Kinetic evidence for hydrophobically stabilized encounter complexes formed by hydrophobic esters in aqueous solutions containing monohydric alcohols

Article abstract of DOI:10.1021/ja010617w

The pH-independent hydrolysis of four esters, p-methoxyphenyl 2,2-dichloroethanoate (1a), p-methoxyphenyl 2,2-dichloropropanoate (1b), p-methoxyphenyl 2,2-dichlorobutanoate (1c), and p-methoxyphenyl 2,2-dichloropentanoate (1d), in dilute aqueous solution

Full text of DOI:10.1021/ja010617w

11848
J. Am. Chem. Soc. 2001, 123, 11848-11853
Kinetic Evidence for Hydrophobically Stabilized Encounter
Complexes Formed by Hydrophobic Esters in Aqueous Solutions
Containing Monohydric Alcohols
Niklaas J. Buurma, Laura Pastorello, Michael J. Blandamer,† and Jan B. F. N. Engberts*
Contribution from the Physical Organic Chemistry Unit, Stratingh Institute, UniVersity of Groningen,
Nijenborgh 4, 9747 AG Groningen, The Netherlands
ReceiVed March 7, 2001
Abstract: The pH-independent hydrolysis of four esters, p-methoxyphenyl 2,2-dichloroethanoate (1a),
p-methoxyphenyl 2,2-dichloropropanoate (1b), p-methoxyphenyl 2,2-dichlorobutanoate (1c), and p-methox-
yphenyl 2,2-dichloropentanoate (1d), in dilute aqueous solution has been studied as a function of the molality
of added cosolutes ethanol, 1-propanol, and 1-butanol. The rate constants for the neutral hydrolysis decrease
with increasing cosolute concentration. These kinetic medium effects respond to both the hydrophobicity of
the ester and of the monohydric alcohol. The observed rate effects were analyzed using both a thermodynamic
and a kinetic model. The kinetic model suggests a molecular picture of a hydrophobically stabilized encounter
complex, with equilibrium constants Kec often smaller than unity, in which the cosolute blocks the reaction
center of the hydrolytic ester for attack by water. The formation of these encounter complexes leads to a
dominant initial-state stabilization as follows from the thermodynamic model. Decreases in both apparent
enthalpies and entropies of activation for these hydrolysis reactions correspond to unfavorable enthalpies and
favorable entropies of complexation, which confirms that the encounter complexes are stabilized by hydrophobic
interactions.
Introduction
Scheme 1
Hydrophobic interactions are important noncovalent driving
forces for inter- and intramolecular binding and assembly
processes in aqueous chemistry and biochemistry.1 These
interactions vary from relatively weak pairwise intermolecular
contacts to cooperative bulk association processes. The driving
force for these hydrophobic interactions usually originates from
a delicate balance between enthalpic and entropic effects, largely
due to changes in hydration of the interacting solutes. Both
experimental2 and computational studies3 have contributed to
our present understanding of these rather complex phenomena.
In addition to chemical equilibria, hydrophobic interactions
often play a key role in chemical reactions4 and catalytic
processes.5
hydrophobic cosolutes, was chosen for detailed analysis. The
hydrolyses of 1a-d proceed via the mechanism shown in
Scheme 1.6,7
All these reactions are water-catalyzed between pH 2.0 and
5.5. The reactions proceed via a dipolar activated complex in
which two water molecules, one of which, acting as a general
base, are involved with three protons in flight.
The pH-independent hydrolysis of activated esters, p-meth-
oxyphenyl 2,2-dichloroalkanoates 1a-d, in the presence of
Detailed computer simulations, using both quantum and
classical dynamics, revealed that proton tunneling is involved
in the rate-determining step,8 the water molecules involved in
the activated complex being therefore subject to severe orien-
tational requirements. Consistent with these views, strongly
negative entropies of activation have been found for this
reaction.7
Previous studies9 have shown that the hydrolysis of activated
esters structurally similar to 1a-d, but also of similar activated
amides,10 is retarded by most hydrophobic cosolutes.7,9,11-16 In
* To whom correspondence should be addressed. Fax: +31 50 363 4296.
E-mail: J.B.F.N.Engberts@chem.rug.nl.
† Current address: Department of Chemistry, University of Leicester,
Leicester LE1 7RH, U.K. E-mail: MJB@le.ac.uk.
