Equilibrium constants for dehydration of water adducts of aromatic carbon-carbon double bonds

Article abstract of DOI:10.1021/ja0126125

Equilibrium constants (K_de) are reported for the dehydration of hydrates of benzene, naphthalene, phenanthrene, and anthracene. Free energies of formation of the hydrates (ΔG°_f(aq)) are derived by combining free energies of formation of the parent (dihydroaromatic) hydrocarbon with estimates of the increment in free energy (ΔG^OH) accompanying replacement of a hydrogen atom of the hydrocarbon by a hydroxyl group. Combining these in turn with free energies of formation of H₂O and of the aromatic hydrocarbon products furnishes the desired equilibrium constants, The method depends on the availability of thermodynamic data (i) for the hydrocarbons from which the hydrates are derived by hydroxyl substitution and (ii) for a sufficient range of alcohols to assess the structural dependence of ΔG^OH. The data comprise chiefly heats of formation and standard entropies in the gas phase and free energies of transfer from the gas phase to aqueous solution (the latter being derived from vapor pressures and solubilities). They also include experimental measurements of equilibrium constants for dehydration of alcohols, especially cyclic, allylic, and benzylic alcohols. In general ΔG^OH depends on whether the alcohol is (a) primary, secondary, or tertiary; (b) allylic or benzylic; and (c) open chain or cyclic. Differences in geminal interactions of the hydroxyl group of the alcohol with α-alkyl and vinyl or phenyl groups account for variations in ΔG^OH of 5 kcal mol^-1. Weaker variations which arise from β-vinyl/OH or β-phenyl/OH interactions present in the aromatic hydrates but not in experimentally studied analogues are estimated as 1.0 kcal mol^-1. Equilibrium constants for dehydration may be expressed as their negative logs (pK_de). Reactions yielding the following aliphatic, aromatic, and antiaromatic unsaturated products then have pK_de values: +4.8, ethene; +15.0, ethyne; +22.1, cyclopropene; +28.4 cyclobutadiene; -22.2, benzene; -14.6, naphthalene; -9.2, phenanthrene; -7.4, anthracene. Large positive values are associated with formation of strained or antiaromatic double bonds and large negative values with aromatic double bonds. Trends in pK_de parallel those of heats of hydrogenation. The results illustrate the usefulness of a substituent treatment for extending the range of currently available free energies of formation. In addition to hydroxyl substituent effects, ΔG^OH, values of ΔG^π for substitution of a π-bond in a hydrocarbon are reported.

Full text of DOI:10.1021/ja0126125

Published on Web 06/27/2002
Equilibrium Constants for Dehydration of Water Adducts of
Aromatic Carbon-Carbon Double Bonds
Joykrishna Dey, AnnMarie C. O’Donoghue, and Rory A. More O’Ferrall*
Contribution from the Department of Chemistry and Centre for Synthesis and Chemical Biology,
UniVersity College Dublin, Belfield, Dublin 4, Ireland
Received November 28, 2001
Abstract: Equilibrium constants (Kde) are reported for the dehydration of hydrates of benzene, naphthalene,
phenanthrene, and anthracene. Free energies of formation of the hydrates (∆G°f(aq)) are derived by
combining free energies of formation of the parent (dihydroaromatic) hydrocarbon with estimates of the
increment in free energy (∆GOH) accompanying replacement of a hydrogen atom of the hydrocarbon by a
hydroxyl group. Combining these in turn with free energies of formation of H2O and of the aromatic
hydrocarbon products furnishes the desired equilibrium constants. The method depends on the availability
of thermodynamic data (i) for the hydrocarbons from which the hydrates are derived by hydroxyl substitution
and (ii) for a sufficient range of alcohols to assess the structural dependence of ∆GOH. The data comprise
chiefly heats of formation and standard entropies in the gas phase and free energies of transfer from the
gas phase to aqueous solution (the latter being derived from vapor pressures and solubilities). They also
include experimental measurements of equilibrium constants for dehydration of alcohols, especially cyclic,
allylic, and benzylic alcohols. In general ∆GOH depends on whether the alcohol is (a) primary, secondary,
or tertiary; (b) allylic or benzylic; and (c) open chain or cyclic. Differences in geminal interactions of the
hydroxyl group of the alcohol with R-alkyl and vinyl or phenyl groups account for variations in ∆GOH of 5
kcal mol-1. Weaker variations which arise from â-vinyl/OH or â-phenyl/OH interactions present in the aromatic
hydrates but not in experimentally studied analogues are estimated as 1.0 kcal mol-1. Equilibrium constants
for dehydration may be expressed as their negative logs (pKde). Reactions yielding the following aliphatic,
aromatic, and antiaromatic unsaturated products then have pKde values: +4.8, ethene; +15.0, ethyne;
+22.1, cyclopropene; +28.4 cyclobutadiene; -22.2, benzene; -14.6, naphthalene; -9.2, phenanthrene;
-7.4, anthracene. Large positive values are associated with formation of strained or antiaromatic double
bonds and large negative values with aromatic double bonds. Trends in pKde parallel those of heats of
hydrogenation. The results illustrate the usefulness of a substituent treatment for extending the range of
currently available free energies of formation. In addition to hydroxyl substituent effects, ∆GOH, values of
∆Gπ for substitution of a π-bond in a hydrocarbon are reported.
Scheme 1
Introduction
The aim of this paper is to derive equilibrium constants for
the dehydration of water adducts (hydrates) of aromatic
molecules. Experimentally, dehydration occurs readily in the
presence of an acid catalyst to yield the aromatic product.1,3
The reaction is illustrated by the aromatization of benzene
hydrate (1) shown in Scheme 1 and is closely related to the
acid-catalyzed dehydration of alcohols to form alkenes.2 Rate
constants for dehydration have been measured for hydrates of
benzene, naphthalene, phenanthrene, and anthracene1 and also
benzofuran and benzothiophene.4 However, because the reaction
lies strongly in favor of the aromatic product, direct measure-
ments of equilibrium constants have not been possible.
The equilibrium constant for the dehydration reaction Kde may
be expressed as a difference in free energies of formation of
reactants and products, as shown in eq 1. Insofar as the free
energies of formation of water and benzene (and other aromatic
molecules) are known,5,6 evaluation of Kde reduces to measure-
ment or estimation of a free energy of formation of the aromatic
hydrate. Measurements of unknown free energies are not easily
undertaken, and the approach of this paper is to estimate ∆G°f
for the hydrates from existing thermodynamic data or measure-
ments of equilibrium constants of related reactions.
* Corresponding author. E-mail: rmof@ucd.ie.
(1) Rao, S. N.; More O’Ferrall, R. A.; Kelly, S. C.; Boyd, D. R.; Agarwal, R.
J. Am. Chem. Soc. 1993, 115, 5485.
(2) More O’Ferrall, R. A.; Rao, S. N. Croat. Chem. Acta 1992, 65, 593. 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.
-RT ln K ) ∆G°f(C6H6) +
∆G°f(H2O) - ∆G°f(hydrate) (1)
(3) Pirinccioglu, N.; Thibblin, A. J. Am. Chem. Soc. 1998, 120, 106.
(4) Kelly, S. C.; McDonnell, C. M.; More O’Ferrall, R. A.; Rao, S. N.; Boyd,
D. R.; Brannigan, I. N.; Sharma, N. D. Gazz. Chim. Ital. 1996, 126, 747.
The free energies of formation required refer to aqueous
solution at 25°. Hitherto, thermodynamic data have been more
9
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J. AM. CHEM. SOC. 2002, 124, 8561-8574
8561

