Kinetic and thermodynamic stability of naphthalene oxide and related compounds. A comparative microcalorimetric and computational (DFT) study

Article abstract of DOI:10.1021/jo025953p

The kinetics of the acid-catalyzed ring opening of naphthalene 1,2-oxide (5) in highly aqueous media to give naphthols has been measured by heat-flow microcalorimetry. The reaction enthalpy of this aromatization reaction was measured as ΔH = -51.3 ± 1.7 kcal mol^-1. The unexpectedly low reactivity of naphthalene oxide is suggested to be due to an unusually large thermodynamic stability. A crude estimate of the stabilization effect, ～1 kcal mol^-1(not a significant stabilization), is obtained by using the measured reaction enthalpies of structurally related substrates as references. A larger value (2.7 kcal mol^-1) was obtained by calculation using the B3LYP hybrid functional corrected with solvation energies derived from semiempirical AM1/SM2 calculations. The origin of this effect is discussed in terms of homoconjugative stabilization and homoaromaticity. There is a good linear correlation (with slope = 0.63) between the experimentally measured free energy of activation and the calculated enthalpy of carbocation formation in water.

Full text of DOI:10.1021/jo025953p

Kin etic a n d Th er m od yn a m ic Sta bility of Na p h th a len e Oxid e a n d
Rela ted Com p ou n d s. A Com p a r a tive Micr oca lor im etr ic a n d
Com p u ta tion a l (DF T) Stu d y
Peter Brandt,† Zhi Sheng J ia,‡ and Alf Thibblin*,‡
Organic Chemistry, Department of Chemistry, Royal Institute of Technology, SE-100 44 Stockholm,
Sweden and Institute of Chemistry, University of Uppsala, Box 531, SE-751 21 Uppsala, Sweden
Alf.Thibblin@kemi.uu.se
Received May 22, 2002
The kinetics of the acid-catalyzed ring opening of naphthalene 1,2-oxide (5) in highly aqueous media
to give naphthols has been measured by heat-flow microcalorimetry. The reaction enthalpy of this
aromatization reaction was measured as ∆H ) -51.3 ( 1.7 kcal mol-1. The unexpectedly low
reactivity of naphthalene oxide is suggested to be due to an unusually large thermodynamic stability.
A crude estimate of the stabilization effect, ∼1 kcal mol-1(not a significant stabilization), is obtained
by using the measured reaction enthalpies of structurally related substrates as references. A larger
value (2.7 kcal mol-1) was obtained by calculation using the B3LYP hybrid functional corrected
with solvation energies derived from semiempirical AM1/SM2 calculations. The origin of this effect
is discussed in terms of homoconjugative stabilization and homoaromaticity. There is a good linear
correlation (with slope ) 0.63) between the experimentally measured free energy of activation and
the calculated enthalpy of carbocation formation in water.
In tr od u ction
Aromatic hydrocarbons undergo metabolism through
arene oxides. These oxidized products are known inter-
mediates in the metabolic reaction path leading to
carcinogenic aromatic diol epoxides that are know to react
with nucleic acids.1 These reactive species are therefore
of great biochemical interest.2-5 Surprisingly, the reactiv-
ity of benzene oxide and other arene oxides toward acid-
catalyzed ring opening is relatively low compared to the
reactivities of structurally related compounds which do
not undergo aromatization (Figure 1).6,7 Thus, the acid-
catalyzed ring opening of the cyclohexadiene oxide (3) is
7 orders of magnitude faster than the acid-catalyzed
dehydration of cyclohexenol (4), in sharp contrast to
† Royal Institute of Technology.
‡ University of Uppsala.
(1) Harvey, R. G.; Geacintov, N. E. Acc. Chem. Res. 1988, 21, 66.
(2) Lehr, R. E.; Wood, A. W.; Levin, W.; Conney, A. H.; J erina, 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.
(3) (a) Boyd, D. R.; McMordie, R. A. S.; Sharma, N. D.; Dalton, H.;
Williams, P.; J enkins, 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.
F IGURE 1. Second-order rate constants for acid-catalyzed
carbocation formation in water (compounds 1-4 and 8) and
in 25 vol % glycerol in water (compounds 5-7) at 25 °C. The
rate constant for formation of 2-naphthol (via the R-cation) is
given within brackets.
(4) J erina, D. M.; Daly, J . W.; Wikop, B.; Zaltzman-Nirenberg, P.;
Udenfriend, S. Biochemistry 1970, 9, 147.
(5) Boyd, D. R.; J erina, D. M. In Small Ring Heterocycles - Part III
“Oxiranes, Arene Oxides, Oxaziridines, Thietanes, Thietes, Thiazeles”;
A Hassner, A., Ed.; J ohn Wiley: NY, 1985; p 197.