(1) Blokzijl, W.; Engberts, J. B. F. N. Angew. Chem., Int. Ed. Engl. 1993,
32, 1545.
(2) (a) Mayele, M.; Holz, M. Phys. Chem. Chem. Phys. 2000, 2, 2429.
(b) Bagno, A.; Campulla, M.; Pirana, M.; Scorrano, G.; Stiz, S. Chem. Eur.
J. 1999, 5, 1291. (c) Shulgin, I.; Ruckenstein, E. J. Phys. Chem. B 1999,
103, 2496. (d) Pertsemlidis, A.; Saxena, A. M.; Soper, A. K.; Head-Gordon,
T.; Glaeser, R. M. Proc. Natl. Acad. Sci. U.S.A. 1996, 93, 10769. (e) Ansell,
S.; Cser, L.; Grosz, T.; Jancso, G.; Jovari, P.; Soper, A. K. Physica B 1997,
234, 347. (f) Finney, J. L.; Bowron, D. T.; Soper, A. K. J. Phys.: Condens.
Matter 2000, 12, A123. (g) Castronuovo, G.; Elia, V.; Moniello, V.; Velleca,
F.; Perez, C. S. Phys. Chem. Chem. Phys. 1999, 1, 1887. (h) Mochizuki,
S.; Usui, Y.; Wakisaka, A. J. Chem. Soc., Faraday Trans. 1998, 94, 547.
(3) Lazaridis, T. J. Phys. Chem. B 2000, 104, 4964.
(6) Fife, T. H.; McMahon, D. M. J. Am. Chem. Soc. 1969, 91, 7481.
(7) Engbersen, J. F. J.; Engberts, J. B. F. N. J. Am. Chem. Soc. 1975,
97, 1563.
(4) (a) Lubineau, A.; Auge´, J.; Queneau, Y. Synthesis 1994, 741. (b)
Grieco, P. A. Organic Synthesis in Water; Blackie: London, 1998. (c) Li,
C. Chem. ReV. 1993, 93, 2023.
(8) Lensink, M. F.; Mavri, J.; Berendsen, H. J. C. J. Comput. Chem.
1999, 20, 886.
(9) Engberts, J. B. F. N.; Blandamer, M. J. J. Phys. Org. Chem. 1998,
11, 841.
(5) Otto, S.; Engberts, J. B. F. N. J. Am. Chem. Soc. 1999, 121, 6798.
10.1021/ja010617w CCC: $20.00 © 2001 American Chemical Society
Published on Web 11/03/2001

Encounter Complexes Formed by Hydrophobic Esters
J. Am. Chem. Soc., Vol. 123, No. 48, 2001 11849
the present study, these rate retardations are interpreted using
both a thermodynamic model and a kinetic model.17
Scheme 2
The thermodynamic model for interactions between a reacting
molecule and an inert hydrophobic cosolute was developed
several years ago.13,18 This model interprets the rate retardations
in terms of the effect of added cosolute on activity coefficients
of initial and transition states of the ester undergoing hydrolysis.
These coefficients were re-expressed using the procedures
described by Wood19 in terms of pairwise solute-solute
interaction parameters. The analysis leads to
In the analysis of previously reported kinetic data for this
class of systems, emphasis was placed mainly on the hydro-
phobicity of the cosolutes.9 Rate retardations by added cosolutes
follow an additivity scheme in which each methylene unit makes
a common contribution to G(c), the SWAG approach (Savage-
Wood additivity of group interactions).19 For hydrolysis reac-
tions similar to those involving esters 1a-d, the change in
standard Gibbs energy of activation is largely caused by a
stabilization of the initial state by hydrophobic interactions.24
In a different approach, the severe orientational requirements
on the water orientation in the activated complex8 prompts the
idea of formation of an encounter complex between ester and
added solute, in which the cosolute blocks the reaction center
from attack by water.