A R T I C L E S
Dey et al.
readily available for the gas phase than solution. However, Hine7
and Guthrie5 have used vapor pressures and aqueous solubilities
of organic compounds to evaluate free energies of transfer from
gas to solution. By this means Guthrie has calculated a wide
range of free energies of formation and has developed a scheme
of additive group contributions,5 similar to those of Benson and
others for heats of formation in the gas phase,8 which allow
estimation of values of ∆G°f(aq) for compounds for which
measurements are lacking.
In principle, Guthrie’s group contributions may be used to
estimate values of ∆G°f(aq) for aromatic hydrates. However,
as emphasized by Benson,8 the generality of the group method
makes some sacrifice of precision. Thus, for prediction of
equilibrium constants, which are commonly measured with a
higher precision than thermodynamic data, a more precise
method is desirable.
This means that an unmeasured value of ∆G°f(aq) for an alcohol
may be estimated from the effect of replacing a hydrogen atom
of the hydrocarbon by a hydroxyl group. This replacement is
shown in eq 4, where R-H and R-OH represent the hydro-
carbon and alcohol and ∆GOH is a difference in their free
energies of formation.
∆GOH
R-H
8 R-OH
(4)
The unknown ∆G°f(aq) is derived by examining values of
∆GOH for alcohols and hydrocarbons for which ∆G°f(aq) is
known. Data are available from thermodynamic measurements
and from measurements of equilibrium constants for the
hydration of alkenes14-16 (if ∆G°f(aq) for the alkene is known).
The structure of the target alcohol is matched to that of alcohols
in this database and ∆GOH assigned accordingly.
The plan of this paper reflects this objective. First, equilibrium
constants for dehydration of simple alcohols and their depen-
dence on the structure of the alcohol and alkene are examined.
Then values of ∆G°f(aq) for alcohols, from conventional
thermodynamic measurements and from the hydration equilibria,
are converted to free energies for hydroxyl replacement [∆GOH
) ∆G°f(R-OH) - ∆G°f(R-H)], and their dependence on the
structure of R is evaluated. Finally, appropriate values of ∆GOH
are combined with measured or estimated values of ∆G°f(aq)
for dihydroaromatic molecules to obtain ∆G°f(aq) for the arene
hydrates. We can then derive equilibrium constants for dehydra-
tion from eq 1 provided ∆G°f(aq) for the corresponding aromatic
molecules are known.
For not too large molecules it is possible to calculate heats
of formation for the gas phase by molecular mechanics or by
semiempirical or ab initio quantum methods9-11 and to combine
these with measured, computed, or estimated entropies, vapor
pressures, and solubilities to obtain ∆G°f(aq). Such methods
have been applied to calculations of equilibrium constants for
the addition of water to carbon-oxygen10 and carbon-nitrogen11
double bonds, including hydration of polyazanaphthalenes (eq
2) and nucleotide bases.12
Results
In this paper, however, we adopt an approach used by Guthrie
for extrapolating the free energy of formation of an unknown
compound from that of a known related structure. Guthrie
estimated equilibrium constants for addition of water or alcohols
to the carbonyl groups of esters and amides to form tetrahedral
adducts by taking as models for the adducts corresponding ortho
esters or orthoamides in which a proacyl hydroxyl group had
been replaced by methoxyl and for which a free energy of
formation was available or could be measured. He evaluated
the effect of replacing the methoxyl by hydroxyl (e.g. as in eq
3) from differences in experimentally measured free energies
of formation of alcohols and methyl ethers and of acetals and
hemiacetals.13
Free Energies of Formation. Free energies of formation in
aqueous solution of a number of the alcohols and corresponding
hydrocarbons used in this paper have been compiled by Guthrie.5
They include straight chain primary alcohols and secondary and
tertiary alcohols containing one or two R-methyl substituents,
respectively. They also include secondary cycloalkanols, allyl
alcohol, benzyl alcohol, and vinyl alcohols (enols). Free energies
of formation of further alcohols, including acyclic and cyclic
allylic or benzylic alcohols, have been evaluated from experi-
mental measurements of equilibrium constants for dehydra-
tion14-17 combined with a knowledge a of ∆G°f(aq) for the
unsaturated product.
Free energies of formation of the products of dehydration,
and of the hydrocarbons from which ∆G°f(aq) for alcohols were
derived by hydroxyl substitution, were generally available for
the gas phase from the TRC compilation “Thermodynamics of
Organic Compounds in the Gas State”.6 This compilation
provides a source of experimental and calculated entropies which
allows conversion of heats of formation compiled by Pedley18
into free energies of formation for the gas phase and updates
the previously collected source of such data by Stull et al.19
Many of the values of interest here are for cycloalkyl rings for
RCH(OMe)2 + H2O h RCH(OMe)OH + MeOH (3)
To estimate free energies of formation of aromatic hydrates,
we have used a similar approach. A substantially larger body
of thermodynamic data exists for hydrocarbons than alcohols.
(5) Guthrie, J. P. Can. J. Chem. 1992, 70, 1042.
(6) Frenkel, M.; Kabo, G. J.; Marsh, K. N.; Rogonov, G. N.; Wilhoit, R. C.
Thermodynamics of Organic Compounds in the Gas State, Vols. I and II;
Thermodynamics Research Center: College Station, TX, 1994.
(7) Hine, J.; Mookerjee, P. J. J. Org. Chem. 1975, 40, 292.
(8) Benson, S. W. Thermochemical Kinetics; Wiley: New York, 1976. Benson,
S. W.; Buss, J. H. J. Chem. Phys. 1958, 29, 546. Cohen, N.; Benson, S.
W. Chem. ReV. 1993, 93, 2419.
(14) Schubert, W. M.; Keeffe, J. R. J. Am. Chem. Soc. 1972, 94, 559.
(15) Jensen, J. L.; Carre, D. J. J. Org. Chem. 1971, 36, 3180. Jensen, J. L.,
Uaprasert, V.; Fuji, C. R. J. Org. Chem. 1976, 41, 1675. Jensen, J. L.;
Uaprasert, V. J. Org. Chem. 1976, 41, 649.
(16) Gold, V.; Gruen, L. C. J. Chem. Soc. B 1966, 600.
(17) Chiang, Y.; Kresge, A. J.; Obraztsov, P. A.; Tobin, J. B. Croat. Chem.
Acta 1972, 65, 615, and references cited therein.
(18) Pedley, J. B. Thermochemical Data and Structures of Organic Compounds,
Vol. I; Thermodynamics Research Center: College Station, TX, 1994.
(19) Stull, D. R.; Westrum, E. F., Jr.; Sinke, G. C. The Chemical Thermodynam-
ics of Organic Compounds; Wiley: New York, 1967.
(9) Pople, J. A., Beveridge, D. L. Approximate Molecular Orbital Theory;
McGraw-Hill: New York, 1972.
(10) Wiberg, K. B.; Marquez, M. J. Am. Chem. Soc. 1998, 120, 2932.
(11) Wiberg, K. B.; Morgan, K. M.; Maltz, H. J. Am. Chem. Soc. 1994, 116,
11067. Wiberg, K. B.; Nakaji, D. Y.; Morgan, J. M. J. Am. Chem. Soc.
1993, 115, 3527.
(12) Erion, M. D.; Reddy, M. R. J. Am. Chem. Soc. 1998, 120, 3295.
(13) Guthrie, J. P. Can. J. Chem. 1975, 53, 898. Guthrie, J. P. J. Am. Chem.
Soc. 1978, 100, 5892.
9
8562 J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002