(6) (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.
benzene oxide (1), which is about 1 order of magnitude
less reactive than benzene hydrate (2). Similarly, the
reaction of naphthalene oxide (5) to give 2-naphthol
through the R-carbocation is only about 14 times faster
than the naphthalene hydrate 6a , but carbocation forma-
tion from 1,2-dihydronaphthalene oxide (7) is 5 × 105
(7) J ia, Z. S.; Brandt P.; Thibblin, A. J . Am. Chem. Soc. 2001, 123,
10147.
10.1021/jo025953p CCC: $22.00 © 2002 American Chemical Society
Published on Web 10/03/2002
7676
J . Org. Chem. 2002, 67, 7676-7682

Kinetic and Thermodynamic Stability of Naphthalene Oxide
TABLE 1. Ra te Con sta n ts a n d Hea ts of Rea ction for th e Rea ction s of 1-8 a t 25 °C (Sch em e 1)
solventa
glycerol-waterb
glycerol-waterc
MeCN-waterd
watere
106kobs, s-1
kH, M-1s-1
Benzene Oxide (1) f Phenol7
∼22
Benzene Hydrate (2) f Benzene7
58
P0, µW
∆H, kcal mol-1
-57.0
-38.7
∆∆H,q kcal mol-1
-18.3
1,3-Cyclohexadiene oxide (3) f (1-hydroxy-1,3-cyclohexadiene)7
-23.7
-24.9j
2-Cyclohexenol (4) f 1,3-Cyclohexadiene7
∼10-3
Naphthalene 1,2-Oxide (5) F Naphthols
70.4 (3.5)o
∼0
glycerol-waterf
MeCN-waterg
11080
129 ( 4
33
-51.3 ( 1.7 (-52.2)h
-29.7 (-30.6)h
1-Hydroxy-1,2-dihydronaphthalene (6a ) f Naphthalene
glycerol-wateri
glycerol-waterj
MeCN-waterk
3043
235
134 ( 4
0.25
0.25
0.030
24
33
-21.6 ( 0.6
-21.6 ( 0.6
2-Hydroxy-1,2-dihydronaphthalene (6b)16 f Naphthalene
1.22 -18.4
1-Hydroxy-1,4-dihydronaphthalene (6c)16 f Naphthalene
254 -23.7
glycerol-waterb
glycerol-waterb
1,2-Dihydronaphthalene Oxide (7) f 11
300
glycerol-waterl
glycerol-waterm
6414
283
27
-25.6 ( 1.2
-31.6p
1-Hydroxy-1,2,3,4-tetrahydronaphthalene (8) f 14
glycerol-watern
6.4
6.4 × 10-6
∼ 0
0 ( 2.0
a
b
The buffer concentrations and pH values: before mixing with the organic cosolvent. 25 vol % glycerol in water. c 50 vol % glycerol
in water. 50 vol % acetonitrile in water. e In water. f 25 vol % glycerol in water, acetate buffer; k0 ) 1.3 × 10-3 s-1 (eq 1). 50 vol %
d
g
h
acetonitrile in water, 0.10 M acetate buffer pH 3.85. Enthalpy for the reaction to give 2-naphthol estimated by using the calculated
energy difference (0.9 kcal-1 mol-1) between 1- and 2-naphthol. 25 vol % glycerol in water, perchlorate buffer pH 1.92; kinetics were
i
j
k
studied by UV spectrophotometry. 25 vol % glycerol in water, perchloric acid buffer pH 3.03. 50 vol % acetonitrile in water in water,
perchloric acid buffer pH 2.35. 25 vol % glycerol in water, 0.10 M acetate buffer, k0 ) 154 × 10-6 s-1 (eq 1). 25 vol % glycerol in water,
l
m
n
0.10 M phosphate buffer pH 6.23. 50 vol % glycerol in water, 2.00 M perchloric acid. Rate constant measured with UV spectrophotometry.
p
o Rate constant for formation of 2-naphthol. Including the enthalpy of dehydration to give 13, calculated as -6.0 kcal mol-1 (Scheme 1).
q
Relative to the corresponding hydrate.
times faster than from 1,2,3,4-tetrahydro-1-naphthol (8).