A kinetic scheme based on this molecular picture (Scheme
2) emerges in which ester molecules that are not solvated by
cosolute molecules react with a rate constant k(mc)0). The
hydrolysis rate constant for the ester in the encounter complex
is assumed to be zero. This assumption leads to the following
expression for the observed rate constant
k(mc)
2
ln
)
2[gcx - gcq]mc - NφMlmc (1.1)
[
]
k(mc)0)
RTmo
Here k(mc) is the (pseudo-)first-order rate constant in an mc
molal aqueous solution of cosolute c, k(mc)0) the rate constant
in the absence of added cosolute, R the gas constant, and T the
temperature in Kelvin. Significantly, [gcx - gcq] is the difference
in interaction Gibbs energies between the cosolute c and the
reactants x on one hand and the activated complex q on the
other hand. Furthermore, Ml is the molar mass of water, N is
the number of water molecules involved in the rate-determining
step, and φ is the practical osmotic coefficient for the aqueous
solution where the molality of added solute is mc. In the present
study, N ) 2 (vide supra). Further, the solutions are very dilute,20
and hence, φ can be taken as unity; mo is the (hypothetical)
ideal reference state and corresponds to 1 mol kg-1. The
difference [gcx - gcq] is denoted as G(c). This analysis of the
kinetic results, which involves a direct link between thermo-
dynamics and transition-state theory, has also been employed
for completely different reactions, including keto-enol tau-
tomerization,21 rate-determining electron-transfer reactions,22 and
aquation of iron(II) complexes in aqueous solutions.23
k(mc)0)
1 + Kecmc
k(mc) )
(1.2)
(10) Karzijn, W.; Engberts, J. B. F. N. Tetrahedron Lett. 1978, 1978,
1787.
(11) Blokzijl, W.; Jager, J.; Engberts, J. B. F. N.; Blandamer, M. J. J.
Am. Chem. Soc. 1986, 108, 6411.
(12) Engbersen, J. F. J.; Engberts, J. B. F. N. J. Am. Chem. Soc. 1974,
96, 1231.
Here Kec is the equilibrium constant for encounter complex
formation in kilograms per mole, mc the molality of added
cosolute, and k(mc) the observed (pseudo-)first-order rate
constant in an mc molal solution.
In the present study, both the hydrophobicity of the cosolute
molecules and of the reacting ester were varied. The results of
the analysis based on both eqns 1.1 and 1.2 are reported.
Furthermore, the isobaric activation enthalpies and entropies for
the hydrolysis of 1c in the presence of hydrophobic cosolutes
were determined in order to obtain more information on the
thermodynamics of encounter complex formation and to un-
derstand the relation between the thermodynamic description
and the molecular picture of rate inhibition. We show that both
approaches account for the kinetic data.
The study of the thermodynamics and kinetics of encounter
complexes in aqueous solution has immediate relevance for a
mechanistic understanding of reactions in aqueous media.
Generally, the formation of an encounter complex constitutes
the first step in the activation process of a bimolecular reaction.
Insight into factors governing encounter complex formation aids
in a quantitative analysis of second-order rate constants for such
chemical transformations.
(13) Blokzijl, W.; Engberts, J. B. F. N.; Blandamer, M. J. J. Phys. Chem.
1987, 91, 6022.
(14) (a) Streefland, L.; Blandamer, M. J.; Engberts, J. B. F. N. J. Phys.
Chem. 1995, 99, 5769. (b) Kerstholt, R.; Engberts, J. B. F. N.; Blandamer,
M. J. J. Chem. Soc., Perkin Trans. 2 1993, 49. (c) Engberts, J. B. F. N.;
Kerstholt, R.; Blandamer, M. J. J. Chem. Soc., Chem. Commun. 1991, 1230.
(15) Benak, H.; Engberts, J. B. F. N.; Blandamer, M. J. J. Chem. Soc.,
Perkin Trans. 2 1992, 2035.
(16) (a) Apperloo, J. J.; Streefland, L.; Engberts, J. B. F. N.; Blandamer,
M. J. J. Org. Chem. 2000, 65, 411. (b) Noordman, W. H.; Blokzijl, W.;
Engberts, J. B. F. N.; Blandamer, M. J. J. Org. Chem. 1993, 58, 7111. (c)
Hol, P.; Streefland, L.; Blandamer, M. J.; Engberts, J. B. F. N. J. Chem.
Soc., Perkin Trans. 2 1997, 485. (d) Blokzijl, W.; Engberts, J. B. F. N.;
Blandamer, M. J. J. Am. Chem. Soc. 1990, 112, 1197.
