Ketonization equilibria of phenol in aqueous solution

Article abstract of DOI:10.1139/v99-048

The two keto tautomers of phenol (1), cyclohexa-2,4-dienone (2) and cyclohexa-2,5-dienone (3), were generated by flash photolysis of appropriate precursors in aqueous solution, and the pH-rate profiles of their enolization reactions, 2 → 1 and 3 → 1, were measured. The rates of the reverse reactions, 1 → 2 and 1 → 3, were determined from the rates of acid-catalyzed hydron exchange at the ortho- and para-positions of 1; the magnitude of the kinetic isotope effect was assessed by comparing the rates of hydrogenation of phenol-2t and -2d. The ratios of the enolization and ketonization rate constants provide the equilibrium constants of enolization, pK_E(2, aq, 25°C) = -12.73 ± 0.12 and pK_E(3, aq, 25°C) = -10.98 ± 0.15. Combination with the acidity constant of phenol also defines the acidity constants of 2 and 3 through a thermodynamic cycle. These ketones are remarkably strong carbon acids: pK_a(2) = -2.89 ± 0.12 and pK_a(3) = -1.14 ± 0.15. They disappear by proton transfer to the solvent with lifetimes, τ(2) = 260 μs and τ(3) = 13 ms, that are insensitive to pH in the range from 3-10.

Full text of DOI:10.1139/v99-048

605
Ketonization equilibria of phenol in aqueous
solution1
Marco Capponi, Ivo G. Gut, Bruno Hellrung, Gaby Persy, and Jakob Wirz
Abstract: The two keto tautomers of phenol (1), cyclohexa-2,4-dienone (2) and cyclohexa-2,5-dienone (3), were
generated by flash photolysis of appropriate precursors in aqueous solution, and the pH–rate profiles of their
enolization reactions, 2 → 1 and 3 → 1, were measured. The rates of the reverse reactions, 1 → 2 and 1 → 3, were
determined from the rates of acid-catalyzed hydron exchange at the ortho- and para-positions of 1; the magnitude of
the kinetic isotope effect was assessed by comparing the rates of hydrogenation of phenol-2t and -2d. The ratios of the
enolization and ketonization rate constants provide the equilibrium constants of enolization, pKE(2, aq, 25°C) = –12.73
± 0.12 and pKE(3, aq, 25°C) = –10.98 ± 0.15. Combination with the acidity constant of phenol also defines the acidity
constants of 2 and 3 through a thermodynamic cycle. These ketones are remarkably strong carbon acids: pKa(2) =
–2.89 ± 0.12 and pKa(3) = –1.14 ± 0.15. They disappear by proton transfer to the solvent with lifetimes, τ(2) = 260 µs
and τ(3) = 13 ms, that are insensitive to pH in the range from 3–10.
Key words: proton transfer, tautomers, flash photolysis, kinetic isotope effect, pH–rate profiles.
Résumé : Opérant en solution aqueuse, on a généré les deux tautomères cétoniques du phénol (1), la cyclohexa-2,4-
diénone (2) et la cyclohexa-2,5-diénone (3), par photolyse éclair de précurseurs appropriés et on a mesuré les profils
pH de la vitesse pour leurs réactions d’énolisation, 2 → 1 et 3 → 1. On a déterminé les vitesses des réactions inverses
1 → 2 et 1 → 3 à partir des vitesses d’échange d’hydron acidocatalysées dans les positions ortho- et para- de 1; on a
évalué l’amplitude de l’effet isotopique cinétique en comparant les vitesses d’hydrogénation du phénol-2t et -2d. Les
rapports des constantes de vitesse d’énolisation et de cétonisation permettent d’arriver aux constantes d’équilibre
d’énolisation, pKE(2, aq, 25°C) = –12,73 ± 0,12 et pKE(3, aq, 25°C) = –10,98 ± 0,15. Une combinaison avec la
constante d’acidité du phénol permet de définir aussi les constantes d’acidité de 2 et de 3 à travers un cycle
thermodynamique. Ces cétones sont des acides carbonés relativement forts: pKa(2) = –2,89 ± 0,12 et pKa(3) = –1,14 ±
0,15. Elles disparaissent par transfert de proton vers le solvant avec des temps de vie, τ(2) = 260 µs et τ(3) = 13 µs
qui sont insensibles au pH entre 3 et 10.
Mots clés : transfert de proton, tautomères, photolyse éclair, effet isotopique cinétique, profils pH–vitesse.
[Traduit par la Rédaction] Capponi et al. 613
Introduction
beatable research team of Jerry Kresge and his wife, Yvonne
Chiang. On the scientific side, we were rewarded by partici-
pating in the kinetic and thermodynamic characterization of
the enolization equilibria of a representative selection of car-
bonyl compounds over the last two decades (2).
Twenty years ago, we reported on the observation of ace-
tophenone enol by flash photolysis and determined the acid-
ity constant of the enol from the pH–rate profile of its
ketonization in aqueous solution (1). That work was a side-
line of studies aimed primarily at the characterization of
diradical intermediates in Norrish Type 2 reactions, and would
not have been pursued, had Prof. A.J. Kresge not visited
Basel and convinced us that the photochemical approach to
study enol kinetics should be further exploited. Thus began a
growing friendship and continued collaboration with the un-
The protomeric equilibria of thermodynamically stable
enols have received less attention. Phenol (1) is much more
stable than its keto tautomers cyclohexa-2,4-dienone (2) and
cyclohexa-2,5-dienone (3), because proton transfer from ox-
ygen to carbon disrupts π-conjugation of the phenyl ring.
Nevertheless, keto tautomers of phenols are frequently in-
voked as reactive intermediates, e.g., in the Photo–Fries
rearrangement, the Kolbe–Schmitt and Reimer–Tiemann re-
actions, in electrophilic substitutions of phenols, and in the
oxidative metabolism of aromatic compounds (“NIH-shift”).
In preliminary work (3), we generated cyclohexa-2,4-
dienone (2) by flash photolysis and observed the kinetics of
its enolization, 2 → 1, in acidic and neutral aqueous solu-
tions. The rate constant of ketonization, kK(1 → 2), was esti-
mated from the rate of deuteration of 1 in 1 N D2SO4 at
80°C. Combination with the rate constant of enolization,
kE(2 → 1), provided a preliminary estimate of the equilib-
rium constant of enolization of 2 in aqueous solution, KE ϵ
[1]/[2] = kE(2 → 1)/kK(1 → 2) ≈ 10(13 ± 1). We now report the
Received January 12, 1999.