Dehydration of Adducts of Aromatic Double Bonds
A R T I C L E S
which values were critically reviewed and new calculations of
S° undertaken by Dorofeeva et al.20 In a few instances
experimental values of ∆H°f(g) for these compounds were not
available. Thus for cyclobutadiene ∆H°f(g) was a calculated
value.21 For 1,2-dihydronaphthalene ∆H°f(liq) was available22
and was combined with ∆Hv ) 13.6 kcal mol-1 estimated from
the bp (as described below) to give ∆H°f(g), which could be
combined with a calculated value20 of S° to obtain ∆G°f(g) )
55.1 kcal mol-1 from eq 5. For 1,4-dihydronaphthalene ∆G°f(g)
) 59.3 kcal mol-1 was similarly derived from21 ∆H°f(l) ) 20.1
kcal mol-1 using the same value for ∆Hv as that used for its
isomer; again S° was a calculated value.6
a direct measurement, by extrapolation of a plot of log p against
T, or from coefficients of the Antoine equation calculated from
the temperature dependence of the vapor pressure in a temper-
ature range removed from 25 °C.26 For cyclopropane and
butadiene, for which vapor pressure data are lacking, the
required value was derived from the (estimated) boiling point
and the heat of vaporization of the liquid (∆Hv) by making use
of eq 9.30 In this equation, p is the vapor pressure at 25 °C and
Tb is the boiling point of the liquid (K). In the absence of
satisfactory models for estimating ∆Cvap 31 for the hydrocarbons
concerned, the values were taken as -12.30
∆Hv
1
1
ln p ) ln 760 -
-
+
∆G°f(g) ) ∆H°f(liq) + ∆Hv - 298.16∆S°f(g)
(5)
{
}
R
298 Tb
For cyclobutanol, a value of ∆H°f(g) ) -34.6 kcal mol-1 is
available and S° can be estimated from the value for cyclobutane
plus an average of the increment in entropy for replacement of
H by OH in cyclopentane and cyclohexane. This led to ∆G°f(g)
∆Cvap
298 ln 298
1 -
-
(9)
{
}
R
Tb
Tb
Evaluation of the vapor pressure requires a knowledge of
∆Hv, which itself will normally be unavailable in the absence
of vapor pressure measurements as a function of temperature.
Fortunately, accurate estimates of ∆Hv can be obtained from
correlations of ∆Hv with bp32 or by group contributions,33 which
have been well-developed for cyclic hydrocarbons.33
) -10.4 kcal mol-1 23
.
For the unsaturated cyclic hydrocarbons, in most instances,
values of free energies of transfer ∆Gt were required to convert
gas-phase data into free energies of formation for aqueous
solution. These were obtained from the vapor pressure and
aqueous solubility at 25 °C following the procedure of Hine
and Guthrie summarized in eqs 6-8. In these equations γ is
Hine’s activity coefficient referring to transfer of a mole of gas
from an ideal 1 M concentration in the gas phase to aqueous
solution at 25 °C, p is the vapor pressure of the gas in
millimeters of mercury at 25 °C, cg is the corresponding molar
concentration of the gas, and cw is the (molar) solubility in water
(assuming that this is low enough for an ideal solution to
prevail).
Even where the bp of a compound is lacking, as for
cyclobutadiene, a value may be inferred from a correlation of
boiling points with the molecular connectivity parameter ø used
by Kier and Hall.34 White has used such a correlation to predict
bps for polycyclic aromatic hydrocarbons.35 For small and
medium ring hydrocarbons, the correlation is not linear and there
are separate correlation lines for cycloalkanes, cycloalkenes, and
cycloalkadienes, etc. However, these correlations are nearly
parallel, and it is possible to infer a bp ∼ -15 °C for
cyclobutadiene (if it is assumed that there is no difference in
bp between singlet and triplet structures).
Values of vapor pressure p and corresponding values of cg
for these compounds are listed in Table 1 together with values
of γ and ∆Gt obtained from their combination with cw according
to eqs 6-8. For a few bi- or tricyclic hydrocarbons, notably
1,2-dihydronaphthalene, indene, indane, 9,10-dihydronaphtha-
lene, and 9,10-dihydroanthracene, solubility measurements are
lacking. Our own attempts at spectrophotometric measurements
of the solubility of the lower molecular weight bicyclic
molecules were poorly reproducible, partly because the com-
pounds are liquids and less easily handled than solids. Although
dihydroanthracene and dihydrophenanthrene are solids, they
were too insoluble to allow straightforward measurements. In
these cases therefore values of log γ were again interpolated,
from comparisons within a matrix of similar structures, as shown
in Charts 1 and 2, in which values of log γ are displayed below
the structures and the interpolated values are shown in brackets.
log cg ) log p - 4.269
γ ) cw/cg
(6)
(7)
(8)
∆Gt ) -1.364{log γ - 1.39}
The data required therefore are vapor pressures and aqueous
solubilities at 25 °C.7 Solubilities of a number of cyclic and
acyclic hydrocarbons have been measured by McAuliffe,24 who
has shown that linear correlations exist between solubility and
molar volume for different ring sizes or with the degree of
unsaturation at constant ring size. Where the molar volume
(molecular weight/density) is not available, as for example for
cyclobutadiene, similar correlations of molar volume with ring
size or number of double bonds exist. These correlations allow
solubilities of all cyclic hydrocarbons of ring sizes from 3 to 8
to be inferred if measurements are lacking. Also inferred was
the solubility of trans-2-butene. Further solubility data have been
reported by Deno and Berkheimer.25
There are a number of compilations of vapor pressures26-29
from which vapor pressures at 25 °C may be obtained either as
(27) Smith, B. D.; Srivastava, R. Thermodynamic Data for Pure Compounds,
Part A, Hydrocarbons and Ketones; Elsevier: Amsterdam, 1986.
(28) Boublik, T.; Fried, V.; Hala, E. The Vapour Pressures of Pure Substances;
Elsevier: Amsterdam, 1984.
(29) Boyd, R. H.; Samuel, S. N.; Shary-Tehrany, S.; McNally, D. J. Phys. Chem.
1971, 75, 1264.
(20) Dorofeeva, O. V.; Gurvich, L. V.; Jorish, V. S. J. Phys. Chem. Ref. Data
1986, 15, 437.
(21) Rogers, D. W.; McCafferty, F. J.; Padison, A. V. J. Phys. Chem. 1996,
100, 7148. Glukotsev, N. N.; Laiter, S.; Pross, A. J. Phys. Chem. 1995,
99, 6828.
(22) Williams, R. B. J. Am. Chem. Soc. 1942, 64, 1395.
(23) Rocek, J.; Radkowsky, A. E. J. Am. Chem. Soc. 1974, 95, 7123.
(24) McAuliffe, C. J. Phys. Chem. 1966, 70, 1267.
(25) Deno, N. C.; Berkheimer, H. E. J. Chem. Eng. Data 1960, 5, 1.
(26) Stull, D. R. Ind. Eng. Chem. 1947, 39, 517.
(30) Benson, S. W.; Mendenhall, G. D. J. Am. Chem. Soc. 1976, 98, 2046.
(31) Guthrie, J. P.; Taylor, K. F. Can. J. Chem. 1984, 62, 363.
(32) Wadso, I. Acta Chem. Scand. 1966, 20, 544.
(33) Ducros, M.; Gruson, J. F.; Sannier, H. Thermochim. Acta 1980, 36, 39.
Ducros, M.; Sannier, H. Thermochim. Acta 1982, 54, 153.
(34) Kier, L. B.; Hall, L. H. Molecular ConnectiVity in Chemistry and Drug
Research; Academic Press: New York, 1976.
(35) White, S. M. J. Chem. Eng. Data 1986, 31, 198.
9
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A R T I C L E S
Dey et al.
Table 1. Solubilities, Free Energies of Transfer and ∆G°f(aq) (kcal mol-1) for Cyclic Hydrocarbons
log cw
pa
log cg
log γ
∆G
t
∆G°f(g)
∆G°f(aq)
cyclobutaneb,c
-2.20
-1.15
-1.60
-1.12
-1.62
-2.10
-2.06
-2.26j
-3.49
1175
-1.2
-1.00
-0.62
-0.56
-0.14
0.01
0.17
0.38
-2.26
1.41
3.26
2.74
2.66
2.09
1.88
1.66
1.38
3.09
-0.03
-0.08
0.19
-0.35
-0.49
-1.79
-2.33
2.52
26.80
68.50
48.33
103.3
42.55
43.5
44.60
15.13
40.07
56.2
40.04
55.10
59.1
69.15
70.76
34.90
30.06
71.2
51.0
105.4
44.43
45.2
cyclopropeneb,d
5320
1695
2030
437
-0.54
-1.04
-0.96
-1.63
-2.27
-2.44
-1.39
-0.03
cyclobuteneb,c
cyclobutadieneb,d,f,g
cyclopentadieneb,e
1,3-cyclohexadieneb,e
1,4-cyclohexadieneb,h
trans-2-butene
98
67.5
760k
0.233
46.0
18.22
40.04
56.1
40.2
54.75
58.6
67.36
68.43
37.42
1,2,3,4-tetrahydronaphthalene
indene
indane
1,2-dihydronaphthalene
1,4-dihydronaphthalene
9,10-dihydrophenanthrene
9,10-dihydroanthracene
trans-1,3-pentadiene
1.45p
1.25p
1.65p
1.75p
2.70p
3.10p
-0.46q
a Vapor pressure (mmHg) measurements at 25 °C. b Solubility from correlation of cw with molecular volume.24 c Vapor pressure from the Antoine
equation. d Vapor pressure at 25 °C from bp and a value of ∆Hv estimated from a correlation of ∆Hv and bp for cyclic and noncyclic hydrocarbons, or as
otherwise indicated. e The vapor pressure at 25 °C from Smith and Srivatsava.27 f Molecular volume predicted from correlation with ring size and degree
of unsaturation. g Bp from a plot of bp versus ø for cycloalkenes, alkenes, and alkadienes.35 h Solubility from McAuliffe.24 j Solubility estimated from
value for 1-butene multiplied by the ratio of values for trans-2-pentene and 1-pentene.24 k Gas solubility measured at 1 atm pressure. m Solubility from ref
25; vapor pressure from ref 29. p log γ interpolated as described in text. q log γ estimated as the value for 1,4-pentadiene corrected for the difference
between 1-pentene and 2-pentene (values from ref 5).
Chart 1
Chart 2
In Chart 1 the value for dihydronaphthalene comes from
allowance for the combined effects of (a) the degree of
unsaturation and (b) benzoannelation. Thus, the ratio of incre-
ments between compounds in the second row of the chart is
taken to be the same as that for the last three compounds of the
first.
In Chart 2 other values of log γ are interpolated from
consideration of the effects of benzoannelation and unsaturation
on five- and six-membered rings. These interpolations make
use of the value for dihydronaphthalene from Chart 1 and
recognize that the more compact structure of two isomeric
molecules (e.g. phenanthrene compared with anthracene) is
characterized by a smaller value of log γ.
More specifically the increments in log γ between adjacent
compounds in row two were assumed to be the same as in row
one (with interpolated values rounded to 0.05 log units). For
rows three and four it was assumed that the ratio of increments
between the first and the second and the second and the third
compounds was the same in the two rows. This ratio is similar
to but not quite the same as the corresponding ratio in row 5.
The remaining value to be assigned in Chart 2 is log γ for
indane which is taken as 1.25 by assuming that the difference
from indene is the same as the difference between 1,2-
dihydronaphthalene and 1,2,3,4-tetrahydronaphthalene. Despite
the arbitrary nature of these interpolations the error in log γ
should not be greater than 0.05 which amounts to a 10% change
in equilibrium constant.
with 2-5 carbons and cyclopentanol and cyclohexanol this
difference falls between 4.90 and 5.06: for cyclobutanol a value
of 4.9 was assigned, equal to that for the cyclic alcohols.
Experimental Measurements of Equilibrium Constants.
Direct measurements of equilibrium constants for dehydration
are often made only with difficulty because of the slow rate at
which equilibrium is achieved and the small amount of alkene
present at equilibrium with the dominant alcohol. For allylic
and benzylic alcohols, however, the rate of dehydration is faster
than for aliphatic alcohols, and there is a substantial fraction of
alkene formed at equilibrium. Moreover a strong chromophore
for the alkene allows the equilibrium constant to be conveniently
measured using UV-vis spectrophotometry. Equilibrium con-
stants for reactions relevant to this paper have been measured
by Schubert and Keeffe,14 Jensen and co-workers,15 Gold and
Gruen,16 and ourselves. Earlier measurements are cited in ref
17.
For cyclobutanol a value of log γ was inferred from inspection
of the difference in log γ between alcohols and the correspond-
ing hydrocarbons. For primary, secondary and tertiary alcohols
9
8564 J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002