(17) Correlations between ln(k) and several solvent parameters yield less
satisfactory results. For example, ln(k) for individual probes correlates
reasonably well with the relative permittivity ꢀ for aqueous solutions within
a series of concentrations using only one cosolute. Plotting ln(k) vs relative
permittivity for solutions of different alcohols, however, results in different
correlations for different alcohols.
(18) Blandamer, M. J.; Burgess, J.; Engberts, J. B. F. N.; Blokzijl, W.
Annu. Rep. R. Soc. Chem., Sect. C 1990, 45.
(19) Savage, J. J.; Wood, R. H. J. Solution Chem. 1976, 5, 733.
(20) The concentration range of the cosolute was deliberately kept small
in order to avoid complexities in the kinetic data due to 2:1 and higher
order interactions.
Results and Discussion
(21) Blokzijl, W.; Engberts, J. B. F. N.; Blandamer, M. J. J. Chem. Soc.,
Perkin Trans. 2 1994, 455.
(22) Bietti, M.; Baciocchi, E.; Engberts, J. B. F. N. J. Chem. Soc., Chem.
Commun. 1996, 1307.
(23) Blandamer, M. J.; Burgess, J.; Cowles, H. J.; De Young, A. J.;
Engberts, J. B. F. N.; Galema, S. A.; Hill, S. J.; Horn, I. M. J. Chem. Soc.,
Chem. Commun. 1988, 1141.
Hydrolysis of 1a-d in the Absence of Cosolutes. (Pseudo)-
first-order rate constants at 298.2 K for the hydrolysis of 1a-d
in water are summarized in Table 1.
(24) Karzijn, W.; Engberts, J. B. F. N. Recl. TraV. Chim. Pays-Bas 1983,
102, 513.

11850 J. Am. Chem. Soc., Vol. 123, No. 48, 2001
Buurma et al.
Figure 1. Hydrolysis of 1b (O), 1c (b), and 1d (4) as a function of
the concentration of 1-butanol. The lines are the best fits using eq 1.1.
Figure 2. Absolute values of G(c) in J kg mol-2 for different probe-
cosolute combinations.
Table 1. (Pseudo-)First-Order Rate Constants for the
Water-Catalyzed Hydrolysis of 1a-d in Water at 298.2 K
For the small range of cosolutes studied, however, definite
conclusions about additivity cannot be drawn. The effects of
added longer chain alcohols were not examined because their
solubility ranges are small. As was observed for similar
hydrolytic probe and cosolute systems, the methylene units
closest to the hydrophilic group are partially shielded by the
hydrophilic hydration shell of the polar moiety, reducing their
hydrophobic effect.16 Similar shielding effects by nonionic
hydrophilic groups have been found by other authors.27
compd
104k (mc)0)/s-1
compd
104k (mc)0)/s-1
1a
1b
30.9
11.7
1c
1d
3.06
2.73
Table 2. G(c) Values for the Hydrolysis of Esters 1a-d in
Aqueous Solution at 298.2 K in the Presence of Short-Chain
Alcoholsa
G(c) for given cosolute (J kg mol-2
)
It is possible to write the observed G(c) values (Table 2) as
matrices that can be written as a matrix product (eq 3 is a least-
squares analysis),
ester
EtOH
n-PrOH
n-BuOH
1a
1b
1c
1d
-304(5)
-338(9)
-400(4)
-466(22)
-474(8)
-709(10)
-833(29)
-1044(54)
-1213(70)
-555(22)
-592(22)
-634(52)
-304 -474 -709
-338 -555 -833
-400 -592 -1044
-466 -634 -1213
1
1.15
a The numbers in brackets are standard errors based on a least-squares
fit of the kinetic data using eq 1.1.
) -301.1
1.34
1.51
(
)
(
)
Increasing hydrophobicity of the alkyl chain in the alkanoate
moiety of the ester retards the rate of hydrolysis. Previously,
the influence of alkyl groups on the water-catalyzed hydrolysis
of activated amides was studied using Charton’s expanded
branching equation.25,26 The size of the data set in Table 1 does
not allow a similar analysis. Unfortunately, the data set cannot
be expanded due to the severe solubility problems encountered
with more hydrophobic esters.