This paper is dedicated to Jerry Kresge in recognition of his
many achievements in chemistry.
M. Capponi, I.G. Gut, B. Hellrung, G. Persy, and J.
Wirz.2 Institut für Physikalische Chemie der Universität
Basel, Klingelbergstrasse 80, CH-4056 Basel, Switzerland.
1Tautomerization of phenols, part 2; see ref. 3 for part 1.
2Author to whom correspondence may be addressed.
Telephone: +41 61 267 38 42. Fax: +41 61 267 38 55.
e-mail: wirz2@ubaclu.unibas.ch
Can. J. Chem. 77: 605–613 (1999)
© 1999 NRC Canada

606
Can. J. Chem. Vol. 77, 1999
Fig. 1. pH–rate profile for the enolization of 2,4-cyclo-
Scheme 1.
hexadienone (2) in aqueous solution at 25.0°C and I = 0.10 M.
5.5
log(kobs/s–1
)
5.0
4.5
4.0
3.5
3.0
O
OH
H2O
pH
2
4
6
8
10
12
complete pH–rate profile of the enolization reaction 2 → 1
in aqueous solution and an accurate value for KE. A brief
study of the other keto tautomer of 1, cyclohexa-2,5-dienone
(3), is included. The photoreactions that were used to gener-
ate dienones 2 and 3 are shown in Scheme 1.
profile of the enolization reaction 2 → 1 is shown in Fig. 1.
As is well known from many studies of ketone enolization
(2), the rate of enolization is accelerated by acid and by
base, but, unlike the behavior of simple ketones, the pH-
independent contribution dominates the overall reaction over
a large range, 3 ≤ pH ≤ 10. The solid line corresponds to
function [1], which was fitted to the observed rates of eno-
lization by nonlinear least-squares adjustment of the parame-
Results
Continuous irradiation of compounds 4–6
kE
,
ters kE and OH- the rate coefficients for acid and base
+
Direct irradiation of bicyclo[2.2.2]oct-2-ene-5,7-dione (4)
eliminates ketene to yield 2 as a primary product that tauto-
merizes to phenol (3). Oxadi-π-methane rearrangement is the
predominant reaction initiated by triplet sensitization of 4
(4). Irradiation of tricyclo[6.2.1.02,7]undeca-5,9-dien-4-one
(5) or tricyclo[6.2.1.02,7]undeca-4,9-dien-3-one (6) also gave
phenol as a major photoproduct that was identified by GC–
MS analysis. The flash photolysis experiments described be-
low provided direct evidence that phenol is formed via the
primary photoproducts 2 and 3 after light-induced cyclo-
elimination of cyclopentadiene from 5 and 6, respectively, as
indicated in Scheme 1.
catalyHsis, and
for the reaction catalyzed by the solvent
k0E
(Table 2). The dissociation constant of water at ionic
strength I = 0.1 M was taken as Kw = 1.59 × 10–14 M2 (5).
Weighting of the data was not required, because the relative
errors of the observed rate constants, and hence, standard er-
rors of log (kobs) values, were approximately constant.
[1]
log (koEbs͞s−1) = log {(k0E + kHE+ [H+]
+ kOEH K ͞[H+ ])͞s−1}
-
w
Acetate and tris(hydroxymethyl)methyl amine (Tris) buff-
ers were used to determine some rate constants at pH values
around 5 and 8. The decay rates increased linearly with
buffer concentration, and the intercepts obtained by least-
squares fitting of these individual plots are given in Table 1.
From the slopes of the buffer dilution plots only general
base catalysis is detectable, kAc = (9.5 ± 0.4) × 103 M–1 s–1,
kTris = (1.8 ± 0.2) × 105 M–1 s–1. Solvent isotope effects (Ta-
ble 2) were found to be in the normal direction (kH > kD) for
the decay rate constants k0E of cyclohexa-2,4-dienone and of
its 6,6-dideuterated derivative 2-d2, but strongly inverse on
the coefficient kHE+ of the acid-catalyzed reaction.
Flash photolysis; rates of enolization of the keto
tautomers 2 and 3
Aqueous solutions of compounds 4–6 were flashed with
25 ns pulses from an excimer laser operated on XeCl (200
mJ per pulse, 308 nm), or 20 µs pulses from a conventional
electric discharge flashlamp of up to 1000 J electrical en-
ergy. The cycloadducts 4 and 5, presumed to be photochemi-
cal precursors of 2, both gave rise to weak transient
absorption in the range of 280–330 nm. The decays of these
transients obeyed first-order kinetics with rate constants de-
pending on pH and buffer concentrations. Degassing had no
influence on the yield or lifetime of the transient intermedi-
ates. Ketones 4 and 5 are sensitive to base, but decomposi-
tion was sufficiently slow to permit measurements in freshly
prepared, weakly basic solutions. The decay rates of the
300 nm transients produced from 4 and 5 were equal, within
the limits of error, in neutral water (kobs = 3800 ± 300 s–1)
and faster, but again equal, in 2.0 mM aqueous KOH (7600
± 800 s–1).
Flash photolysis of 6 gave very weak transient absorption
around 250 nm, which is the wavelength expected for the
absorption of 2,5-cyclohexadienone (3) (6). Eight traces
were averaged for each measurement to improve the signal-
to-noise ratio. Due to the instability of 6 to base, experi-
ments were limited to acidic and neutral solution. The decay
rates of 3 were more than an order of magnitude slower than
kE
for 3
Because 4 gave somewhat stronger transient signals and
was available in larger quantity, this precursor was used to
collect a more extensive set of data at various pH (25°C,
ionic strength I = 0.1 M, Table 1). The resulting pH–rate
those of 2 (Table 1). The rate coefficients k0E and
+
were determined by nonlinear least-squares fitting oHf eq. [1]
(omitting the last term for base catalysis), and are also given
in Table 2.
© 1999 NRC Canada

Capponi et al.
607
Table 1. Rates of enolization of 2 and 3 in aqueous solution at 25.0 ± 0.1°C.