All these reactions have been concluded to be subject to
specific acid catalysis and to proceed through carbocation
intermediates.6-8
Resu lts
Syn th esis. Naphthalene 1,2-oxide (5) was prepared
from 1,2-dihydronaphthalene in a three-step procedure
according to the method by Yagi and J erina.9 The same
precursor was used for preparation of 1,2-dihydronaph-
thalene oxide (7).10 The starting material for the four-
step synthesis of 1-hydroxy-1,2-dihydronaphthalene (6a )
was R-tetralone.11
Kin etic a n d Ca lor im etr ic Mea su r em en ts. All reac-
tions, except the reaction of 1,2,3,4-tetrahydro-1-naphthol
(8), produce large amounts of heat (Table 1) and are
easily followed by heat-flow microcalorimetry. The details
of the calorimetric measurements are given in the
experimental part.
The reaction of naphthalene 1,2-oxide (5) is fast in 25
vol % glycerol in water at 25 °C, and the kinetics were
therefore only studied by UV spectrophotometry (Table
1). The reaction has been studied previously in water and
in 10 vol % dioxane in water, respectively.8,12 Slower
reaction conditions were required for studies with our
The low reactivity of arene oxides has been proposed
to be due to an exceptionally large thermodynamic
stability of these compounds.6 This stability, which has
been attributed to homoaromaticity, is of considerable
interest in connection with studies of liver enzyme
metabolism of aromatic compounds and their possibility
to provide dihydrodiols and glutathione adducts.4 Recent
work supports the hypothesis that benzene oxide has an
additional ground-state stabilization. Thus, we have
estimated the extra stabilization energy for benzene oxide
to be ca. 7 kcal mol-1 in a combined heat-flow microcalo-
rimetric and computational study.7
We have now performed a similar study for naphtha-
lene oxide (5) to determine the effect of benzoannelation
on the reactivity and reaction heats. The measured low
reactivity of this compound suggests an additional ground-
state stabilization of about 6 kcal mol-1, which in the
present work is compared with estimated stabilization
energies based upon measured reaction heats and cal-
culations.
(9) Yagi, H.; J erina, D. M. J . Am. Chem. Soc. 1975, 97, 3185.
(10) Pedragosa-Moreau, S.; Archelas, A.; Furstoss, R. Tetrahedron
1996, 52, 4593.
(11) Boyd, D. R.; Sharma, N. D.; Kerley, N. A.; McMordie, R. A. S.;
Sheldrake, G. N.; Williams P.; Dalton, H. J . Chem. Soc., Perkin Trans.
1 1996, 67.
(8) Kasperek, G. J .; Bruice, T. C. J . Am. Chem. Soc. 1972, 94,
198.
(12) Bushman, D. R.; Sayer, J . M.; Boyd, D. R.; J erina, D. M. J .
Am. Chem. Soc. 1989, 111, 2688.
J . Org. Chem, Vol. 67, No. 22, 2002 7677

Brandt et al.
SCHEME 1
compounds 5-7 in aqueous glycerol (this work, Table 1),
are shown in Figure 1.
The rate constant for 8 refers to the carbocation
formation measured in aqueous trifluoroethanol extrapo-
lated to water.16 A very similar value is expected for 25%
glycerol in water. Carbocation formation is not rate
limiting so the rate of product formation is smaller (Table
1). The equilibrium constant has been measured in
aqueous solution at 25 °C as [14]eq/[8]eq ) 22.17
Ca lcu la tion s. The thermodynamics of the reactions
depicted in Scheme 1, including the intermediate car-
bocations, was calculated using the Gaussian 98 pro-
gram18 and the B3LYP hybrid functional19 along with the
continuum solvation method AM1/SM220 in Spartan21
(see the Experimental Section for more details). The
results are given in Tables 2 and 3.
Discu ssion
The acid-catalyzed aromatization of naphthalene oxide
(5) has been the subject of several mechanistic stud-
ies.8,12,22 There seems to be a consensus that the ring-
opening is subject to specific acid catalysis and proceeds
through carbocation intermediates (Scheme 2). The reac-
tion, which is analogous to that of benzene oxide, involves
a rate-limiting cleavage of one of the carbon-oxygen
bonds to give a carbocation that rapidly loses a hydron
to give naphthol. This final aromatization step may be
preceded by a hydride shift (NIH shift). There is direct
evidence of the cation character of the transition state
of acid-catalyzed epoxide ring-opening reactions.8,22-24
There is also a noncatalyzed reaction path (eq 1), but that
is not the subject of the present study.
heat-flow microcalorimeter equipment. Therefore, 50 vol
% acetonitrile in water was used as solvent (Table 1).