Dehydration of Adducts of Aromatic Double Bonds
A R T I C L E S
These measurements are important in offering access to cyclic
allylic or benzylic alcohols, which are close analogues of arene
hydrates. No thermodynamic measurements are available for
such alcohols, but the combination of a measurement of an
equilibrium constant, for example for dehydration of cyclohex-
enol (2) to cyclohexadiene (3, eq 10), with a value of ∆G°f(aq)
for the diene product allows us to derive ∆G°f(aq) for the
alcohol, as illustrated in eq 11.
were extrapolated from measurements in concentrated acid
solutions based on plots of logs of second-order rate constants
against the medium acidity parameter X. The measurements
refer to 25° and gave 4.5 × 10-4 (cyclopentenol), 1.6 × 10-4
(1-phenyl-1-propanol), 5.3 × 10-6 (cyclohexenol), and 1.7 ×
10-5 (3-penten-2-ol) with units M-1 s-1. The measurements
could be made by taking either the alcohol or alkene as reactant,
and the rate constants represent a sum of values for the hydration
and dehydration reactions. In the case of 3-penten-2-ol the rate
constant (and equilibrium constant) measured refers to formation
of the trans-1,3-pentadiene (piperylene), as equilibrium is
established much more rapidly with the trans- than cis-
pentadiene product. An approximate rate constant for conversion
of cis-pentadiene to an equilibrium mixture of trans-3-penten-
2-ol and cis- and trans-pentadienes was measured as 3.0 × 10-6
∆G°f(C6H9OH)(aq) ) -RT ln K +
∆G°f(C6H8)(aq) + ∆G°f(H2O) (11)
M-1 s-1
.
1,2-Substituent Interactions. The main structural difference
between the alcohols for which equilibrium studies have been
carried out and the arene hydrates for which equilibrium
constants are to be estimated is illustrated by comparison of
benzene hydrate (cyclohexadienol) 1 and cyclohexenol 2. There
is an additional double bond in the hydrate not present in
cyclohexenol. If we consider the difference in free energy
changes accompanying introduction of a hydroxyl substituent
into cyclohexene and cyclohexadiene, respectively, it is apparent
that this corresponds to the energy of what is formally a 1,2-
interaction between the hydroxyl group and the remote (4, 5)
double bond of cyclohexadienol, as expressed by the isodesmic
reaction of eq 12. If the measurable free energy change between
cyclohexene and cyclohexenol is to be transformed into that
between cyclohexadiene and the hydrate, therefore account must
be taken of this interaction.
The alcohols for which equilibrium measurements were
carried out were 2-cyclohexenol (2), 2-cyclopentenol (4),
1-phenylpropan-1-ol (5), and trans-3-penten-2-ol (6); measure-
ments for tetralol (7) and indanol (8) were available from a
previous study.1 Values of ∆G°f(aq) for the products of
dehydration, namely, cyclohexadiene, cyclopentadiene, trans-
1,3-pentadiene, 1,2-dihydronaphthalene, and indene are listed
in Table 1.
There are experimental difficulties even with these measure-
ments, however. At 25 °C the acid-catalyzed approach to
equilibrium still occurs rather slowly. Equilibrium constants have
to be extrapolated from measurements in fairly concentrated
solutions of strong acids therefore, and, as the value appears to
be sensitive to the amount of acid in the medium, this
extrapolation is subject to some uncertainty. This difficulty is
compounded by the low solubility of the alkenes, their tendency
to polymerize, and in some cases their volatility or high
concentration relative to the alcohol at equilibrium. The equi-
librium constants are thus subject to a significant experimental
error (e.g. 20%). Nevertheless, the uncertainty is generally small
compared with the variations in magnitude of equilibrium
constants with changes of structures of the alkenes or alcohols
under examination.
The measurements were based on spectrophotometric moni-
toring of the reactions and are described in the Experimental
Section of the paper. By combining the equilibrium constants
found with values of ∆G°f(aq) for the appropriate dienes or
styrenes the following values of ∆G°f(aq) were calculated for
the alcohols (kcal mol-1) using eq 11: 2-cyclohexenol (2),
-12.85; 2-cyclopentenol (4), -11.2; tetralol (7), -0.1; indanol
(8), 0.8; 1-phenylpropan-1-ol (5), -7.1; trans-3-penten-2-ol (6),
-20.3. In addition, values of ∆G°f(aq) were derived for
1-penten-3-ol (-17.8) and 1-buten-3-ol (-19.8), by assuming
that the substituent effect of replacing H by OH in 2-pentene
to give 3-penten-2-ol was the same as the corresponding
replacement in 1-pentene and 1-butene to give 1-penten-3-ol
and 1-buten-3-ol, respectively.
It appears that no thermodynamic data or equilibrium
measurements exist for 3,4-unsaturated alcohols, i.e., alchohols
which include the structural feature of interest. The magnitude
of the interaction must be inferred therefore by considering other
1,2-interactions. In general, for substituents X and Y this may
be expressed by the isodesmic equilibrium of eq 13. If the
equilibrium constant for the reaction is unity (∆G° ) 0 kcal
mol-1), there is no net stabilization or destabilization. If the
equilibrium constant is <1 (∆G° positive), the interaction is
stabilizing, and if the equilibrium constant is >1 (∆G° negative),
the interaction is destabilizing.
Available measurements based on eq 13 are listed in Table
2. The table includes values of ∆H° based on data for the gas
phase18,19 in addition to values of ∆G° for the gas phase6,19 and
aqueous solution.5 For X ) Y ) Cl it is found that ∆H°(g) )
-3.2 kcal mol-1, ∆G°(g) ) -3.4 kcal mol-1, and ∆G°(aq) )
-4.6 kcal mol-1; i.e., the 1,2-interaction of two chlorine atoms
is quite strongly repulsive. The same is true to a lesser degree
of two bromine atoms or two iodine atoms. On the other hand,
a chlorine atom and a methyl group are attractive (∆G°(aq) )
Rate constants were also measured for the acid-catalyzed
hydration or dehydration reactions. Values for aqueous solution
9
J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002 8565

A R T I C L E S
Dey et al.
Table 2. 1,2-Interaction Energies (kcal mol-1) between X and Y
Table 3. Equilibrium Constants (Kde) for Dehydration of Aliphatic
Alcohols in Aqueous Solution at 25 °C
Based on Equation 13a
X
Y
∆H° (g)
∆G°(g)
∆G°(aq)
CN
Cl
Br
F
CN
Cl
Br
F
-5.9
-3.2
-0.6
0.6
-3.4
-2.4
0.2
-4.6
-3.5
NH2
HO
CH2dCH
Ph
NH2
HO
CH2dCH
Ph
1.6
1.0
0.2
0.2
-2.7
-0.75
-1.0
-0.65
CH3
CN
Cl
Cl
F
NH2
HO
CH2dCH
Ph
CH3
CH3
CH3
CH2Cl
CH3
CH3
CH3
CH3
CH3
0
0.6
-0.7
4.3
1.0
-0.5
0
-0.7
2.0
-0.2
-0.4
0.4
-0.2
0.4
0.7
3.8
0.1
-0.1
-0.25
1.2
0.05
-0.1
a Data from refs 5, 6, and 18: positive values indicate favorable
interactions.
3.8 kcal mol-1). In general, the interactions seem to show a
dependence on the electronegativity (and possibly polarizability)
of X and Y.
a Kde is the equilibrium constant for dehydration: [alkene]/[alcohol].
c
b pKde is the negative log of Kde. KH O is the equilibrium constant for
2
hydration; values are calculated from values of ∆G°f(aq) (especially from
The behavior of two OH (or NH2) groups is a little different
from that of the halogens insofar as ∆H° in the gas phase is
positive; i.e., the interaction is attractive. This is presumably
because HOCH2CH2OH and NH2CH2CH2NH2 are stabilized by
hydrogen bonding. For two ether oxygen atoms, for example
in dioxan, for which no hydrogen bonding is possible, the
interaction is again repulsive. Thus, for the equilibrium of eq
ref 5). d Measured experimentally.16
group (as it is when a 3,4 double bond is introduced into a
cycloalkanol, e.g. in transforming 2 to 1). It is difficult to judge
the magnitude of this interaction, but it seems likely that it is
less than that between two oxygen atoms. We have chosen -1.0
kcal mol-1 for both vinyl and phenyl groups, therefore. Although
this is an arbitrary choice, it seems clear that the free energy of
interaction between hydroxyl and vinyl or phenyl must be small
and negative so the uncertainty in the assignment should also
be small.
14, ∆H° ) -2 kcal mol-1
.
On the other hand ∆G° for HOCH2CH2OH, in the gas phase
is -2.7 kcal mol-1, despite the positive value of ∆H°.
Presumably this reflects hindering of rotation about the C-C
bond by the hydrogen bonding. ∆G° remains negative in
aqueous solution (∆G°(aq) ) -1.0 kcal mol-1), and in the
absence of intramolecular hydrogen bonding apparently the
mutual repulsion of the electronegative oxygen atoms reasserts
itself.
Discussion
Dehydration of Aliphatic Alcohols. Equilibrium constants
for dehydration (Kde) of a series of aliphatic alcohols to form
alkenes are shown in Table 3. The equilibrium constants refer
to the reaction in the dehydration direction, although it is often
useful to consider values for the reverse hydration (KH2O). The
constants are expressed as pKs, i.e., the negative logarithms of
the dehydration constants (which correspond to log KH O for
2
However, the repulsion between two oxygen atoms is
noticeably less than between two chlorine atoms. Similarly, the
hydration). To emphasize the direction to which Kde and pKde
refer, the normal double arrows for the equilibria are replaced
by single arrows. It should be noted that because water is the
solvent, its concentration is not included in the equilibrium and
interaction between OH and CH3 (∆G°(aq) ) 1.2 kcal mol-1
)
is less favorable than between Cl and CH3 (∆G°(aq) ) 3.8 kcal
mol-1). In this connection, it is noteworthy that the fluorine-
fluorine interaction in FCH2CH2F is positive rather than negative
(∆G°(g) ) +0.6 kcal mol-1) and the interaction between F and
CH3 is negative rather than positive (∆G°(g) ) -0.4 kcal
mol-1). Thus, the oxygen atoms appear to be intermediate in
behavior between Cl and F.
Kde and KH O are expressed as in eqs 15 and 16.
2
Kde ) [alkene]/[alcohol]
(15)
(16)
KH O ) 1/Kde
2
Our interest is in the interaction between OH and a double
bond or phenyl group. Two vinyl groups (in 1,5-hexadiene) are
mildly repulsive in aqueous solution (∆G°(aq) ) -0.65 kcal
mol-1), and if this is an electronegativity effect, it seems likely
that the same is true of vinyl and OH, especially if this
interaction is compared with the interaction of OH and an alkyl
Examination of the structural dependence of the equilibrium
constants shows that, as is well-known, for simple aliphatic
alcohols the alcohol is favored over the alkene, but that the
equilibrium is shifted to the right by methyl substitution of the
double bond. Thus, the equilibrium constants for dehydration
9
8566 J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002