Compound
Acid/base
Buffer ratio
Conc./M
log ([H+]/M)
koEbs/(103 s–1)
No. of expts.
a
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
2
HClO4
HCl
HCl
HCl
HCl
HCl
HCl
HCl
HAc/NaAc
HAc/NaAc
HAc/NaAc
HAc/NaAc
HAc/NaAc
HCl/Tris
KOH
KOH
—
—
—
—
HCl
0.100
0.100
0.050
0.032
0.010
0.005
0.0025
0.001
0.02–0.10
0.01–0.09
0.01–0.09
0.01–0.09
0.01–0.09
0.01–0.10
0.002
0.004
—
1.00
1.00
1.30
1.50
2.00
2.30
2.60
3.00
3.87b
4.37b
4.87b
5.36b
5.87b
8.07c
11.10d
11.40d
≈6
73.6 ± 1.9
60.8 ± 1.5
30.6 ± 0.7
10
10
10
10
10
10
10
18
23
24
18
23
25
10
10
5
18.5 ± 0.3
7.65 ± 0.08
5.80 ± 0.08
4.80 ± 0.06
4.00 ± 0.06
3.82 ± 0.15e
3.90 ± 0.06e
3.98 ± 0.07e
3.99 ± 0.10e
3.95 ± 0.03e
3.80 ± 0.15e
7.59 ± 0.17
12.0 ± 0.6
3.83 ± 0.04
2.67 ± 0.02
0.58 ± 0.01
0.40 ± 0.01
0.62 ± 0.07
0.224 ± 0.010
0.079 ± 0.004
5.00
1.584
0.500
0.159
0.050
1.12
2
2
2-d2
2-d2
3
(D2O)
(D2O)
—
—
—
0.020
0.005
—
≈6
≈6
≈6
1.70
2.30
≈6
6
5
9
8
8
8
3
3
HCl
—
aValue calculated from the regression line of an Arrhenius plot (data in Table 4).
bCalculated using the thermodynamic acidity constant pKa = 4.76 (38), and activity coefficients recommended by Bates (5) to give the concentration
quotient pKa,c(I = 0.1 M) = 4.57.
s
cCalculated from pKa,c(I = 0.1 M) = 8.12 (39).
dCalculated from the nominal base concentration with Kw(I = 0.1 M) ϵ [H+][HO–] = 1.59 × 10–14 (ref. 5).
eIntercept of buffer dilution plot. From the slopes of the buffer dilution plots only general base catalysis is detectable, kAc = (9.5 ± 0.4)103 M–1 s–1, kTris
= (1.8 ± 0.2)105 M–1 s–1.
Table 2. Rate constants of enolization and ketonization at 25 ± 0.1°C, I = 0.10 M.
Reaction
Rate coefficient
k/s–1
(3.8 ± 0.1)103
(5.8 ± 0.1)102
(5.4 ± 0.3)105 M–1
(2.0 ± 0.3)106 M–1
80 ± 8
(2.8 ± 0.9)104 M–1
(1.01 ± 0.28)10–7 M–1
(2.96 ± 0.27)10–7 M–1
k(H2O)/k(D2O)
2 → 1
2-d2 → 1
2 → 1
2 → 1
3 → 1
3 → 1
1 → 2
1 → 3
k0E
k0E
kHE +
1.43 ± 0.03a
1.45 ± 0.05a
0.50 ± 0.05b
kOEH
-
k0E
kHE +
kHK+
kHK+
aAverage from ≥ 5 measurements each in H2O and D2O (no buffer); 25°C, I = 0.1 M.
bAverage from 9 measurements each of HCl/H2O and DCl/D2O at matching nominal acid concentrations in the
range of 0.02–0.1 N; 23°C, I = 0.1 M.
the isotope effect and to extrapolate to the rate constant kHK+
for the acid-catalyzed ketonization reaction 1 → 2 is now
described.
Acid-catalyzed exchange of aromatic hydrogen proceeds
by a two-step reaction mechanism with a C-protonated inter-
mediate (7), Scheme 2. Solutions of the substrate 2-L in H2O
are assumed to be dilute, such that that the second step, 2-L+
→ 1 + L+, is practically irreversible. Following the notation
Isotope exchange kinetics; rates of ketonization
At room temperature the exchange of the aromatic hydro-
gen atoms of 1 is slow. Solutions of phenol-2d and phenol-2t
in 0.1 N HClO4 were kept at 80.0 ± 0.5°C, and the loss of
their deuterium or tritium label to the solvent was monitored
2
by 400 MHz H NMR and by scintillation counting, respec-
tively (Table 3). Details of the procedures are given in the
experimental section. The analysis of these data to determine
© 1999 NRC Canada

608
Can. J. Chem. Vol. 77, 1999
Table 3. Isotopic exchange rates of phenol.
Reaction
Symbol
kobs/s–1
Acid
Acid conc./M
T/K
1–2t → 1
1–2d → 1
1 → 1-2d
1 → 1-2d
1 → 1-2d
1 → 1-4d
1 → 1-4d
1 → 1-4d
k–T
(5.63 ± 0.03)10–7
(9.04 ± 0.55)10–7
(3.22 ± 0.03)10–5
(1.11 ± 0.01)10–5
(3.48 ± 0.03)10–6
(5.39 ± 0.05)10–5
(1.93 ± 0.02)10–5
(7.02 ± 0.07)10–6
HClO4
HClO4
DCl
DCl
DCl
DCl
DCl
DCl
0.10
0.10
1.00
1.00
1.00
1.00
1.00
1.00
353.2
353.2
353.2
343.2
333.2
353.2
343.2
333.2
k–D
kD
2
kD
2
kD
2
kD
4
kD
4
kD
4
process is given by eq. [2]. The statistical factor of 2 arises
Scheme 2.
because the ketonization rate constant kK refers to
+
protonation at either of the ortho positionHs C2 or C6,
whereas H k1H refers to protonation at C2 only.
+ OH
OH
OH
H
L
Hk2
L
+ H+
L
H
Lk1
H
+ L+
[2]
kHK+ = 2H k1H
Lk2
H
Carbon protonation does, however, not fully determine the
rate of isotopic exchange, because only a fraction of the
protonated intermediate 2-L+ eventually loses the label by
dissociating to 1 + L+. The rates of the forward reaction
(loss of label) and backward reaction (retention of label) are
of the same order of magnitude, H k2L ≈ Lk2H. Since
protonation of 1–L is much slower than the subsequent
1–L
2–L+
1
O
L
H
+ H+
steps, L k1H[H+] << H k2L
+
L k2H, the steady-state (Bodenstein)
approximation is applicable to 2-L+, and the observed rate
constant for acid-catalyzed loss of label L may be expressed
as the product of the rate constant of C-protonation of 1-L,
the proton concentration, and the partitioning ratio of inter-
mediate 2-L+, eq. [3].