The reaction yields 95% of 1-naphthol and 5% of 2-naph-
thol (Scheme 1). The kinetics of the reactions involves
an uncatalyzed path and a specific acid-catalyzed path
according to:
kobs ) k0 + kH[H+]
(1)
The carbon-oxygen bond cleavage of unsymmetrical
arene oxides, such as naphthalene oxide, yields generally
two isomeric cations. The product composition is expected
to reflect the relative stabilities of the R- and â-carboca-
tions. The ratio decreases by benzoannelation adjacent
(R) to the cation center from a kR/kâ ratio of 1.0 for the
symmetrical benzene oxide to kR/kâ ) 0.05 for naphtha-
lene oxide (Table 1), and 1,2-anthracene oxide shows an
even lower ratio.6a This effect has been rationalized by
less resonance stabilization owing to the benzoannelation
The measured second-order rate constant, based upon
the observed rate constants in 25 vol % glycerol in water
given in Table 1, is kH ) 70.4 M-1s-1 at 25 °C. The
uncatalyzed path has a rate constant of k0 ) 1.3 × 10-3
s-1
.
The cleavage reaction of 1,2-dihydronaphthalene oxide
(7) in 25 vol % glycerol in water gives a mixture of diols
with the trans 1,2-diol (11) as the main product (Scheme
1). The rate constants in aqueous solution at 25 °C and
30 °C and product composition of diols (94% trans and
6% cis) have been reported previously.13,14 The kinetics
is in accord with eq 1. To obtain the reaction heat for
the production of the mono-hydroxy substituted com-
pounds (Scheme 1), the calculated dehydration enthalpies
(-21.7 - (-18.5) ) -3.2 and -24.5 - (-18.5) ) -6.0
kcal mol-1, respectively, Table 1) were used to correct the
measured value.
The dehydration of 1-hydroxy-1,2-dihydronaphthalene
(6a ) has been studied in aqueous acetonitrile and in
aqueous glycerol (Table 1). The kinetics of the dehydra-
tion has been studied previously in aqueous solution.4
Table 1 also shows previously reported reaction heats for
the two other naphthalene hydrate isomers 6b and 6c.15
The second-order rate constants for the acid-catalyzed
reactions of compounds 1-4 and 8 in water,6 and of
(16) 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.
(17) Dey, J .; O’Donoghue, A. C.; More O’Ferrall, R. A. J . Am. Chem.
Soc. 2002, 124, 8561.
(18) Frisch, M. J .; Trucks, G. W.; Schlegel, H. B.; Scuseria, G. E.;
Robb, M. A.; Cheeseman, J . R.; Zakrzewski, V. G.; Montgomery, J . A.,
J r.; Stratmann, R. E.; 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.; J ohnson, B. G.; Chen,
W.; Wong, M. W.; Andres, J . L.; Head-Gordon, M.; Replogle, E. S.;
Pople, J . A. Gaussian 98, Gaussian, Inc.: Pittsburgh, PA, 1998.
(19) (a) Becke, A. D. J . Chem. Phys. 1993, 98, 5648. (b) Lee, C.; Yang,
W.; Parr, R. G. Phys. Rev. B 1988, 37, 785.
(20) Cramer, C. J .; Truhlar, D. G. Science 1992, 256, 213. Cramer,
C. J .; Truhlar, D. G. In Rev. Comput. Chem. Lipkowitz, K. B., Boyd,
D. B., Eds.; VCH: New York, 1995; Vol. 6, pp 1-72.
(21) Spartan version 5.0.3, Wavefunction, Inc., Irvine, CA 1997.
(22) (a) Kasperek, G. J .; Bruice, T. C. Yagi, H.; J erina, D. M. J .
Chem. Soc., Chem. Commun. 1972, 784. (b) Bruice P. Y.; Bruice, T. C.
J . Am. Chem. Soc. 1976, 98, 2023. (c) Boyd, D. R.; Daly, J . W.; J erina,
D. M. Biochemistry 1972, 11, 1961.
(13) Becher, A. R.; J anusz, J . M.; Bruice, T. C. J . Am. Chem. Soc.
1979, 101, 5679.
(14) Gillilan, R. E.; Pohl, T. M.; Whalen, D. L. J . Am. Chem. Soc.
1982, 104, 4481.
(15) Pirinccioglu, N.; Thibblin, A. J . Am. Chem. Soc. 1998, 120, 6512.
7678 J . Org. Chem., Vol. 67, No. 22, 2002

Kinetic and Thermodynamic Stability of Naphthalene Oxide
TABLE 2. Ca lcu la ted En th a lp ies of th e Rea ction s of 1-8 Usin g B3LYP /6-311+G**//B3LYP /6-31G*
substrate
product
∆H,a kcal mol-1
∆∆H, kcal mol-1
∆H,b kcal mol-1
∆∆H,b kcal mol-1
-17.7
17
27
37
47
5
phenol
benzene
-48.0
-29.8
-17.6
+5.2
-37.9
-38.2
-18.6
-17.6
-21.4
-17.6
-17.4
-19.4
+2.0
-18.2
-49.8
-32.1
-19.7
+3.4
-38.9
-39.7
-21.6
-20.3
-24.5
-15.5
-18.7
-21.5
-0.7
1-hydroxy-1,3-cyclohexadiene
-22.8
-23.1
cyclohexadiene
1-naphthol
2-naphthol
naphthalene
naphthalene
naphthalene
11
-19.6c
-18.1c
6a
6b
6c
7
7
7
12
13
14
-21.4d
-20.9d
8
a
b
d
Without solvation. With solvation (see the Experimental Section). c Reference: reaction of 6a . Reference: reaction of 8.