Thermodynamic Model. Rate constants for the hydrolysis
of the esters 1a-d decrease upon increasing cosolute concentra-
tion (e.g. Figure 1).20 The decrease in rate constant is more
pronounced for the more hydrophobic cosolutes, in accord with
previous observations.9,14-16
(1 1.50 2.49 ) (1.3)
or, in matrix notation,
G(c) ) aec
(1.4)
Here a is a constant denoting the interaction between p-
methoxyphenyl 2,2-dichloroacetate and ethanol. The vectors e
and c identify the increment in interaction upon increasing the
hydrophobicity of ester and cosolute, respectively, as a multi-
plication factor. In this matrix notation, the SWAG theory should
lead to constant increments in both e and c. Indeed, this pattern
is effectively followed in e, the differences being 0.15, 0.19,
and 0.17 (0.17 ( 0.02), indicating that the interactions between
probe and cosolute are additive with respect to the probe.
However, interactions between probe and cosolute do not seem
to be additive with respect to the cosolutes, most probably as a
result of the small range of cosolutes, which was constrained
by the solubility of the higher alcohols.
Analysis of the kinetic data using eq 1.1 yields the G(c)-
values summarized in Table 2.
Assuming the (standard) chemical potential of the transition
state to be largely unaffected by the cosolute, a negative G(c)
signifies a lowering of the chemical potential of the initial state.
G(c) decreases upon increasing the hydrophobicity of both added
cosolute and ester, indicating increasing stabilization of the
initial state ester. The results are summarized in Figure 2(please
note that -G(c) is plotted in Figure 2).
In Figure 2, the ordinate shows a scale having constant
increments for one methylene unit. The coordinate records G(c)-
values which follow an approximate additivity scheme in accord
with the SWAG theory,19 leading to nearly constant decreases
in G(c) upon lengthening the alkyl moiety in the ester with one
methylene unit (i.e. stepping either down in Table 2 or sideways
in Figure 2).
Molecular Description. The observed decrease in rate
constant upon increasing the hydrophobicity of the ester is
accounted for in terms of the formation of an encounter complex
by hydrophobic probe and cosolute. Encounter complexes are
formed in solution as a result of random movements of
molecules and (de)solvation processes. The chances of encounter
complex formation increase with increasing size and concentra-
tion of the solutes. In fact, the occurrence of encounter
complexes is necessary for any bimolecular reaction to occur
and the concept of encounter complexes is commonly used in
(25) Charton, M. J. Chem. Soc., Perkin Trans. 2 1983, 97.
(26) Mooij, H. J.; Engberts, J. B. F. N.; Charton, M. Recl. TraV. Chim.
Pays-Bas 1988, 107, 185.
(27) Cheng, Y. K.; Rossky, P. J. Biopolymers 1999, 50, 742.

Encounter Complexes Formed by Hydrophobic Esters
J. Am. Chem. Soc., Vol. 123, No. 48, 2001 11851
Figure 3. Hydrolysis of 1b-d as a function of the concentration of 1-butanol. The lines are the best fits using eq 1.2.
Table 3. Thermodynamic Parameters for the Hydrolysis of 1a-d in the Presence of Short-Chain Alcohols
for given cosolute
EtOH
∆ecG° (kJ mol-1
n-PrOH
∆ecG° (kJ mol-1
n-BuOH
∆ecG° (kJ mol-1
)
compd
Kec (mc-1
)
)
Kec (mc-1
)
)
Kec (mc-1
)
1b
1c
1d
0.34 ( 0.02
0.45 ( 0.02
0.51 ( 0.03
2.67 ( 0.15
1.98 ( 0.11
1.67 ( 0.15
0.56 ( 0.05
0.64 ( 0.01
0.71 ( 0.04
1.44 ( 0.22
1.11 ( 0.08
0.85 ( 0.14
0.86 ( 0.07
1.09 ( 0.10
1.21 ( 0.12
0.37 ( 0.20
-0.21 ( 0.23
-0.47 ( 0.25
bimolecular photochemical reactions.28 On the basis of typical
sizes of solvents and solutes, equilibrium constants for formation
of these randomly formed complexes are commonly estimated
to range from 0.2 L mol-1 to values slightly larger than unity.29
In aqueous solution, encounter complexes will be stabilized by
hydrophobic interactions and the stabilization will increase with
an increased hydrophobicity of the encounter complexes con-
stituents.