2–L
used by Kresge and Chiang (7), the rate coefficients shown
in Scheme 2 carry two superscripts to facilitate the book-
keeping of isotope effects. The right-hand superscript desig-
nates the isotope involved in the hydron transfer reaction
between C2 and the solvent (primary isotope effect). The
left-hand superscript designates the spectator isotope re-
maining on C2 during this reaction (secondary isotope ef-
fect).
H k2L
-L
kobs = k1H[H+ ]
L
[3]
H k2L
+
Lk2H
The ratio of the experimental rate constants of detritiation
and dedeuteration in aqueous acid of a given concentration
is then given by eq. [4].
-T
kobs Tk1H
H k2T
H k2D
+
Dk2H
Dissociation equilibria of oxygen acids are, in general, es-
tablished more rapidly than those of carbon acids (8). Ac-
cordingly, we assume that OH dissociation of 2-L+ (the
vertical reaction shown in Scheme 2) is faster than hydron
dissociation from carbon to form 1 or 1-L (horizontal reac-
tions). This assumption conforms to the accepted mechanism
for hydronium-ion-catalyzed enolization of simple carbonyl
compounds, which consists of rapid and reversible
hydronation of the substrate on oxygen followed by rate-
determining removal of a carbon-bound hydron by a mole-
cule of water (9). Application to the present case might be
questioned on the grounds that CH dissociation of 2-L+ is
thermodynamically favored over OH dissociation by 2.3RT
× pKE which, as we shall see below, amounts to 72.6 ± 0.7
kJ mol–1. However, we have recently shown that even larger
thermodynamic bias is insufficient to overcome the kinetic
preference for OH dissociation (10). Moreover, the inverse
solvent isotope effect observed on the acid-catalyzed
enolization of 2 (kH/kD = 0.5, Table 2, row 3) is strong evi-
dence for pre-equilibrium protonation on oxygen.
[4]
=
-D
T
kobs Dk1H H k2T + k2H
H k2D
Tk1H 1 + Dk2H͞Hk2D
=
T
Dk1H 1 + k2H͞Hk2T
Our goal is to calculate the rate constant for acid-
catalyzed ketonization of 1, kHK+ , from the observed rates of
-D
dedeuteration and detritiation, kobs and k-oTbs. Extrapolation
from observed isotope exchange rates kT and kD to kH is
commonly done on the basis of the Swain relation (eq. [5])
(11). However, eq. [4] contains contributions from both pri-
mary and secondary kinetic isotope effects. In a detailed
study of acid-catalyzed hydron exchange in 1,3,5-
trimethoxybenzene, Kresge and Chiang (7) have established
that the Swain relation gives reliable results when it is ap-
plied individually to primary and secondary isotope effects.
[5]
(kH/kD)1.442 = kH/kT
The first factor on the right-hand side of eq. [4],
Tk1H͞Dk1H, is a purely secondary isotope effect: hybridization
of carbon atom 2 changes from sp2 to sp3 upon protonation
(step 1 in Scheme 2), which influences bonding to the isoto-
If OH dissociation of 2-L+ is faster than CH dissociation,
then carbon protonation is the rate-determining step in acid-
catalyzed ketonization of 1 → 2, and the rate constant of this
© 1999 NRC Canada

Capponi et al.
609
Table 4. Temperature dependence of the rate
constants of enolization of 2 in 0.1 N HClO4.
The temperature dependence of the isotopic exchange
rates was determined by monitoring the deuteration of phe-
nol in 1 N DCl/D2O at 60.0, 70.0, and 80.0 ± 0.5°C by
300 MHz 1H NMR. This was easier than repeating the
dedeuteration of phenol-2d at several temperatures, and at
the same time, provided exchange rates at C4 of phenol.
Secondary isotope effects from isotopic substitution at posi-
tions remote from the site of protonation (C4 and C6 for
deuteration at C2) have been demonstrated to be negligible
log (koEbs/s–1)
t/°C
4.85 ± 0.02
5.05 ± 0.02
5.21 ± 0.02
5.37 ± 0.02
5.49 ± 0.02
5.59 ± 0.02
25.0 ± 0.1
37.0 ± 0.1
47.0 ± 0.1
62.5 ± 0.2
71.0 ± 0.2
79.0 ± 0.3
(7). The resulting exchange rate constants kD and kD are
given in Table 3. These values are two to three times higher
than those obtained in preliminary work by deuteration of
phenol in 1 N D2SO4 at the same temperature (3). The dis-
crepancy may in part be attributed to incomplete dissocia-
tion of D2SO4 in 1 N solution. Linear least-squares
regression of log (k/s–1) vs. 1/T gave the Arrhenius parame-
ters log (A/s–1) = 11.61 ± 0.13 and Ea = 109 ± 1 kJ mol–1 for
ortho-deuteration, and log (A/s–1) = 10.47 ± 0.32 and Ea =
100 ± 2 kJ mol–1 for para-deuteration. The rate constants
2
4
pic label L. Such secondary isotope effects are known to be
inverse, and to have a relatively constant value of about 12%
per deuterium atom, i.e., H k1H͞Dk1H = 0.88 ± 0.04 (7, 12).
The factor Tk1H͞Dk1H can then be estimated using the Swain
relation
(1.442−1)
Tk1H
Dk1H
Tk1H H k1H
H k1H Dk1H
Dk1H
H k1H
[6]
=
=
calculated for 25°C from these regressions are kD = (3.47 ±
2
0.16)10–8 M–1 s–1 and kD = (1.02 ± 0.12)10–7 M–1 s–1.
4
= (0.88 ± 0.04)-0.442 = 1.06 ± 0.02
The ratio between the rate constant of ketonization 1 → 2,
eq. [8], and the observed rate constant of ortho-deuteration
amounts to
The fractions appearing in the second term of eq. [4] are a
product of primary and secondary isotope effects. In the nu-
merator we have Dk2H͞Hk2Dϵx. By application of the Swain
relation we obtain Tk2H͞Hk2T = x1.442, and eq. [4] is replaced
by eq. [7].