TABLE 3. Ca lcu la ted En th a lp ies of th e Ca r boca tion
In ter m ed ia tes (r to th e Ben zen e Rin g) of th e Rea ction s
of 5-8, Usin g B3LYP /6-311+G**//B3LYP /6-31G*
SCHEME 3
∆H,a
∆∆H,a
∆H,b
∆∆H,b
c
substrate kcal mol-1 kcal mol-1 kcal mol-1 kcal mol-1 ∆∆Gq
exp
5
-54.6
4.1
-58.3 (â)
-48.3
0.8 (â)
6a
6b
6c
7
-6.3d
-8.0e
-5.5e
7.7
-3.5d
-5.7e
-2.4e
-1.6d
-50.3 (â)
-52.8 (â)
-54.4
6.5 (â)
3.2 (â)
3.9
-2.4e
0.8e
8
-42.4
-15.1
14.4
-13.4
-7.8
a
b
Without solvation. With solvation (see the Experimental
Section). c ∆∆Gq
) -RT ln(koxide/khydrate). Carbocation com-
d
exp
pared to the carbocation intermediate of 5 en route to 2-naphthol.
e Carbocation compared to the carbocation intermediate of 5 en
route to 1-naphthol.
SCHEME 2
SCHEME 4
and has been discussed thoroughly by More O’Ferrall and
co-workers.6a
The acid-catalyzed cleavage of the structurally related
epoxide 7, but which does not involve aromatization, is
expected to give only the benzylic R-carbocation. The
â-carbocation should be much less stable (Scheme 3).
There are close similarities between the transition
state of the acid-catalyzed reaction of naphthalene oxide
and the transition state of the acid-catalyzed aromati-
zation of the naphthalene hydrates 6a and 6b (Schemes
3 and 4). It has been concluded that the high reactivity
of these hydrates is mainly related to the stability of their
cation-forming transition states, and that the stability
of the reactant is of less importance.15,16 The carbocation
stability in the reaction of naphthalene oxide is disfa-
vored by the presence of the hydroxy substituent. How-
ever, this effect is generally small when comparing the
reactivity of arene oxides and hydrates since the major
factor should be the gain in energy brought about by the
release in strain from the cleavage of the epoxide ring.
Thus, the acid-catalyzed reaction of naphthalene oxide
J . Org. Chem, Vol. 67, No. 22, 2002 7679

Brandt et al.
SCHEME 5
accounted for by considering a homoconjugative effect
between the pair of epoxide C-O molecular orbitals and
the π- and π* orbitals of the butadiene moiety. In the
case of norcaradiene 9, this effect has been thoroughly
discussed by J orgensen.26
Our computational results support this latter explana-
tion. A second-order perturbation theory analysis of the
Fock matrix in a natural bond order basis for benzene
oxide indeed shows that the most important stabilizing
interaction is electron donation from the C-C π bonds
to the nearest C-O antibonding orbital and from the
C-O bonding orbitals to the π*-orbital system of the
diene. The symmetry requirement of the participating
molecular orbitals determining whether the system
becomes homoaromatic or homoantiaromatic has previ-
ously been discussed for cyclopropane derivatives.26
Do the calorimetric data of Table 1 support the as-
sumption of an increased stability of naphthalene oxide?
It is more complicated to compare the reactions of
naphthalene oxide with suitable references owing to the
lack of symmetry of the reactant. The reaction heat (∆H)
of -51.3 kcal mol-1 refers to the formation of a mixture
of 95% of 1-naphthol and 5% of 2-naphthol. The experi-
mental results do not give information about the indi-
vidual reaction enthalpies. However, the calculations
indicate that the two enthalpies are similar, 2-naphthol
is only 0.9 kcal mol-1 more stable than 1-naphthol in
water (Table 2).