Dehydration of Adducts of Aromatic Double Bonds
A R T I C L E S
Scheme 2
Scheme 3
is practically no influence of R-vinyl or R-phenyl substitution
upon the stability of the alcohol.
of ethanol, n-propanol, and isobutanol to give ethene, propene,
and isobutene are 1.6 × 10-5, 0.011, and 0.83, respectively.
The effect of constraining double bond formation within a
five- or six-membered ring is shown in Scheme 3. It can be
seen that reaction is more favorable within the ring than in the
open-chain alcohol by a factor of 6 for cyclopentanol and 30
for cyclohexanol.
It may seem surprising that dehydration occurs more easily
in the six- than in the five-membered ring, but this reflects not
the respective stabilities of the double bonds but the relative
stabilities of the hydroxyl groups in the five- and six-membered
rings. Thus introduction of a double bond into cyclohexane is
energetically less favorable than into cyclopentane, but this is
more than compensated by the corresponding difference for the
hydroxyl group. The relative stabilities of the open-chain
compound and five- and six-membered rings containing a double
bond compared with the corresponding saturated hydrocarbons
are 1:6:2.5.
Perhaps less well appreciated is that a methyl substituent on
the double bond of the alkene which is R rather than â to the
hydroxyl group of the alcohol does not have the same effect
and that the equilibrium constant for dehydration of 2-propanol
is hardly more favorable than that for ethanol (5.9 × 10-5
compared with 1.6 × 10-5). This is a manifestation of the
stabilizing effect of a “geminal” interaction between the methyl
group and hydroxyl group in the alcohol reactant. The magnitude
of this interaction is expressed by ∆G°(aq) ) 5.5 kcal mol-1
and equilibrium constant K ) 10-4 for the “bond separation”
reaction of eq 17. It arises from the favorable interaction between
carbon-oxygen and carbon-carbon σ-bonds and has been
discussed in detail in the literature, especially for the stronger
mutual interaction of carbon-oxygen bonds in acetals which
forms the basis of the anomeric effect.36,37
However, of principal interest in this paper are equilibrium
constants for dehydration of unsaturated and benzoannelated
cyclic alcohols rather than simple alicycles. When considering
the combined effect of double bond and ring structure, therefore,
it should be noted that Scheme 2 shows the effect of a double
bond (or phenyl group) replacing a hydrogen atom in ethanol
as reactant and ethylene as product. In Scheme 3 by comparison
ethanol has been replaced by 3-pentanol (11) as the reference
open-chain alcohol, in recognition of the fact that the cyclic
alcohols are secondary alcohols yielding alkenes with alkyl
substituents at each carbon atom of the double bond. Introduc-
tion of a double bond into the alcohols in Scheme 3 therefore
means that it replaces an alkyl group rather than a hydrogen
atom in both the alcohol and alkene, as may be seen by
comparing 11 (in Scheme 3) with 1-penten-3-ol (12) in eq 18
(cf. also footnote c in Table 5).
CH3CH2OH + CH4 h CH3OH + CH3CH3
(17)
It is worth noting that the equilibria are also influenced by
statistical factors. Thus, dehydration of 2-propanol is favored
by a factor of 6 corresponding to the six equivalent hydrogen
atoms in the alcohol. There are three equivalent hydrogens in
ethanol, but the influence of these is compensated by a factor
of 2 in the hydration direction arising from the equivalence of
the two carbon atoms in the ethene product. In the absence of
statistical factors there would be practically no difference in
equilibrium constant for the two alcohols. Coincidentally, when
corrected for the presence of nine equivalent hydrogen atoms
in the alcohol,17 the equilibrium constant for dehydration of tert-
butyl alcohol also becomes the same as that for ethanol and
2-propanol.
Allylic, Benzylic, and Cyclic Alcohols. If the range of
equilibrium constants is extended to allylic and benzylic
alcohols, we find, again as expected, that the alkene product is
stabilized and its concentration at equilibrium is increased. This
is illustrated by the three reactions of Scheme 2, which lead to
formation of ethylene, butadiene, and styrene, respectively. It
can be seen that the vinyl and phenyl substituents in 3-buten-
2-ol (9) and R-phenylethanol (10) favor dehydration by factors
of 4400 and 1500, respectively. This arises from the expected
stabilization of the alkene. In contrast to alkyl substitution, there
In practice, both the open-chain alkene and alcohol are
stabilized by alkyl substitution. This means that replacing an
alkyl group rather than hydrogen by a double bond or phenyl
group is less stabilizing for both the alkene and the alcohol.
The effects on the equilibrium constant for dehydration are
compensating, therefore, and in practice the double bond in
1-penten-3-ol increases the equilibrium constant for dehydration
(eq 18) by 5450-fold which differs little from the 4400-fold
difference between 9 and ethanol in Scheme 2.
(36) Chang, Y.-P.; Su, T.-M. J. Phys Chem. A 1999, 103, 8706, and references
cited therein.
(37) Harcourt, M. P.; More O’Ferrall, R. A. Bull. Soc. Chim. Fr. 1988, 407;
Richard, J. P.; Aymes, T. L.; Rice, D. J. J. Am. Chem. Soc. 1993, 115,
2523.
We are now in a position to examine the effect of ring
formation upon dehydration of saturated, unsaturated, and phenyl
9
J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002 8567

A R T I C L E S
Dey et al.
Table 5. Relative Equilibrium Constants (Kde) for Dehydration of
Open-Chain and Cyclic Saturated, Unsaturated, and
Benzoannelated Alcohols
Table 4. Equilibrium Constants (pKde) for Dehydration of Allylic,
Benzylic, and Cyclic Alcohols in Aqueous Solution at 25 °C
a Reaction of open-chain alcohol to form trans-alkene. b Equilbrium
constant, 3-pentanol f 2-pentene, Kde ) 1.1 × 10-3 (Scheme 3). c 1-Penten-
3-ol f 1,3-pentadiene. Note that replacement of 1-penten-3-ol by 2-hexen-
4-ol increases this factor of 5450 to 20 000 (if the influence of the extra
methyl group is modeled by the difference in the effect of a methyl group
on the double bond of propene and butadiene). The larger value should
provide a better comparison with the cyclic alcohols. d 3-Cyclopentenol f
cyclopentadiene, Kde ) 3.5. e Indanol f indene, Kde ) 10.0. f 3-Cyclohex-
enol f 1,3-cyclohexadiene, K ) 0.1. g Tetralol f 1,2-dihydronaphthalene,
Kde ) 22.
reaction of cyclohexenol (2) is 60 times less favorable than that
of its open-chain analogue 1-penten-3-ol (12) in eq 18.
This difference is presumably a result of ring strain in the
cycloalkadiene. However, as can be seen by comparing columns
one and three in Table 5, surprisingly, the influence of ring
formation in benzoannelated cycloalkenols reverts to the pattern
of the saturated alcohols, with ring formation weakly favoring
dehydration! This effect of benzoannelation is reflected in a
difference in equilibrium constants for isodesmic reactions
comparing the effect of replacing a double bond by a benzene
ring in open-chain reactions (eq 19) and cyclic reactions (eq
20), which amounts to a factor of nearly 600-fold.
a Equilibrium constant for dehydration. b pKde is the negative log of Kde.
c KH O is the equilibrium constant for hydration. d Calculated from ∆G°f
(aq)2for alcohol and alkene. e ∆G°f(aq) for alcohol estimated by assuming
∆GOH (eq 21) is the same as that for 3-penten-2-ol. f This work. g See also
ref 15. h Reference 14.
The surprisingly large value of the equilibrium constant for
eq 20 was previously interpreted as partly reflecting a significant
difference in geminal interactions between the hydroxyl group
and a double bond and a hydroxyl group and a phenyl ring.1
As discussed below, it now seems likely that the contribution
from this source amounts only to a factor of 3 and that the main
contribution to the difference is a greater stabilization of a
conjugated annular double bond by a phenyl ring than a second
double bond.
substituted alcohols. Absolute and relative equilibrium constants
for dehydration of open-chain and five- and six-membered ring
alcohols are shown in Table 4 and Table 5, respectively. In
Table 5 the unsaturated cyclic alcohols are 4 and 2, and their
phenyl substituted counterparts are the benzoannelated deriva-
tives, 8 and 7. The equilibrium constants in the table are
measured relative to the dehydration of 11 (the second reaction
in Table 4).
It is remarkable that contrary behavior is seen for the open-
chain equilibrium (eq 19). In this case the conjugated π-bond
is more stabilizing than the phenyl group. It seems best to ascribe
this difference to a greater steric hindrance to achieving a planar
configuration favoring conjugation between a double bond and
a phenyl ring than between two double bonds in the open-chain
structures.
For the cyclic as for the acyclic alcohols the presence of the
double bond increases the extent of dehydration. This can be
seen in Table 5 by reading from left to right. On the other hand,
reading down the table it can be seen (by comparing the first
two columns) that the transformation from open-chain to cyclic
structure has a different effect upon saturated and unsaturated
alcohols. Thus, while reaction within a ring favors dehydration
of saturated alcohols, the corresponding reaction of cycloalk-
enols to form cycloalkadienes is strongly disfavored. Thus,
In the five-membered ring, benzoannelation displays a smaller
advantage over the π-bond (i.e. in comparing indanol and
9
8568 J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002