kHK+͞kD
= (8.1 ± 1.3)10–5/(3.22 ± 0.03)10–5
=
2
2.5 ± 0.4 at 80°C. This factor is a compound of primary and
secondary isotope effects of the substrate as well as ionic
strength and isotope effects of the solvent. The temperature
dependence of this factor is estimated from the formula
kHK+͞kD = exp (–∆E0/RT), assuming that it may be attributed
-T
2
kobs
1 + x
1 + x1.442
[7]
= (1.06 ± 0.02)
-D
kobs
to a single dominant zero-point vibrational energy difference
∆E0. That is an overly crude assumption, but the correction
of +0.4 for extrapolation to room temperature is small; the
The experimental rate constants for detritiation and
dedeuteration of 1-2l at 80°C in 0.1 N HClO4 are given in
Table 3. Equation [7] can be solved for x by iterative ap-
proximation to give x = 4.1 ± 0.5. Insertion of x into eq. [3]
ratio resulting for 25°C is kHK+͞kD = 2.9 ± 0.8 (the error
2
margin was increased by 0.4, the amount of the above cor-
rection), and the rate constant for acid-catalyzed
ketonization 1 → 2 at 25°C is then calculated as kHK+ = (2.9 ±
0.8)(3.47 ± 0.16)10–8 = (1.01 ± 0.28)10–7 M–1 s–1. The same
factor was used to calculate the acid-catalyzed rate constant
(L = D) gives Dk1H = ko-Dbs(1 + x)/[H+] and using H k1H
=
Dk1H(0.88 ± 0.04), we can replace eq. [2] by eq. [8], in
which the desired rate constant of ketonization is expressed
fully in terms of experimentally determined quantities
of ketonization 1 → 3, kHK+ = (2.9 ± 0.8)kD = (2.96 ±
4
0.27)10–7 M–1 s–1.
[8]
kHK+ = 2k-oDbs(x +1)(0.88 ± 0.04)͞[H+ ]
kHK+
We obtain
(353 K) = (8.1 ± 1.3)10–5 M–1 s–1 for the
Discussion
rate constant of acid-catalyzed ketonization, 1 → 2. The er-
ror limits were determined by propagation of the standard
errors of the experimental data and of the assumed range for
Evidence for keto tautomers of phenol
In a review article on the theory of unsaturated and aro-
matic compounds, published in 1899, Thiele (13) speculated
that the propensity of phenol (1) and its anion to undergo
substitution and oxidation reactions might be attributed to
rapid equilibration of 1 with the transient keto forms 2 and
3. At that time, the concepts of protomerism, as in
Scheme 3, and mesomerism, Scheme 4, were not clearly dis-
tinguished. From the present viewpoint, it is clear that the
ketonization of phenol, Scheme 3, is much too slow to ex-
plain the reactivity of phenol towards, e.g., bromine, but
mesomeric charge delocalization, as shown in Scheme 4, is
still used to explain the enhanced reactivity of phenolate at
the ortho and para positions.
the secondary isotope effect, taking account of the high
-D
covariance between x and k through eq. [7].
obs
Temperature dependence of the kinetic data
Flash photolysis of 4 in 0.1 N aqueous HClO4 was done at
various temperatures from 25 to 80°C (Table 4). The rate
constants accurately obeyed the Arrhenius equation, k = A
exp (–Ea/RT), and linear regression of log (koEbs/s–1) vs. 1/T
gave the Arrhenius parameters log (A/s–1) = 9.58 ± 0.11 and
Ea = 27.0 ± 0.7 kJ mol–1. The rate coefficient for the acid-
catalyzed enolization of 2 at 80°C is calculated as kE
=
+
koEbs/[H+] = (3.94 ± 0.06)106 M–1 s–1 from this regresHsion.
Combination with the rate constant of ketonization provides
the equilibrium constant of enolization of 2 at this tempera-
ture, KE(353 K) = kHE+͞kHK+ = (3.94 ± 0.06)106/(8.1 ± 1.3)10–5
= (4.9 ± 0.8)1010, pKE(353 K) = –10.69 ± 0.07.
Only 2 years later, Lapworth (14) put forth the hypothesis
that bromocyclohexadienones are the primary products
formed in the aqueous bromination of phenol to rationalize
the exclusive bromination of 1 at the o- and p-positions.
© 1999 NRC Canada

610
Can. J. Chem. Vol. 77, 1999
Scheme 3.
Scheme 5.
OH
OH
O
O
O
H
H
pKE = –12.73 ± 0.12
–
2
1
O
H
H
pKaE = 9.84 ± 0.02
pKaK = –2.89 ± 0.12
1
2
3
Scheme 4.
–
O
O
O
common cation (24). ∆2→1H (g, 298 K) = –15.2 kcal mol–1
s
–
is obtained using Benson’s group increments (25). Tee and
s
Iyengar (15) predicted ∆3→1G (aq, 298 K) = –15 kcal mol–1
based on an earlier estimate of the pK for C-protonated phe-
nol. Shiner et al. (26) generated the cyclohexadienones 2
and 3 by pyrolytic cycloreversion of their formal Diels–Al-
der adducts with cyclopentadiene, compounds 5 and 6, and
measured their gas phase acidities by the flowing afterglow
–
These bromocyclohexadienone intermediates were detected
15 years ago by Tee et al. (15) using the stopped-flow tech-
nique. Cyclohexadienones were also proposed (16) and later
characterized (17, 18) as intermediates in the Photo–Fries re-
arrangement of aryl esters. 4-Benzylcyclohexa-2,5-dienone
was identified by CIDNP as a transient intermediate in the
photorearrangement of benzyl phenyl ether; it had a half-life
of about 4 s in cyclohexane (19).
4H-Keto tautomers of heavily substituted phenols have
been isolated in crystalline form (20). The tautomers of par-
ent phenol, cyclohexa-2,4-dienone (2), and cyclohexa-2,5-
dienone (3), were generated by flash vacuum pyrolysis of
suitable Diels–Alder adducts, trapped at 77 K, and charac-
terized by their IR and UV spectra (21). Pentaammineosmium
complexes of 1–3 have been identified, and it was shown
that all three tautomeric complexes are of comparable ther-
modynamic stability (22).
In the present work, cyclohexadienones 2 and 3 were formed
by flash photolysis of suitable precursors (4–6, Scheme 1) in
aqueous solution, which allowed their enolization rates to be
measured directly. The transient formed by flash photolysis
of 4 has previously been identified as cyclohexa-2,4-dienone
(2) based on the observed decay kinetics, the position of its
absorption (the first π,π* absorption band of stable cyclo-
hexa-2,4-dienones is around 300 nm, log ⑀ ≈ 3.7 (6)), the
identification of phenol as the final product, and the forma-
tion of 1,3,5-hexatrienone upon irradiation of 4 in a glassy
matrix (3). The present finding that precursor 5 gives the
same transient intermediate (absorption and pH-dependent
decay rate) confirms that assignment. Formation of
cyclohexa-2,5-dienone (3) from precursor 6 was assumed by
analogy, and that is supported by the observed transient ki-
netics and the identification of phenol as the final product.
s
technique. The derived heats of tautomerization, ∆2→1H =
– 6 ± 3 and ∆3→1H = –10 ± 3 kcal mol–1 are, however, not
s
compatible with present data.