is expected to be much faster than the corresponding
reactions of the naphthalene hydrates 6a and 6b. How-
ever, the reaction to give 2-naphthol, which occurs
through the R-cation, is only about 14 times faster than
the reaction of 6a . The reactions through the â-carboca-
tion show a similar behavior; the reaction of 5 producing
1-naphthol is only 55 times faster than the reaction of
6b. This is in sharp contrast with the corresponding
reactions of 7 and 8, which exhibit a rate-ratio for acid-
catalyzed carbocation formation of ∼5 × 105 (Table 1,
Figure 1). The rate-ratio of the latter reactions, which
are assumed to occur exclusively through the R-carboca-
tions (vide supra, Scheme 3), should be a reasonable
reference for prediction of the relative reactivity of 5 (to
give 2-naphthol) and 6a through the R-carbocations. This
estimated rate ratio of ∼5 × 105 is about 3.6 × 104 times
larger than the measured ratio. This corresponds to a
Let us assume that the reaction enthalpy for formation
of 2-naphthol is -51.3-0.9 ) -52.2 kcal mol-1. This could
be compared with the reaction enthalpy of -21.6 kcal
mol-1 for the aromatization of 6a by acid-catalyzed
dehydration. Thus, as expected, the ring-opening aroma-
tization of naphthalene oxide has a driving force that is
much larger than that of aromatization of the naphtha-
lene hydrate 6a , ∆∆H ) -52.2 - (-21.6) ) -30.6 kcal
difference in free activation energy of about 6 kcal mol-1
.
Since there is no reason to expect anomalous reactivity
of the naphthalene hydrates, we attribute the difference
in predicted and measured free activation energy to an
unusually low reactivity of naphthalene oxide. The reason
for the unusually slow reaction is discussed below.
Hom oa r om a tic Sta biliza tion of Na p h th a len e Ox-
id e? In analogy with what has been proposed for benzene
oxide, the slow reaction of naphthalene oxide may reflect
an increased stability of this compound caused by ho-
moaromaticity.6,7 Thus, we have recently reported that
the low reactivity of benzene oxide in acid-catalyzed
aromatization, corresponding to an increase in reaction
barrier of about 10 kcal mol-1, is accompanied by an
experimentally determined stabilization energy of 6.6
mol-1
.
The analogous reactions of 7 and 8, but which do not
involve aromatization, may be used for estimation of the
aromatization energies involved in the reactions of 5 and
6a . However, first we have to correct the measured
enthalpy of the reaction of 7, which experimentally yields
the diol, for the elimination of one water molecule to give
the enol 13 (Scheme 3 and Table 1). Calculations estimate
this enthalpy to be -6.0 kcal mol-1.27Accordingly, the
corrected enthalpy is ∆H ) -25.6 + (- 6.0) ) -31.6 kcal
mol-1. The aromatization energies for the naphthalene
oxide and naphthalene hydrate reactions are then esti-
mated as ∆∆H ) -52.2 - (-31.6) ) -20.6 kcal mol-1
and ∆∆H ) -21.6 - (0) ) -21.6 kcal mol-1, respectively.
According to this estimation, there is experimental
support for some homoaromatic stabilization of naph-
kcal mol-1 7
. This value was derived from the calori-
metrically measured reaction enthalpies. A similar value
(5.4 kcal mol-1) was obtained by calculations.7
Two decades ago Kollman and co-workers proposed
that an interaction in the benzene oxide (1) between the
π-orbital system of the diene with a lone pair of the
oxygen results in stabilization of the ground state (Scheme
5).25 Thus, as was pointed out by More O’Ferrall and co-
workers,6 benzene oxide is a 6-π analogue of the 2-π
homoaromatic 7-norbornenyl cation 10. However, they
concluded that the stabilization should be dominated by
σ-π interactions. These orbital interactions could be
thalene oxide, ∆∆∆H ) -20.6 - (-21.6) ) 1.0 kcal mol-1
.
However, because this estimate is based upon several
reaction heats, the maximum error in this value is larger
than 2 kcal mol-1 so there is no conclusive evidence for
a stabilization. The calculated gas phase enthalpies of
the same four reactions (Table 2) suggest a slightly
stronger stabilization: ∆∆∆H ) (-38.2 - (-19.4)) -
(-18.6 - (2.0)) ) 1.8 kcal mol-1. Correcting for solvation
in water gives ∆∆∆H ) (-39.7 - (-21.5)) - (-21.6 -
(-0.7)) ) 2.7 kcal mol-1. These values should be com-
(23) Pritchard, J . G.; Siddiqui, I. A. J . Chem. Soc., Perkin Trans. 2
1973, 452.
(24) Lamaty, G.; Maleq, R.; Selve, C.; Sivade A.; Wylde, J . J . Chem.
Soc., Perkin Trans. 2 1975, 1119.
(25) Hayes, D. M.; Nelson, S. D.; Garland, W. A.; Kollman, P. A. J .