Dehydration of Adducts of Aromatic Double Bonds
A R T I C L E S
cyclopentenol), and the behavior here is intermediate between
that of the open-chain and six-membered ring structures.
Hydroxyl and π-Bond Replacements. This detailed discus-
sion of dehydration equilibria introduces an attempt to estimate
the effect of exchanging a hydroxyl group for a hydrogen atom
upon the free energy of formation of different hydrocarbons.
This effect may be denoted ∆GOH and is expressed by eq 21.
Our objective is to evaluate (and then predict) the influence of
alkyl and aryl substituents R upon this free energy change.
∆GOH ) ∆G°f(R-OH) - ∆G°f(R-H)
(21)
In practice, it is more convenient to consider the effect of a
change in the alkyl group than absolute values of ∆GOH
.
Adopting Grunwald’s notation38 this may be represented as
δR∆GOH and corresponds to a substituent effect on the process
of hydroxyl substitution. It is also convenient to choose a
reference structure (or substituent) with which different R groups
can be compared. If the reference is taken as a hydrogen atom
(R ) H), δR∆GOH corresponds to the reaction R-H + H2O f
R-OH + H2 and the analysis of structural changes achieves
most of the elements of a free energy relationship.
Figure 1. Values of δOH (kcal mol-1) (eqs 22 and 24) for replacement of
H by OH in an alkane (relative to ethane and ethanol, ∆GOH ) -39.5 kcal
mol-1).
In practice, use of H as a reference is less useful here than in
other free energy relationships. This is because the reaction then
includes a bond between two hydrogen atoms, whereas we will
be concerned solely with C-H (and C-OH) bonds. A methyl
group might seem a suitable alternative choice, but in the context
of dehydration reactions ethyl is to be preferred, because ethyl
alcohol is the simplest alcohol that can undergo dehydration.
As shown in eq 22, with ethyl as reference δR∆GOH corresponds
to the free energy change for the reaction shown in eq 23 with
equilibrium constant KS. At the risk of abusing Grunwald’s
notation,38 δR∆GOH in eq 22 is abbreviated as δOH (or δROH).
The relationship between δOH and ∆GOH is then given by eq
24 in which -39.5 (kcal mol-1) is the value of ∆GOH for ethyl
alcohol.
Although alkenes are not the principal concern of this paper,
eq 27 in combination with eq 24 provides a convenient
framework for factoring structural effects upon dehydration
equilibria into contributions from the alcohol reactant and alkene
product, as was shown above to be necessary for interpreting
substituent effects on these equilibria. The difference in sub-
stituent effects upon introduction of a hydroxyl group and a
double bond (δOH - δπ) corresponds to the difference in (the
log of) the equilibrium constants for dehydration of ROH and
ethanol. For this reason, and to provide a comparison with the
treatment of alcohols, a brief summary of alkyl and aryl
substituent effects upon the stabilities of double bonds as well
as of alcohols is presented.
At the outset, it should be recognized that these substituent
effects cannot be represented in the normal graphical form of a
free energy relationship. Although δπ and δOH correspond to
clearly defined reactions (eqs 23 and 26) and substituent effects,
no reference reaction has been defined. In principle the process
corresponding to δOH (e.g.) could be defined as a reference
“reaction” and then used to correlate substitution of hydrogen
by other groups or atoms such as chlorine (δCl). However, for
unknown values of δOH (or δπ) themselves we rely on inspection
of structural analogies.
Representative substituent effects δOH and δπ are shown in
Figures 1 and 2, respectively. Considering first the values for
alcohols, it can be seen in the top row of Figure 1 that methanol
has a positive value of δOH ) 5.5 kcal mol-1. This implies that
the free energy change for the reaction shown in eq 28 is
unfavorable by this amount. The positive value of δOH is a
consequence of the geminal interaction between methyl and
hydroxyl groups attached to the same carbon atoms present in
ethanol but not in methanol, as noted above. Moving from left
to right in the first row of the figure, and to 2-propanol in the
second row, it can be seen that extension of the alkyl chain
length in a primary alcohol has little influence on this geminal
effect, but that secondary (second row) and tertiary alcohols
δOH ) δR∆GOH ) -RT ln KS
(22)
KS
R-H + CH3CH2OH y\z R-OH + CH3CH3 (23)
δOH ) ∆GOH + 39.5
(24)
An advantage of choosing the ethyl group as reference is that
it can be used not only for introducing an OH group into a
hydrocarbon but also for introducing a double bond. The
influence of alkyl structure upon the stability of a π-bond can
then be considered analogously to that of a hydroxyl group.
This is shown by defining δπ in analogy with δOH as the free
energy for replacing two hydrogen atoms on adjacent carbon
atoms by a π-bond, as is summarized in eqs 25-27. In eq 27,
the free energy change 23.4 (kcal mol-1) is ∆Gπ for ethylene,
i.e., the counterpart of ∆GOH for ethanol (-39.5 kcal mol-1) in
eq 24.
δπ ) -RT ln KS
(25)
KS
R-CH2CH3 + CH2dCH2 y\z
R-CHdCH2 + CH3CH3 (26)
(38) Leffler, J. E.; Grunwald, E. Rates and Equilibria of Organic Reactions,
δπ ) ∆Gπ - 23.4
(27)
McGraw-Hill: New York, 1963.
9
J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002 8569

A R T I C L E S
Dey et al.
the more subtle effects of cyclic structures are shown in rows
one and two. Consistent with the discussion of equilibria, the
opposite effects of passing from an open-chain structure to five-
or six-membered rings upon the stabilizing effect of conjugation
of two double bonds on the one hand and upon conjugation of
a double bond and a phenyl group on the other are evident from
comparing rows three and four.
The general conclusion to be drawn from Figures 1 and 2 is
that although there is considerable variation in the effects of
incorporating a hydroxyl group or double bond into a hydro-
carbon molecule, these variations can be interpreted in a fairly
straightforward manner. For the alcohols, which are of principal
concern in this study, the effects arise from geminal interactions
of alkyl, vinyl, and phenyl groups and a smaller influence of
cyclization. The results are sufficiently systematic to give some
confidence in the prediction of ∆G°f (aq) for alcohols which
are not structurally too far removed from the examples of Figure
1.
Figure 2. Values of δπ (kcal mol-1) for replacement of two hydrogens by
a double bond in an alkane (relative to ethane and ethylene, ∆Gπ ) 23.4
kcal mol-1).
introduce additional, although attenuated, stabilizing interactions,
as is apparent from the magnitude and sign of their δOH values
(e.g. -2.9 kcal mol-1 for 2-propanol). In row two, the smaller
differences of open-chain from cyclic alcohols can also be seen,
although it should be recognized that cyclopentanol and
cyclohexanol are favored over their open-chain counterparts by
statistical factors of 5/2 and 6/2, respectively.
Free Energies of Formation of Aromatic Hydrates. If we
turn now to aromatic hydrates, such as benzene hydrate
(cyclohexadienol) 1 and naphthalene hydrate 14, it is clear that
their free energies of formation can be estimated if ∆G°f(aq)
for the corresponding hydrocarbon is known and the value of
δ
OH (or ∆GOH) for replacement of hydrogen by a hydroxyl group
can be estimated, on the basis of a satisfactory structural model.
For benzene hydrate and naphthalene hydrate the relevant
hydrocarbons are 1,3-cyclohexadiene 3 and 1,2-dihydronaph-
thalene 13. Evaluation of ∆G°f (aq) for these structures has been
described above, and the transformation to their hydrates is
shown in eqs 29 and 30. Inspection of Figure 1 suggests that
the closest models for estimating ∆GOH for these transformations
are provided by the conversion of cyclohexene to 3-cyclohexenol
(2) and 1,2,3,4-tetrahydronaphthalene to tetralol (7).
∆G°(aq)
CH3CH2OH + CH4 y
z CH3OH + CH2CH3 (28)
So far, the values of δOH considered have been based on
values of ∆G°f (aq) for alcohols and alkanes from Guthrie’s
compilation of free energies of formation for organic molecules.5
The same is true for the allyl and benzyl alcohols in row three:
these alcohols show positive values of δOH reflecting the less
favorable geminal interaction of a hydroxyl group with a vinyl
or phenyl substituent than an alkyl group. Row four shows
values of δOH for cyclic allylic and benzylic alcohols for which
values of ∆G°f(aq) are not available in the literature and which
have been derived, as described above, by combining measure-
ments of equilibrium constants for dehydration with known or
estimated values of ∆G°f (aq) for the alkene products.
From comparison of rows two and four it can be seen that,
as for the open-chain alcohols, a double bond or phenyl group
R to the hydroxyl group reduces δOH relative to the cycloalkanol.
However, there are significant deviations from additivity of the
effects of cyclization and incorporation of double bonds or
phenyl groups for five- and six-membered ring alcohols. These
were considered in detail in the discussion of dehydration
equilibria, and Figure 1 summarizes the contribution to the
equilibria from changes in the stability of the alcohol: thus
hydroxyl substitution at an allylic carbon is consistently more
favorable than at a benzylic carbon atom by between 0.4 and
1.3 kcal mol-1. Finally, in rows five and six the corresponding
influence of structure upon δOH for enolic (and phenolic)
hydroxyl groups is shown. Apparently the stabilizing effect of
an R-(geminal) substituent on the double bond is much smaller
than for alphatic alcohols. One and especially two â-alkyl groups
are destabilizing, probably because of a steric effect.
The single structural difference of the models from the desired
hydrocarbons and alcohols is the presence of an extra double
bond with a 1,2-relationship to the hydroxyl group. Unfortu-
nately, as noted above, no thermodynamic measurements appear
to have been made for molecules possessing this structural
feature. A correction for the interaction of a hydroxyl group
and double bond has been crudely estimated therefore on the
basis of inspection of available measured 1,2-interactions of
substituents (Table 2). As already described, although interac-
tions between electronegative atoms can introduce a correction
as large as 3-4 kcal mol-1, the best approximation for hydroxyl
and vinyl (or phenyl) substituents at present appears to be
approximately 1 kcal mol-1. Thus, ∆G°f(aq) for benzene and
naphthalene hydrates may be based directly on values of ∆GOH
for cyclohexenol and tetralol, respectively, with a correction of
+1.0 kcal mol-1. The only further correction to be made is a
statistical factor of 2 arising from the fact that there are four
equivalent H atoms that may be substituted by hydroxyl in
Figure 2 likewise summarizes the contribution of changes in
the stability of the alkene to structural effects on equilibrium
constants for dehydration. The stabilizing effect of alkyl
substitution and of conjugation upon a double bond as well as
9
8570 J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002