From the present data, the enolization constant of 2 at
25°C is determined as the ratio of the enolization and ket-
onization rate constants given in Table 2, KE(2, aq, 298 K) =
kHE+͞kHK+
pKE
= (5.4 ± 0.3)105/(1.01 ± 0.28)10–7 = (5.4 ± 1.5)1012,
= –12.73 ± 0.12. This value, and the acidity constant of
phenol, pKaE(1, 298 K) = 9.84 ± 0.02 (expt. sect.), form two
legs of a thermodynamic cycle, from which the C–H acidity
constant of ketone 2 can be calculated, pKaK = pKE + pKaE =
–2.89 ± 0.12 (Scheme 5). It turns out that ketone 2 is one of
the strongest carbon acids known (27).
The enthalpy of enolization can be calculated by the sec-
ond-law method from the temperature dependence of the
s
enolization constant, d ln KE/dT = ∆EH /RT2, ∆EH
s
=
2.303R[pKE(T2 = 353 K) – pKE(T1 = 298 K)]/(1/T2 – 1/T1) =
–74.8 ± 5.1 kJ mol–1. The standard free energies of
s
enolization are ∆EG (aq) = 2.303RT × pKE(353 K) = –72.3 ±
0.5 kJ mol–1 at 353 K and –72.7 ± 0.7 kJ mol–1 at 298 K.
s
The entropy of enolization, ∆ES (aq, 298 K) = –7 ± 16 J K–1
mol–1, is very small, too small to be reliably determined
from these values. In our preliminary communication (3),
we had estimated the temperature dependence of the
enolization constant by neglecting the temperature depend-
s
ence of the free energy of enolization ∆EG , i.e., by assum-
s
ing ∆ES ≈ 0.
The enolization constant of 2,5-cyclohexadienone 3 is
similarly obtained as the ratio of the corresponding acid-
catalyzed rate constants of enolization and ketonization
(Table 2), KE(298 K) = kHE+͞kHK+ = (9.5 ± 3.2)1010, pKE =
–10.98 ± 0.15 (Scheme 6).
The thermochemical data obtained in this work are sum-
marized in Table 5. Note that the acidity constants pKaE and
pKaK reported here are concentration quotients determined at
ionic strength I = 0.1 M. The experimental enolization con-
stants are in reasonable, though not perfect, agreement with
the results of density functional calculations given in
Table 6. The (gas phase) enolization constants predicted by
these calculations are pKE(2) = –10.4 and pKE(3) = –9.8.
Thermochemistry of cyclohexadienone–phenol
tautomerization
Several thermochemical estimates have been reported for
the protomeric equilibria of phenol (1) with ketones 2 and 3
s
(Scheme 3): ∆(2 or 3)→1G (g, 298 K) = –18.6 kcal mol–1 was
obtained from
∆
a
bond additivity scheme (23), and
2→1G (aq, 298 K) = –13 ± 3 kcal mol–1 from estimates of
the aqueous protonation equilibria for each tautomer to the
s
© 1999 NRC Canada

Capponi et al.
611
Table 5. Thermochemical data (aq, 25°C, I = 0.1 N).
Reaction
pKE
pKaK
pKaE
∆EH/(kJ mol–1)
∆ES/(J K–1 mol–1)a
2 → 1
3 → 1
–12.73 ± 0.12
–10.85 ± 0.29
–2.89 ± 0.12
–1.01 ± 0.30
9.84 ± 0.02
9.84 ± 0.02
–74.8 ± 5.1
–61.9 ± 6.2b
–2.6
–1.2
aCalculated B3LYP/6-31G* (g, 298 K), cf. Table 6.
bCalculated from pKE assuming |∆ES(aq)| ≤ 20 J K–1 mol–1.
Table 6. Results of ab initio calculations.a
Compound
Method
–E0/hartree
Enthalpy/hartreeb
S /(J K–1 mol–1)
s
1
2
3
1
2
3
B3LYP/6-31G*
B3LYP/6-31G*
B3LYP/6-31G*
6-31G*
6-31G*
6-31G*
–307.464 865
–307.441 132
–307.442 838
–305.558 063
–305.538 575
–305.541 889
0.111 231
0.110 405
0.110 618
0.117 496
0.117 177
0.117 385
74.55
77.14
75.76
73.58
75.63
74.65
aGaussian 94.
bThermal correction to enthalpy including zero-point vibrational energy.
Scheme 6.
OH
H2O and D2O (Table 2) give two closely similar values
for the predominantly primary kinetic isotope effect,
k(2-h2)/k(2-d2) = 6.6 ± 0.2, and for the compound (second-
ary and medium) solvent isotope effect, k(H2O)/k(D2O) =
1.44 ± 0.03. Both of these values are consistent with the
mechanism assigned to the pH-independent enolization of
ketone 2: rate-determining hydron transfer from carbon to
solvent water. A crude a priori estimate for the solvent iso-
tope effect, k(H2O)/k(D2O) ≈ 1.5, is obtained by assuming
KaK(H2O)/KaK(D2O) ≈ 1/l3 = 3 (l = 0.69 is the fractionation
factor for the hydronium ion (29)) and application of the
Brønsted equation k(H2O)/k(D2O) = [KaK(H2O)/KaK(D2O)]α
with α ≈ 0.35 (2b). Finally, the inverse isotope effect ob-
served on the acid-catalyzed enolization of 2-L (Table 2,
row 3), shows that OH fission of 2-L+ (vertical reaction in
Scheme 2) is faster than CH or CL fission, in spite of a
strong thermodynamic bias for the latter.
O
pKE = –10.98 ± 0.15
–
1
3
O
pKaE= 9.84 ± 0.02
pKaK= –1.14 ± 0.15
Exchange kinetics and kinetic isotope effects
Isotope exchange rates of phenol in aqueous sodium hy-
droxyde were already determined in 1936 by Ingold et al.
(28). These authors used 2.2 M solutions of phenol and
found that the exchange rate reached a maximum when the
concentration of NaOH was about half that of phenol.