Am. Chem. Soc. 1980, 102, 1255.
(26) J orgensen, W. L. J . Am. Chem. Soc. 1976, 98, 6784.
7680 J . Org. Chem., Vol. 67, No. 22, 2002

Kinetic and Thermodynamic Stability of Naphthalene Oxide
oxides. For benzene oxide, the calculations support a
homoconjugative effect between the pair of epoxide C-O
bonds and the π-system of the butadiene part of the
molecule. As previously suggested electron donation
between the epoxide oxygen lone-pair and the π-system
could be rejected.
The experiments indicate that benzoannelation de-
creases the effect of homoaromaticity. Thus, the stabili-
zation energy decreases from 7 kcal mol-1 for benzene
oxide to 1 kcal mol-1 for naphthalene oxide. In conflict
with experiments, calculations suggest that benzene
oxide and naphthalene oxide should be stabilized to the
same extent. The reason to this inconsistency is not
understood.
The rates of the reactions of the arene oxides and the
arene hydrates are well described by the calculated car-
bocation stabilities. The results indicate that the carbon-
oxygen bond is substantially broken and that there is a
large carbocation character in the transition state.
F IGURE 2. The experimentally measured free energy of
activation versus the calculated enthalpy of carbocation forma-
tion in water (in kcal mol-1, Table 3) for the reactions of the
naphthalene derivatives 5-8 through the R-cations (filled
circles, slope 0.63) and â-cations (open circles). The corre-
sponding data for the “benzene derivatives” 1-4 are included
in the plot (triangles; the rate constant for 1 is statistically
corrected).7
Exp er im en ta l Section
Gen er a l P r oced u r es. NMR spectra were recorded for
CDCl3 solutions with a 300 MHz spectrometer. Chemical shifts
are indirectly referenced to TMS via the solvent signal
(chloroform-d1 7.26 and 77.0 ppm). The high-performance
liquid chromatography analyses were carried out with a liquid
chromatograph equipped with a diode-array detector on an
Inertsil 5 OSD-2 (3 × 100 mm) reversed-phase column. The
chromatography was performed isocratically using acetonitrile
in water as the mobile phase. The UV spectrophotometry was
performed on a spectrophotometer equipped with an automatic
cell changer kept at constant temperature using a thermo-
stated water bath. 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 micro glass
electrode. The pH values given are those measured before
mixing with the organic solvent. The microcalorimetric experi-
ments were carried out with a dual channel calorimeter
(Thermometric Thermal Activity Monitor 2277). The signals
were recorded on both a two-channel potentiometric recorder
and on a computer.
Ma ter ia ls. Diethyl ether and tetrahydrofuran were distilled
under nitrogen from sodium and benzophenone. Methanol and
acetonitrile 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.
Na p h th a len e 1,2-oxid e (5) was prepared from 1,2-dihy-
dronaphthalene by reaction with N-bromosuccinimide (NBS)
in acetic acid followed by reaction with NBS in carbon
tetrachloride with UV radiation using R,R′-azoisobutyrodini-
trile as an initiator. The final step with sodium methoxide in
dry tetrahydrofuran afforded 5 as product.9 The NMR spec-
trum of the recrystallized product showed 3.0 mol % of
1-naphthol but no other impurities.
pared to the homoaromatic stabilization energy calcu-
lated for benzene oxide; 4.6 and 5.4 kcal mol-1 in the gas
phase and in water, respectively.
To conclude, the combined heat-flow microcalorimetry
and computational study suggests that the homoaromatic
stabilization is lower for naphthalene oxide than for
benzene oxide. This is in contrast to the purely theoretical
study, which suggests that the hyperconjugative effects
should be very similar.
Ca r boca tion Sta bilities a n d Kin etics. The rate-
limiting C-O bond breaking step is preceded by an acid-
base preequilibrium (Schemes 2-4). The basicities of the
benzene derivatives 1-4 have been found to be very
similar.7 There is only a small difference in pKa of the
protonated substrates (<0.4 kcal mol-1 according to
calculations). Very similar pKa values are also expected
for the compounds 5-8. Thus, we conclude that an
unfavorable preequilibrium could only explain a minor
part of the low reactivity of naphthalene oxide.
The rates of the cation-forming step of the studied
reactions are expected to be closely related to the stability
of the carbocation (vide supra). Indeed, the enthalpies
for the formation of the cations (Table 3), which are
calculated in a way similar to that for the total reactions,
show a good correlation with reactivity (Figure 2, the data
for the corresponding benzene derivatives are plotted in
the same diagram). The regression line for the reactions
of the compounds 5-8 through the R-cations has a slope
of 0.63. Thus, the carbocation stability is substantially,
but not fully, reflected in the reactivity.