2.67 1.44 0.37
1
1.98 1.11 -0.21 ) 10.0 - 7.5 1.07 (1 1.11 1.28 )
1.67 0.85 -0.47
(
)
(
)
1.10
(1.5)
or, in matrix notation,
e
c
∆ecG° ) ∆ecG(noninteract) - GGec Gec
(1.6)
The kinetic scheme (Scheme 2), assuming the encounter
complex is inert, is strongly supported by a computer simulation
of the hydrolysis reaction.8 Since the hydrophobic interaction
with the cosolute occurs close to the reaction center, the critical
orientation of the water molecules for attack at the ester carbonyl
group is disturbed and hydrolysis is largely inhibited.30 This
model leads to the kinetic description given by eq 1.2. Nonlinear
least-squares fitting of the observed rate data to eq 1.2 results
in the equilibrium constants and standard Gibbs energies of
encounter complex formation, ∆ecG°, as given in Table 3.
Typical examples of the fits are shown in Figure 3.31
Here, ∆ecG(noninteract) is the unfavorable standard Gibbs
energy term associated with bringing ester and cosolute together
if there were no favorable interactions between the two. In the
present analysis, ∆ecG(noninteract) has been set to 10.0 (-RT
ln(0.018)), corresponding to the chance (based on mole frac-
tions) of finding a cosolute molecule near the reaction center.
∆ecG(noninteract) was restricted as the size of the data set does
not allow independent determination of all variables. G is the
favorable interaction between p-methoxyphenyl 2,2-dichloro-
propanoate and ethanol. Gec and Gec are the increments in
interaction upon increasing the hydrophobicity of ester and
cosolute, respectively. Again, the increment is given as a
multiplication by a number > 1. The interaction becomes more
favorable upon increasing the hydrophobicity of ester and
cosolute, in accord with the encounter complex being increas-
ingly stabilized by hydrophobic interactions.
e
c
The equilibrium constants for formation of pairwise encounter
complexes are in general smaller than unity. The equilibrium
constants increase upon increasing the hydrophobicity of the
ester and/or the hydrophobic cosolute.
Rewriting ∆ecG° in a matrix expression similar to eq 1.3 leads
Activation Parameters. Enthalpies and entropies of activa-
tion for the hydrolysis of 1c as a function of the concentration
of ethanol, 1-propanol, and 1-butanol are summarized in Figure
4. Apparent enthalpies of activation ∆qHapp° for the hydrolysis
reaction according to Scheme 2 are given by
to eq 1.5 as determined using a weighed least-squares fit
(28) (a) Kavarnos, G. J.; Turro, N. J. Chem. ReV. 1986, 86, 401. (b)
Hubig, S. M.; Kochi, J. K. J. Am. Chem. Soc. 1999, 121, 1688. (c) Weng,
H. X.; Roth, H. D. J. Phys. Org. Chem. 1998, 11, 101. (d) Rathore, R.;
Hubig, S. M.; Kochi, J. K. J. Am. Chem. Soc. 1997, 119, 11468.
(29) North, A. M. The Collision Theory of Chemical Reactions in Liquids;
Methuen: London, 1964.
(30) If interaction between ester and cosolute occurs far from the reaction
center, there will be no difference between the interaction with reactant or
activated complex, resulting in the absence of a kinetic effect.
(31) The rate decreases cannot be caused by the decreased water
concentration alone. Based on known densities of aqueous solutions (see
e.g.: Jolicoeur, C., Lacroix, G. Can. J. Chem. 1976, 54, 624), the water
concentration in the dilute aqueous solutions used in the present study can
be calculated. Considering that the hydrolysis reactions are second order
in water, the decreased water concentration in, for example, a 0.57 m solution
of 1-propanol would result in a rate decrease of 6.5%, whereas experimen-
tally rate effects around 25% are found for the different probes.
Kec [R′Y]
∆qHapp° ) ∆qHw° -
∆ecH°
(1.7)
1 + Kec[R′Y]
Here, ∆ecH° is the enthalpy of formation of the encounter
complex and ∆qHw° is the enthalpy of activation for the
hydrolysis reaction in the absence of cosolute.
Using a nonlinear least-squares analysis based on eq 1.7 with
Kec values as obtained from fitting the kinetic data to eq 1.2,
the enthalpies of encounter complex formation were calculated.

11852 J. Am. Chem. Soc., Vol. 123, No. 48, 2001
Buurma et al.