Hydron transfer from phenol as a general oxygen acid to
phenolate as a carbon base was identified as the dominant
mechanism for exchange under these conditions. In the
present work, hydron exchange experiments were done
with dilute solutions and under acid catalysis. Note that,
from the data given in Tables 2 and 5, base is predicted to
reduce the rate of hydron exchange in dilute aqueous solu-
tions. The rate of hydrogen exchange at carbon 2 of phenol
The range covered by experimental data on keto–enol pro-
tomeric equilibria is nearly doubled by the present measure-
ments of phenol ketonization: enolization constants now extend
over a range of about 30 orders of magnitude, and the corre-
sponding rate constants vary by about 20 orders of magni-
tude. The free-energy relationship between these quantities
is clearly curved, and the curvature defines the intrinsic bar-
rier of proton transfer to the β-carbon atom of enols ∆G0‡
=
57 kJ mol–1 (2b). Similar values have been determined previ-
ously from curvature in Brønsted plots (30, 31). The pH-
independent, spontaneous dissociation of 2 as a carbon acid,
k0E = 3.8 × 103 s–1, dominates the pH–rate profile of 2
(Fig. 1) from pH 3 to 10. On the other hand, the contribution
of this mechanism to the pH–rate profiles for the enolization
of simple ketones such as acetophenone is barely noticeable
(31), because carbon protonation of enolates is so highly
exergonic that these rates approach the diffusional limit.
kE K ͞KK
=
E
K
′
in strong base, pH >> pKa (1), is given by k0
-
w
a
= 4 × 10–11 s–1 at 25°C. This is more than an orderOoHf magni-
tude less than the rate calculated for the same process in
neutral solution, k0K = k0EKaE͞KaK = 7 × 10–10 s–1, because the
reaction of protons with phenolate accelerates ketonization
at pH < 12 until proton catalysis saturates at pH ≈ pKaE
=
9.84.
The method used to assess the magnitude of the kinetic
isotope effects and to determine the rate constants of
ketonization from the observed rates of hydron exchange of
phenol in aqueous acid closely follows previous work of
Kresge and Chiang (7) and was described in the results sec-
tion. Some of the observed kinetic isotope effects deserve
mention. The four decay rates determined for 2 and 2-d2 in
Experimental
General
All concentrations are nominal values referring to 25°C.
No corrections were applied for volume expansion at higher
© 1999 NRC Canada

612
Can. J. Chem. Vol. 77, 1999
temperatures. The required, small corrections would affect
acid-catalyzed rate constants of enolization and ketonization
by the same factor that would cancel in their ratio, the
enolization constant.
was kept at 80.0 ± 0.1°C. Deuterium loss from phenol-2d in
samples removed periodically from the bath was determined
2
by integration of the two peaks observed in the H NMR
spectra, at δ 1.13 (HDO) and at 1.28 ppm (phenol-2d). The
first six samples covered two half-lives, the last was drawn
after seven half-lives.
Materials
Samples of tricyclo[6.2.1.02,7]undeca-5,9-dien-4-one (5) and
tricyclo[6.2.1.02,7]undeca-4,9-dien-3-one (6) were obtained
as a gift from Dr. C. Shiner, University of Colorado, Boulder
(23), and samples of 5,7-dioxobicyclo[2.2.2]octane-2,3-
dicarboxylic acid from Profs. M. Demuth, MPI Göttingen,
and C.A. Grob, University of Basel.
Deuteration rates of phenol in 1.0 N DCl in D2O were de-
1
termined by 300 MHz H NMR. The NMR tubes (10 for
each temperature) were heated in a water bath which was
held at 80.0, 70.0, or 60.0 ± 0.1°C, removed at appropriate
time intervals, and stored at in a refrigerator until all sam-
ples had been collected. The signal of the meta hydrogens (δ
= 7.4 ppm; dd, J2,3 = J3,4 = 7 Hz, 2H) was used as an inter-
nal standard. First-order exponential decays were fitted to
the integrated signal intensities of the ortho (δ = 6.9 ppm; d,
J2,3 = 7 Hz, 2H) and para (δ = 7.1 ppm; t, J3,4 = 7 Hz, 1H)
hydrogens (Table 3).
The contents of phenol-2t were determined by liquid scin-
tillation counting of 31 aliquots of 0.5 mL containing 7 mM
of phenol which were drawn periodically over three half-
lives of the exchange reaction. A final sample was measured
after 10 half-lives. The aliquots were extracted with 4 mL of
diethyl ether to separate the tritiated phenol from the tritium
that had exchanged into water. To correct for variations
(±10%) in the efficiency of the extraction, the absorbance of
each extract was measured at 274 and 281 nm, and the ra-
dioactivity of the samples were normalized by these
absorbances. In this way, the standard error of single mea-
surements was reduced to about 3%. The activity of the sam-
ples was determined by scintillation counting on a Packard
scintillation analyzer CA 2000 using 1,4-bis(5-phenyl-2-
oxazolyl)benzene (POPOP) and Packard emulsifier 299 as
scintillation reagents. Prior to the measurements, the instru-
ment was calibrated with a standard of known activity.
Phenol-2d was synthesized from 2-bromophenol as de-
scribed by Tashiro et al. (32). 1H NMR indicated >90%
deuteration at the 2-position. The same procedure was used
to synthesize phenol-2t: NaOH (500 mg) was dissolved in
1 mL of tritium-enriched water and 520 mg of 2-bromo-
phenolate (2.7 mmol) was added. CuAl catalyst (125 mg,
Ventron, Karlsruhe) was added. The solution was stirred vig-
orously for 30 min. The catalyst was removed by filtration.
The solution was acidified with 2 N HCl whereupon phenol
precipitated. The aqueous solution was extracted with 10 mL
of ether. The organic extract was dried with Na2SO4, and the
solvent was evaporated at room temperature under a stream
of N2. The residue was sublimed in vacuo (<1 Pa) at room
temperature. A sample containing 0.6 mg (6.4 µmol) of phe-
nol-2t had an activity of 1750 Bq on a calibrated scintilla-
tion counter which corresponds to a tritium content of
0.25 ppm.
Bicyclo[2.2.2]oct-7-ene-2,5-dione (4) was prepared by ox-
idative bisdecarboxylation of 5,7-dioxobicyclo[2.2.2]octane-
2,3-dicarboxylic acid, samples of which were obtained as
gifts from Profs. Grob (33) and Demuth (5). The following,
improved method was taken from the thesis of Weitemeyer
(34). Pb(OAc)4 (18 g) was added to a stirred solution of the
diacid (4.48 g) in pyridine (60 mL) which was immersed in
a water bath at 20°C. The yellow solution slowly turned or-
ange to deep red under evolution of CO2. After an hour, the
solution was gently warmed to 40°C until the gas evolution
ceased after 90 min. CHCl3 (200 mL) was added and the
mixture shaken with 40 mL of cold 20% HCl. The precipi-
tate was removed and washed with CHCl3. The combined
organic phases were washed four times with 40 mL of 20%
HCl until the color was removed, neutralized with NaHCO3,
and dried. Evaporation of the solvent left a brown–red solid
residue which was chromatographed on silica with an 80:20
mixture of hexane and diethyl ether and sublimed at 90°C in
vacuo to give 560 mg (20.6%) of 4 as colorless crystals.