1-Hyd r oxy-1,2-d ih yd r on a p h th a len e (6a ) was prepared
by a four-step procedure:11 reduction of R-tetralone with
sodium borohydride, acetylation with acetic acid anhydride,
bromination with NBS, and elimination with sodium meth-
oxide. Analysis by NMR and HPLC showed pure material
besides a small amount of naphthalene (3.0 mol %).
1,2-Dih yd r on a p h th a len e oxid e (7)10,13 was prepared by
the following procedure, which is a slight modification of the
previously published methods.10 1,2-Dihydronaphthalene was
reacted with NBS in acetone followed by reaction of the
isolated and purified bromohydrin (flash chromatography) with
The large ∆Gq for benzene oxide (Figure 2) is in accord
with the previous proposal that its unusually low reactiv-
ity is due to a combination of an unusually stable
compound (homoaromaticity) and an unusually high
energy of the carbocationic transition state.7
Con clu sion s. The heat-flow microcalorimetry mea-
surements and computational analysis presented in this
work have been aimed at increasing the understanding
of the homoaromaticity concept is connection with arene
J . Org. Chem, Vol. 67, No. 22, 2002 7681

Brandt et al.
sodium methoxide in tetrahydrofuran. Flash chromatography
gave pure material (HPLC and NMR).
tr a n s-1,2-Dih yd r oxy-1,2,3,4-t et r a h yd r on a p h t h a len e
(11)10,13,14,28 was prepared in accord with a previously published
method.10 The crude product was purified by recrystallization
from chloroform, mp 110-111 °C.
for the cooling constant parameter (κ) of the instrument was
necessary for calculation of the rate constants of the reaction
heat decay.29 Very good first-order rate constants (kobs) were
measured. These agree well with those measured in the kinetic
studies using the UV spectrophotometric technique (Table 1).
The extrapolated heat flow at time zero (P0) was used for
calculation of the reaction heat (∆H) according to eq 2:29
Kin etics. UV Sp ectr op h otom etr ic P r oced u r e. The reac-
tions 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 concentration 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.
Kin etics. Micr oca lor im etr ic P r oced u r e. This technique
has the advantage that both kinetic data and reaction heats
are obtained from the same kinetic experiment.29 The reactions
were run in parallel in the 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 about
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., 1- and 2-naphthol, naphthalene, 1,2-
dihydroxy-1,2,3,4-tetrahydronaphthalene, and 1,2-dihydronaph-
thalene, respectively, after more than 10 half-lives.
P0 ) ∆H kobs n0κ/ (κ - kobs
)
(2)
where n0 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.
Com p u ta tion a l Meth od s. All geometry optimizations
reported in this work were conducted using the Gaussian 98
program18 and the B3LYP hybrid functional19 together with
the 6-31G* basis set. Thermal corrections to enthalpies were
also determined at this level of theory. The final energies were
subsequently determined using B3LYP in combination with
the larger basis set 6-311+G**. The B3LYP results have
recently been compared with calculations at the MP2/6-
311+G**//MP2/6-31G* level of theory without any significant
change in either geometry or thermodynamics for a similar
set of molecules.7
Free energies of solvation were estimated at the gas-phase
geometries using the continuum solvation method AM1/SM220
in Spartan.21 The free energy of solvation has been used
directly as an approximation of the enthalpy of solvation in
the estimation of solvation effects.
This combination of calculations was also used in the
conformational searches of the hydroxy-containing compounds
to ensure that the global minimum was used in the calculation
of homoaromaticity.
Ack n ow led gm en t. We thank the Swedish Natural
Science Research Council for supporting this work.
The microcalorimeter was statically calibrated after
a
kinetic experiment using the reaction solutions. No correction
Su p p or t in g In for m a t ion Ava ila b le: Geometries at
B3LYP/6-31G* and Cartesian coordinates/energies in Hartrees
of optimized structures. This material is available free of
charge via the Internet at http://pubs.acs.org.
(27) The error in the enthalpy of dehydration (∆Hcalc, Scheme 1)
should be less than ( 2 kcal mol-1. This estimation is based upon our
calculation of the enthalpy of dehydration of ethanol in the gas phase
(-9.9 kcal mol-1, which is 1.0 kcal mol-1 lower than the literature
value) and the error given for solvation by Cramer and Truhlar (0.7
kcal mol-1, mean absolute error).
J O025953P
(28) Cook, J . W.; Loudon J . D.; Williamsson, W. F. J . Chem. Soc.
1950, 911.
(29) Thibblin, A. J . Phys. Org. Chem. 2002, 15, 233.
7682 J . Org. Chem., Vol. 67, No. 22, 2002