Figure 4. Activation parameters, ∆qHapp (O) and -T∆qSapp (b), at 25 °C of the hydrolysis of 1c in the presence of ethanol, 1-propanol, and
1-butanol.
Table 4. Thermodynamics of Encounter Complex Formation of 1c
with Short-Chain Alcohols
description explains previous observations of the negative G(c),
signifying initial state stabilization, being accompanied by a
strong enthalpic destabilization of the initial state,15 directly in
terms of the thermodynamics of encounter complex formation.
Using the molecular model of encounter complex formation,
the observed thermodynamics, including the G(c) values, can
be fully accounted for.
∆ecG° (kJ mol-1
)
∆ecH° (kJ mol-1
)
Τ∆ecS° (kJ mol-1
)
ethanol
1-propanol
1-butanol
1.98 ( 0.11
1.11 ( 0.08
-0.21 ( 0.23
4.95 ( 0.50
7.71 ( 0.92
6.68 ( 0.82
2.97 ( 0.52
6.60 ( 0.93
6.89 ( 0.85
The inherent advantage of the molecular description is its
possibility for linking the observed kinetics and thermodynamics
to a molecular picture of two interacting molecules. However,
one has to keep in mind that an important contribution to the
thermodynamics of interaction is caused by water molecules
being released from restricted positions in the hydration shells
of those molecules.
Using the standard Gibbs energies of encounter complex
formation, the entropies of encounter complex formation were
obtained, Table 4.
The formation of encounter complexes is enthalpically
opposed and entropically favored, as expected for hydrophobic
interactions1 in which water molecules are liberated from their
orientationally restricted positions in the hydration shells of the
ester and cosolute.32 Increasing the hydrophobicity of the
cosolute results in a more favorable entropic term, while the
changes in the enthalpy are less pronounced. The entropic effect
being most pronounced, leads to a lowering of the standard
Gibbs energy of encounter complex formation and eventually
even to a favorable standard Gibbs energy of encounter complex
formation (Kec > 1). Moreover, from the standard entropy and
enthalpy of encounter complex formation, it is anticipated that
both the entropy and the enthalpy of the initial state are
increased. This, assuming no change in the standard Gibbs
energy of the activated complex, is in accord with the observa-
tion that, with increasing molality of added alcohol, the decrease
in apparent entropy of activation is more pronounced than the
decrease in apparent enthalpy of activation.
Conclusion
Inert cosolutes can influence reactions in solution by forming
encounter complexes. In aqueous solution, these encounter
complexes can be stabilized by hydrophobic interactions. This
results in enhanced cosolute effects on chemical reactions as
the unfavorable entropy term associated with bringing the
molecules together is partially or completely (depending on
concentration) compensated by the release of water molecules
from the hydration shell. For the water-catalyzed hydrolysis of
the activated esters used in the present study, the formation of
encounter complexes, with equilibrium constants Kec often
smaller than unity, leads to an initial state stabilization as given
by G(c). The stabilization of the encounter complex by
hydrophobic interactions results in a decrease in both apparent
enthalpy and apparent entropy of activation.
Comparison of the Models. Both the thermodynamic model
and the molecular description fit the observed rate decreases.
A link between the two descriptions can be derived for Kecmc
< 1,
Experimental Section
Kinetics. Aqueous solutions were prepared by weight immediately
before use. Water was distilled twice in an all-quartz distillation unit.
All reactions were monitored at 288 nm and at 25.0 ( 0.1 °C (in the
determination of the G(c) values) and at least six different temperatures
in the interval between 20.0 ( 0.1 and 50 ( 0.1 °C (except for the 1.5
mol% 1-propanol and the 0.5 mol% 1-butanol solutions for which
measurements were performed at four and five temperatures, respec-
tively). Reactions were followed for at least six half-lives using a Perkin-
Elmer lambda 2, lambda 5, or lambda 12 spectrophotometer. Good to
excellent first-order kinetics were obtained, the error in the rate constants
being 2% or less. Esters were injected as 20-30 µL of stock solutions
containing 1a-d in acetonitrile into about 15 mL of an aqueous solution
of cosolute in the concentration range of 0-2 mol % (up to 1.68 mol
% for 1-butanol, below the solubility limit of 1.92 mol %) followed
by sonication of the solution for 5 min. The sonicated solutions were
centrifuged, decanted, and diluted to about 20 mL. Of the resulting
solution, 6-7 mL aliquots were transferred into a 2.000 cm path length
stoppered quartz cuvette. The resulting concentrations of hydrolytic
probe were about 10-5 mol dm-3 or less. All these precautions were
taken in order to prevent problems due to the low solubility of the
k(mc)
ln
) - ln{1 + Kecmc} ≈ -Kecmc (1.8)
{
}
k(mc)0)
Comparison of eq 1.8 and eq 1.1 shows that the terms in eq 1.1
describing the interaction between ester and alcohol and the
term for the lowering of the water activity in eq 1.8 are replaced
by an equilibrium constant. Hence, the lowering of the standard
Gibbs energy of encounter complex formation, as given by the
increasing equilibrium constants, is equivalent to a stabilization
of the initial state, as revealed by the negative G(c).