A sample of 4 (50 mg) was deuterated in 2 mL of D2O
containing 0.1 g of NaOD. The solution turned yellow. After
10 min the diketone was extracted with ether. The organic
phase was washed with D2O, evaporated to dryness, and
sublimed at 80°C in vacuo, giving off-white crystals of 4-
Flash photolysis
All kinetic measurement were done with water-jacketed
quartz cells which were kept at 25.0 ± 0.1°C unless stated
otherwise. Temperature readings were taken from
a
thermoelement that was dipped into the sample solutions.
Acid and base strength was adjusted with HCl in the range
of 0.1 to 0.001 N or KOH in the range of 0.01 to 0.001 N.
Values from buffered solutions quoted in this paper were ob-
tained by extrapolation to zero buffer concentration and thus
refer to wholly aqueous solutions. All solutions were ad-
justed to ionic strength I = 0.1 M by the addition of NaCl.
Kinetic data determined by different persons using different
experimental setups differed by up to 20%, much more than
the standard error of sample means determined by a single
person in a series of experiments (cf., e.g., the first two en-
tries of Table 1).
Spectrophotometric titration of phenol
1
6,6,8,8-d4 (40 mg). H NMR (D2O) δ: 2.8 (m, <0.01 H), 3.8
A 0.3 mM solution of phenol in water (3 mL) containing
0.010 M NH4Cl and 0.09 M NaCl was titrated with 30 µL
portions of 0.1 N NaOH in a quartz cell which was left un-
touched in the thermostatted (25 ± 0.1°C) sample holder of a
Perkin–Elmer Lambda 9 UV spectrophotometer. The UV
spectra were recorded and digitized (250–310 nm, 120
absorbance readings per spectrum, 16 spectra per titration)
after each addition and the acid concentration was measured
(m, 1 H, > CH-), 6.8 ppm (m, 1 H, -CHϭ); solvent peak
(HDO) δ: 5.0 ppm. MS: 141 (8.8), 140 (100, M+), 139 (9.1).
Isotopic exchange experiments
The samples of phenol-2d (7 NMR tubes containing
0.5 mL, 0.04 M) and phenol-2t (20 mL, 0.007 M) in 0.1 N
HClO4 were heated simultaneously in a thermostat bath that
© 1999 NRC Canada

Capponi et al.
613
13. J. Thiele. Liebigs Ann. Chem. 306, 87 (1899); 306, 129 (1899).
14. A. Lapworth. J. Chem. Soc. 1265 (1901).
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(1983); O.S. Tee and N.R. Iyengar. J. Am. Chem. Soc. 107,
455 (1985); J. Org. Chem. 51, 2585 (1986).
16. J.W. Meyer and G.S. Hammond. J. Am. Chem. Soc. 94, 2219
(1972).
17. C.E. Kalmus and D.M. Hercules. J. Am. Chem. Soc. 96, 449
(1974); T. Arai, S. Tobita, H. Shizuka. J. Am. Chem. Soc. 117,
3968 (1995).
18. S.M. Beck and L.E. Brus. J. Am. Chem. Soc. 104, 1805 (1982).
19. R. Bausch, H.-P Schuchmann, C. von Sonntag, R. Benn, and
H. Dreeskamp. J. Chem. Soc. Chem. Commun. 418 (1976).
20. V.V. Ershov and A.A. Volod’kin. Izv. Akad. Nauk. SSSR Otd.
Khim. Nauk. 730 (1962); T. Matsuura and K. Ogura. J. Am.
Chem. Soc. 89, 3846 (1967).
with a calibrated (35) Metrohm 6.0236.100 micro glass elec-
trode using a digital Metrohm pH meter 654. The spectra
were analyzed using the program SPECFIT of Zuberbühler
and co-workers (36). The original spectra were reproduced
with a standard error of <0.002 absorbance units using the
two dominant eigenvectors of the factor analysis. The load-
ing coefficients fitted well to a titration curve (standard
errors of single runs were <0.001 pK units). Three independ-
ent runs gave an average dissociation quotient of phenol,
pKa,c = 9.84 ± 0.02 (ionic strength I = 0.1 M). Using activity
coefficients for I = 0.1 M (γH + = 0.83 and γ1 = 0.76 (7)),
−
the thermodynamic acidity constant of phenol is estimated
s
as pKa = 10.04 ± 0.03. This is about 0.05 units higher than
the literature values commonly quoted (37).
21. M.-C. Lasne, J.-P. Ripoll, and J.-M. Denis. Tetrahedron Lett.
21, 463 (1980); V.P. Vitullo. J. Org. Chem. 35, 3975 (1970).
22. M.E. Kopach, W.G. Hipple, and W.D. Harman. J. Am. Chem.
Soc. 114, 1736 (1992).
23. J.B. Conant and G.B. Kistiakowsky. Chem. Rev. 20, 181 (1937).
24. M.J. Cook, A.R. Katritzky, P. Linda, and R.D. Tack. Tetrahe-
dron Lett. 49, 5019 (1972).
25. S.W. Benson. Thermochemical kinetics. 2nd ed. J. Wiley, New
York. 1976.
26. C.S. Shiner, P.E. Vorndam, and S.R. Kass. J. Am. Chem. Soc.
108, 5699 (1986).
27. E.D. Laganis and D.M. Lemal. J. Am. Chem. Soc. 102, 6633
(1980).
28. C.K. Ingold, C.G. Raisin, and C.L. Wilson. J. Chem. Soc.
1637 (1936).
29. A.J. Kresge, Y. Chiang, G.W. Koeppl, and R.A. More O’Ferrall.
J. Am. Chem. Soc. 99, 2245 (1977).
30. P. Pruszynski, Y. Chiang, A.J. Kresge, N.P. Schepp, and P.A.
Walsh. J. Phys. Chem. 90, 3760 (1986).
Acknowledgements
This work was supported the Swiss National Science
Foundation. Helpful advice for the analysis of isotope ef-
fects by Prof. A.J. Kresge is gratefully acknowledged.
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