Dehydration of Adducts of Aromatic Double Bonds
A R T I C L E S
of 0.3 and 1.5 kcal mol-1 for their additional double bond or
phenyl ring, respectively. These are also included in Figure 3.
Before values of ∆G°f(aq) in Figure 3 are used to esti-
mate equilibrium constants, it is of interest to derive values of
∆G°f(aq) for a further group of alcohols which on dehydration
yield the antiaromatic product cyclobutadiene or the nonaromatic
but highly strained small-ring alkenes cyclopropene and cy-
clobutene. Estimates of δOH for these molecules are cruder than
for the aromatic hydrates. However, the gas-phase free energies
of formation of saturated and unsaturated cycloalkanes, from
cyclopropane to cyclooctane, have been compiled and critically
evaluated by Dorofeeva et al.20 These provide data both for the
products of dehydration and for the hydrocarbons from which
∆G°f(aq) for the alcohol can be derived, provided that free
energies of transfer from the gas phase to solution and values
of δOH have been measured or can be estimated. In the case of
cyclobutadiene the value of ∆H°f for the gas phase is based on
a mean of recently calculated values.21
As described above, evaluation of ∆Gt can be satisfactorily
carried out by estimating solubility data and vapor pressures,
the latter based if necessary on bps of liquid hydrocarbons. Even
for cyclobutadiene, a bp can be predicted from a correlation of
bps of cyclic hydrocarbons with the molecular connectivity
parameter ø 35 (although this does not take account of the fact
that the ground state of cyclobutadiene is a triplet rather than a
singlet, i.e., that correlation of the bp with singlet cycloalka-
dienes may be oversimplified).
For cyclobutanol ∆G°f(aq) was based on an experimental
value of ∆H°f(g) and interpolated values of S°(g) ) 75.8 cal
deg-1 mol-1 and log γ. More speculatively δOH ) -4.7 kcal
mol-1 for cyclopranol was assigned by assuming a smooth
variation of statistically corrected values of δOH for cyclic
alcohols upon ring size, with the value for acetaldehyde enol
taken as that for a two-membered ring. This may be compared
with an STO31G calculation which showed that ∆H for
replacement of H by OH in cyclopropane in the gas phase to
be 2.8 kcal mol-1 more favorable than for the 2-position of
propane.39 Taking δOH ) -2.9 kcal mol-1 for cyclopropanol
from Figure 1 our value for this difference (free energy in
aqueous solution) is 1.8 kcal mol-1. For 3-hydroxycyclobutene
Figure 3. Free energies of formation ∆G°f(aq) for cycloalkanols and arene
hydrates. Values in brackets are based on estimated values of ∆Gt (free
energy of transfer from gas to aqueous solution for the parent hydrocarbon).
tetrahydronaphthalene and only two in dihydronaphthalene. The
resulting values of ∆G°f(aq) are shown in Figure 3.
Other aromatic hydrates for which it was possible to estimate
∆G°f(aq) in a similar manner include the isomeric hydrates of
benzene and naphthalene, 15-17 and the 9,10-water adducts
of phenanthrene 18 and anthracene 19. For the naphthalene
hydrate 16 a model for ∆GOH is now provided by cyclohexenol
corrected for a statistical factor of 2 and a 1,2-interaction
between a hydroxyl group and phenyl group, again taken as
+1.0 kcal mol-1. A similar estimate can be made for the
9,10-hydrate of phenanthrene using the value of ∆GOH for
tetralol together with a correction for an OH/phenyl 1,2-
interaction again equal to +1.0 kcal mol-1. The derived value
of ∆G°f(aq) and that for 15 are also shown in Figure 3.
A different estimate is required for the 1,4-hydrates of
benzene 15 and naphthalene 17 and the 9,10-hydrate of
anthracene 19. These molecules have two phenyl groups or
double bonds, or a phenyl group and a double bond, R to the
position of hydroxyl substitution. The hydroxyl group is subject
to geminal interactions with the phenyl or vinyl groups, and
these interactions are modified by incorporation in the cyclo-
hexyl ring. Insofar as geminal effects are not additive, and
ineraction of OH with the saturated cycloalkane is more
favorable than with annular vinyl or phenyl groups, we have
assumed that the second alkyl group (in cyclohexanol) has 60%
of the effect of the first (in cyclohexenol or 1-hydroxy-1,4-
dihydronaphthalene). This is quite a crude assumption but is
based on the observation that accumulation of other geminal
substituents including alkyl leads to a fall off in their effects
relative to hydrogen and probably also to vinyl and phenyl. On
the basis of Figure 1 the presence of a double bond and a phenyl
δ
OH was taken as -1.5 kcal mol-1 on the basis of δOH ) -4.4
kcal mol-1 for cyclobutanol and the effect upon δOH of
introducing a 2,3-double into cyclopentanol and cyclohexanol
(Figure 1).
The free energies of formation estimated for these alcohols
are included in Figure 3. To obtain equilibrium constants for
the dehydration reactions, it remains only to combine these with
the known or already estimated values of ∆G°f(aq) for the
unsaturated reaction products.
Dehydration of Aromatic, Antiaromatic, and Ring-
Strained Hydrates or Alcohols. Negative logs of equilibrium
constants (pKde) for the dehydration reactions of alcohols for
which ∆G°f(aq) has been estimated are shown in Table 6. For
comparison the dehydration equilibria for ethanol, 2-propanol,
cyclohexenol, and vinyl alcohol (ethenol) are included in the
table.
ring in cyclohexanol increases ∆GOH by 0.2 and 0.8 kcal mol-1
,
respectively (allowing for the fact that there are 12 equivalent
hydrogens in cyclohexane compared with four in cyclohexene
or 1,2,3,4-tetrahydronaphthalene). Values of ∆G°f(aq) for the
1,4-hydrates of benzene and naphthalene and 9,10-hydrate of
anthracene are based on values of ∆GOH, reflecting an increase
The dehydration of vinyl alcohol to form acetylene is also
included for comparison with the dehydration of the three-, four-,
(39) Clark, T.; Spitznagel, G. W.; Klose, R.; Schleyer, P. v. R. J. Am. Chem.
Soc. 1984, 106, 4412.
9
J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002 8571

A R T I C L E S
Dey et al.
Table 6. Equilibrium Constants for Dehydration of Alcohols
Yielding Aromatic, Antiaromatic, or Strained Cyclic Products
ring. A fanciful interpretation of this is that the extra two
π-electrons of the unsaturated two-membered ring render it
aromatic!
The elimination within the four-membered ring is not very
much less favorable than within structurally similar open-chain
molecules. The relatively high stability of a double bond in a
four-membered ring is also evident from heats of hydrogenation
and has been attributed to compensation between an increase
in angle strain and decrease in torsional strain between cy-
clobutane and cyclobutene.40
The moderately unfavorable equilibrium constant for cy-
clobutene formation contrasts with the most energetically
unfavorable equilibrium in the table represented by dehydration
of cyclobutenol to form the antiaromatic cyclobutadiene (pKde
) 28.4), which is believed to exist in a ground-state triplet
electronic configuration. Although ∆H°f for this product is a
calculated rather than experimental value, much effort has been
expended in establishing a reliable energy for this small
molecule.21
Moving down Table 6 brings us to dehydration of the
aromatic hydrates. As expected these reactions are all highly
favorable with the ease of dehydration reflecting the stability
of the aromatic product, i.e., benzene > naphthalene >
phenanthrene > anthracene. The calculation of these values
provided the main motivation for this paper. Although none have
been directly evaluated experimentally, in the following paper
it is shown that values for the 9,10-anthracene hydrate and the
naphthalene hydrate 16 can be derived from thermodynamic
cycles based on experimental data and provide surprisingly good
agreement with the calculated values (+7.5 compared with a
calculated value of +7.4 for anthracene and -13.7 compared
with -14.2 for naphthalene).41
The equilibrium constants in Table 6 vary over 40 orders of
magnitude between the extremes of aromatic and strained
antiaromatic molecules represented by benzene and cyclobuta-
diene, respectively. Since there is relatively little variation in
the energy of replacing hydrogen by hydroxyl in generating the
alcohol products, this variation reflects mainly the energy of
hydrogenation of the double bonds, which is usually used to
represent π-bond stability especially in aromatic molecules. The
principal differences between energies of hydration and hydro-
genation as measures of stability are the use of free energies of
hydration compared with heats of hydrogenation and that the
hydration reactions refer to aqueous solution and hydrogenation
to the gas phase. Neither factor is likely to have a great influence
on relative values.
a Calculated from free energies of formation. That for vinyl alcohol is
based on the value for acetaldehyde and a tautomeric constant KT ) 5.89
× 10-7 (Chiang, Y.; Hojatti, M.; Keeffe, J. R.; Kresge, A. J.; Schepp, N.
P.; Wirz, J.: J. Am. Chem. Soc. 1987, 109, 4000.
Relative free energies of hydrogenation may be represented
by the free energies of introduction of a double bond into a
hydrocarbon molecule (in place of vicinal hydrogen atoms) as
shown in Figure 2. The relative stabilities of the double bonds
in the products of elimination in Table 6 are shown in the same
way in Figure 4. Again the measurements are reported relative
to the free energy of double bond formation in ethane (or
hydrogenation of ethene) denoted δπ above. The range of free
energies is numerically a little greater than that of values of
log Kde as a consequence of the relationship ∆G ) 1.364 log
Kde, where the factor 1.364 corresponds to 2.303RT at 25°.
five-, and six-membered rings. Again we consider the double
bond as a limiting case within a comparison of ring sizes as if
it were a two-membered ring. In fact, both vinyl alcohol and
cyclopropanol show a much greater resistance to dehydration
than does ethanol, consistent with the large ring strain present
in a small ring. The magnitude of the ring strain accompanying
formation of cyclopropene from cyclopropanol is revealed by
an equilibrium constant that is 18 orders of magnitude less
favorable than for the formation of propene from 2-propanol.
Interestingly, the reaction of vinyl alcohol to form acetylene is
more favorable than formation of cyclopropene from cyclopro-
panol by 5 orders of magnitude. This implies that there is less
strain in the “two-membered” than three-membered unsaturated
(40) Wiberg, K. B.; Fenoglio, R. G. J. Am. Chem. Soc. 1968, 90, 3395.
(41) MacCormack, A. C.; McDonnell, C. M.; More O’Ferrall, R. A.; O’Donoghue,
A. C.; Rao, S. N. J. Am. Chem. Soc. 2002, 124, 8575 (following paper in
this issue).
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8572 J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002

