies of solvation were estimated at the gas-phase geom-
etries using the continuum solvation method AM1/SM227 in Spar-
tan.28 The free energy has been used directly as a substitute for the
enthalpy of solvation in the estimation of solvation effects. Most of
the AM1/SM2 calculations were repeated using the CPCM29 solva-
tion model with parameters for water in combination with the HF/
6-31+G* wave function. However, these derived energies did not
show as good a correlation to the experimental data as the AM1/SM2
results.
Acknowledgment. We thank Prof. Rory More O’Ferrall for
comments and Prof. Tore Brink for providing us with the
Hardsurf 95 program. Financial support from the Swedish
Natural Science Research Council is gratefully acknowledged.
(25) Frisch, M. J.; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.; Robb,
M. A.; Cheeseman, J. R.; Zakrzewski, V. G.; Montgomery, J. A.; Stratmann,
R. E., Jr.; Burant, J. C.; Dapprich, S.; Millam, J. M.; Daniels, A. D.; Kudin,
K. N.; Strain, M. C.; Farkas, O.; Tomasi, J.; Barone, V.; Cossi, M.; Cammi,
R.; Mennucci, B.; Pomelli, C.; Adamo, C.; Clifford, S.; Ochterski, J.;
Petersson, G. A.; Ayala, P. Y.; Cui, Q.; Morokuma, K.; Malick, D. K.;
Rabuck, A. D.; Raghavachari, K.; Foresman, J. B.; Cioslowski, J.; Ortiz, J.
V.; Stefanov, B. B.; Liu, G.; Liashenko, A.; Piskorz, P.; Komaromi, I.;
Gomperts, R.; Martin, R. L.; Fox, D. J.; Keith, T.; Al-Laham, M. A.; Peng,
C. Y.; Nanayakkara, A.; Gonzalez, C.; Challacombe, M.; Gill, P. M. W.;
Johnson, B.; Chen, W.; Wong, M. W.; Andres, J. L.; Gonzalez, C.; Head-
Gordon, M.; Replogle, E. S.; Pople, J. A. Gaussian 98, Revision A.3;
Gaussian, Inc.: Pittsburgh, PA, 1998.
Supporting Information Available: Tables showing cal-
culations of pKa values using the Hardsurf 95 method and
graphical presentation of some calculated structures (7 pages,
print/PDF). See any current masthead page for ordering infor-
mation and Web access instructions.
JA010226H
(29) Spartan, version 5.0.3; Wavefunction, Inc.: Irvine, CA, 1997.
(30) Brinck, T.; Murray, J. S.; Politzer, P. J. Org. Chem. 1991, 56, 5012.
Brinck, T.; Murray, J. S.; Politzer, P.; Carter, R. E. J. Org. Chem. 1991,
56, 2934.
(31) Scorrano, G.; More O’Ferrall, R. ReV. Chem. Intermed. 1987, 7,
313.
(26) (a) Becke, A. D. J. Chem. Phys. 1993, 98, 5648. (b) Lee, C.; Yang,
W.; Parr, R. G. Phys. ReV. B 1988, 37, 785.
(27) Barone, V.; Cossi, M.; Tomasi, J. J. Comput. Chem. 1998, 19, 404.
(28) Cramer, C. J.; Truhlar, D. G. Science 1992, 256, 213.