Dehydration of Adducts of Aromatic Double Bonds
A R T I C L E S
absorbances determined from a best fit of the dependence of absorbance
upon time to an exponential increase or decay using a weighted least
squares computer program.
Equilibrium Constants for Hydration. For a solution of cyclo-
hexadiene in 2.51 M HClO4 at 40 °C Jensen and Le Carre´15 reported
that the absorbance at λmax ) 257 nm falls to 15% of its value after
eight half-lives of reaction. In a subsequent paper Jensen et al. reported
KH O ) [cyclohexenol]/[cyclohexadiene] ) 8.8 and 13.7 at 4.0 and 5.6
2
M H2SO4, respectively (at 30 °C). We found erratic kinetic behavior
of cyclohexadiene at low acid concentrations, but a satisfactory reaction
of a solution of cyclohexenol (10.3 mg/10 mL CH3CN diluted by 0.1-
2.1 mL with HClO4). This gave limiting absorbances 0.16 at 2.39 M,
0.165 at 4.0 M, and 0.16 at 5.22 M HClO4. Based on ꢀ ) 3000 for
cyclohexadiene an average value of K ) 0.1 was evaluated.
For trans-2-methylstyrene the equilibrium for hydration to 1-phe-
nylpropan-1-ol (5) lies further in favor of the alkene than for
cyclohexadiene, and equilibrium constants were based on initial and
final absorbances for the reaction. Measured changes in absorbances
in concentrated solutions of HClO4 were 1.01 f 0.69 (6.97 M); 0.99
f 0.71 (6.77 M), and 0.99 f 0.68 (5.98 M). Taking 0.006 as the
Figure 4. Additional values of δπ (kcal mol-1) for replacement of two
hydrogen atoms of a hydrocarbon by a double bond (relative to ethane and
ethylene, ∆Gπ ) 23.4 kcal mol-1). *Parent hydrocarbon, 9,10-dihydroan-
thracene.
absorbance of the 1-phenylpropan-1-ol gives an average Kde ) 2.3 (KH O
2
The value of δπ ) -34.7 kcal mol-1 for benzene (compared
with ethylene) is close to the traditional value of 36 kcal mol-1
for the resonance energy of benzene. The large positive energies
of δπ for cyclopropene and cyclobutadiene again emphasize the
high level of strain (and in the latter case presumably avoided
antiaromaticity) inherent in these molecules. The contrasting
stability of the double bond in cyclobutene (3 kcal mol-1 less
stable than in an open-chain structure) is again noteworthy.
In conclusion, we note that in principle any organic functional
group can be submitted to the same treatment as the vinyl and
hydroxyl groups described in this paper and summarized in
Figures 1, 2, and 4. This is a potentially valuable tool for
extending the range of available thermodynamic data and of
equilibrium constants for organic reactions in aqueous solution.
Molecular mechanics and ab initio calculations are now well-
established as methods for accessing unmeasured thermody-
namic data. In this paper we hope to have demonstrated that
extrapolation from existing measurements based on substituent
effects may offer a complementary means for achieving the same
end. As emphasized in the Introduction to the paper a substituent
treatment offers a significantly more precise method of estima-
tion than one based on group contributions. In principle there
is no limitation to the scope of the method provided that a
sufficient body of data exist for related structures incorporating
the functional group of interest, as is the case for alcohols and
alkenes.
) 0.43). The limiting absorbances at completion of reaction were
extrapolated from absorbance measurements over 2.5-3.5 half-lives
of the first-order reactions: no correction was made for isomerization
of trans- to cis-R-methylstyrene.
Cyclopentadiene reacted fairly rapidly in more dilute acid solutions
and multiple determinations of initial and final absorbances at different
acid concentrations were easily made: 0.46-0.386, 0.466-0.365,
0.467-0.355 (1.0 M); 0.480-0.367, 0.497-0.383, 0.489-0.391 (1.5
M); 0.634-0.502, 0.495-0.354, 0.486-0.387 (2.0 M); 0.646-0.499,
0.536-0.420; 0.494-0.392 (2.5 M). These gave an average value of
Kde ) [cyclopentadiene]/[cyclopentenol] ) 3.5.
An equilibrium constant was also measured for the hydration of
trans-piperylene (trans-1,3-pentadiene) to 3-penten-2-ol (6). The identity
of the product was confirmed by the observation that the rate of
approach to equilibrium was the same starting with the 3-penten-2-ol
as starting with the diene. Reaction of the isomeric 1-penten-3-ol (eq
18) was shown to be too slow for this isomer to be implicated in the
equilibrium. Hydration of cis-piperylene was also shown to be
approximately five times slower than the reaction of its trans-isomer.
This is probably because the cis configuration inhibits achievement of
a planar conformation for the incipient allylic carbocation in the
transition state. The difference in stabilities of cis- and trans-piperylenes
in the gas phase corresponds to an equilibrium constant of 7.2 favoring
the trans-isomer. If this is also true in solution, then it is easy to show
that the trans-isomer is equilibrated with 3-penten-2-ol 36 times (5 ×
7.2) more rapidly than with its cis-isomer provided that, as must be
true,2,42 reaction of the carbocation intermediate to form the alcohol is
faster than to form either diene. It follows that the equilibrium observed
is that between 3-penten-2-ol and trans-piperylene only.
Experimental Section
Equilibrium measurements were carried out for reactions in which
the reactant was either the diene or the alkenol. Starting with the diene
a satisfactorily stable dependence of absorbance upon time was observed
only above 2.5-3 M concentration of HClO4, while at concentrations
above 6 M anomalous product absorbances suggested that the diene
was undergoing a further reaction, probably polymerization.
A series of measurements for reactions of aqueous solutions of trans-
piperylene with an initial absorbance of 1.4 gave the following final
absorbances (A) at λmax ) 222 nm at the concentrations of perchloric
acid indicated in brackets (T ) 25 °C): 0.21 (2.15 M); 0.20 (3.15 M);
0.135 (4.2 M); 0.165 (5.52 M). Based on the relationship Kde ) A/(1.4
- A) an average equilibrium constant for dehydration Kde )([diene]/
[alkenol]) ) 0.14 was obtained.
Experimental measurements were confined to measurements of rate
constants and equilibrium constants for hydration of alkadienes and
arylalkenes to alkenols and aryl alcohols, respectively. Most of the
compounds studied were purchased (usually from Aldrich) and used
without further purification. Cyclopentadiene was freshly prepared from
distillation of its dimer. Normally identity and purity were checked by
NMR.
Kinetic and equilibrium measurements were made spectophotometri-
cally using a Phillips PU8600 or Perkin-Elmer Hitachi Model 124
spectrophotometer equipped with a thermostated cell compartment.
Measurements were normally initiated by injection of 20-25 µL of a
10-2 M solution of diene or alkenol substrate in acetonitrile or methanol
into 2 mL of aqueous HClO4 in a spectrophotometric cell prethermo-
stated at 25 °C. Reactions were monitored for at least three half-lives
(more for equilibrium measurements) and rate constants and limiting
Corresponding reactions of 3-penten-2-ol gave the following limiting
absorbances at λ ) 222 nm for a 1.72 × 10-4 M aqueous solution of
(42) Richard, J. P.; Jencks, W. P. J. Am. Chem. Soc. 1984, 106, 1373.
9
J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002 8573

A R T I C L E S
Dey et al.
alkenol and concentrations of perchloric acid indicated in brackets:
0.182 (1.05 M); 0.230 (2.1 M); 0.200 (3.15 M); 0.229 (4.20 M); 0.220
(5.25 M); 0.193 (6.30 M). Based on an extinction coefficient of 1.0 ×
104 for trans-piperylene these values again gave an average value of
Kde ) 0.14.
constants against the medium acidity parameter X.43 The desired rate
constants were obtained as extrapolated values at X ) 0 and are recorded
in the results section of the paper.
In addition to rate constants for the reactions for which equilibrium
measurements have already been described rate constants are reported
for hydration of cis-piperylene. These are not as precise as the other
measurements because hydration is followed by a more rapid but still
kinetically significant conversion of the alcohol to trans-piperylene so
that the reaction is not strictly first order. The measurements demon-
strate, however, that reaction of the cis-piperylene is indeed slower
than that of its trans-isomer, and since the second reaction is five or
six times faster than the first, the error in identifying the measured
rate constant with hydration of the cis-isomer is small.
We assumed that the equilibrium constant for hydration of trans-
piperylene (1/Kde) corresponds to an equilibrium between trans-
piperylene and an equilibrated mixture of cis- and trans-3-penten-2-
ol. Formation of a carbocation intermediate can be expected to occur
considerably more rapidly from the alcohol than the diene because
conversion of the carbocation to alkene is normally rate-determining
in dehydration reactions.2,39 Since the equilibrium constant between cis-
and trans-3-penten-2-ol appears not to have been measured it was
assumed to have a value K ) [trans]/[cis] ) 3.5, which is intermediate
between the gas-phase values of 3.2 and 4.0 for cis- and trans-2-pentene
and cis- and trans-4-methyl-2-pentene, respectively, which could be
calculated from known values of ∆G °f in the gas phase. Correcting
for this equilibration the experimental equilibrium constant Kde ) 0.14
becomes Kde ) 0.18 for the ratio of trans-1,3-pentadiene to trans-3-
penten-2-ol.
Kinetic Measurements. In addition to the equilibrium measurements
rate constants for the approach to equilibrium were measured for the
hydration or dehydration reactions studied. First-order rate constants
(kobs) measured at 25 °C for different concentrations of aqueous HClO4
are listed in Table S1 of the Supporting Information. The measurements
were made spectrophotometrically as described above. Since most of
the measurements refer to concentrated acid solutions values for dilute
aqueous solution were obtained from plots of second-order rate
Acknowledgment. This research was supported by the
National Science Council for Ireland EOLAS (Grant No. SC/
91/204). We thank the reviewers for helpful criticisms and
suggestions.
Supporting Information Available: Table S1, listing first-
order rate constants for hydration or dehydration of various
substrates in concentrated HClO4 (PDF). This information is
available free of charge via the Internet at http://pubs.acs.org.
JA0126125
(43) Cox, R. A.; Yates, K. J. Chem. Soc. 1978, 100, 3861. Kresge, A. J.; Chen,
J. J.; Capen, G. L.; Powell, M. F. Can. J. Chem. 1983, 61, 249. Bagno, A.;
Scorrano, G.; More O’Ferrall, R. A. ReV. Chem. Intermed. 1987, 7, 313.
Cox, R. A. AdV. Phys. Org. Chem. 2000, 34, 1.
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8574 J. AM. CHEM. SOC. VOL. 124, NO. 29, 2002