The entropy and the enthalpy of the initial state both are
increased by the encounter complex formation between ester
and cosolute, as described before. Therefore, the molecular
(32) The changes in ∆qH° and ∆qS° as a function of cosolute concentra-
tion are consistent with recent computer simulations, which show a favorable
entropy of association of two methane molecules in water, provided a
sufficiently close approach in the aqueous medium. See: Smith, D. E.;
Zhang, L.; Haymet, A. D. J. J. Am. Chem. Soc. 1992, 114, 5875.

Encounter Complexes Formed by Hydrophobic Esters
J. Am. Chem. Soc., Vol. 123, No. 48, 2001 11853
OCH3, s). 13C NMR (CDCl3, ppm): δ 11.9, 17.0, 45.6, 52.8, 84.3,
166.6. IR (CCl4, cm-1): 1748, 1768.
Scheme 3
(c) 2,2-Dichlorobutanoic Acid was synthesized from methyl 2,2-
dichlorobutanoate according to a literature procedure.35 1H NMR
(CDCl3): δ (ppm): 1.19 (3H, CH3CH2, t), 2.47 (2H, CH3CH2CCl2, q),
13C NMR (CDCl3, ppm): δ 7.0, 36.1, 83.4, 165.3, IR (CCl4, cm-1):
1732.
(d) 2,2-Dichloropentanoic Acid was synthesized analogously using
1
methyl 2,2-dichloropentanoate. H NMR (CDCl3, ppm): δ 1.00 (3H,
CH2CH3, t), 1.65 (2H, CH3CH2CH2, sextet), 2.41 (2H, CH2CH2CCl2,
m). 13C NMR (CDCl3, ppm): δ 11.9, 17.1, 45.5, 84.3, 166.6. IR (CCl4,
cm-1): 1734.
(e) 2,2-Dichloropentanoyl Chloride. A mixture of 1.94 g (11.4
mmol) of 2,2-dichloropentanoic acid and 2.74 g (23 mmol) of SOCl2
was refluxed for 3 h. Distillation of the reaction mixture under reduced
pressure gave 1.13 g (6 mmol, 53%) of 2,2-dichloropentanoyl chloride.
13C NMR (CDCl3, ppm): δ 13.1, 18.3, 46.6., IR (CCl4, cm-1): 1779,
1799.
more hydrophobic esters. The pH of all solutions was adjusted to 3.6
( 0.3 using aqueous HCl. The pH was checked again at the end of
each kinetic experiment and was found to be 3.6 ( 0.3, well within
the pH range in which solely water-catalyzed hydrolysis takes place.
Materials. Cosolutes were of analytical grade and were purchased
from Merck. The esters were synthesized using the route shown in
Scheme 3.
The starting materials for the syntheses were purchased from Aldrich
and were used as received. NMR spectra were recorded on Varian
Gemini 200 (1H: 200 MHz) and VRX 300 (1H: 300 MHz) spectrom-
eters. IR spectra were recorded using a Perkin-Elmer 841 infrared
spectrophotometer. Methyl dichloroacetate was obtained by reacting
dichloroacetic acid with thionyl chloride and subsequent esterification
with methanol.33 1H NMR (CDCl3, ppm): δ 3.92 (3H, OCH3, s), 5.97
(1H, HCCl2, s). IR (CCl4, cm-1): 1773, 1753.
(a) Methyl 2,2-Dichlorobutanoate. An adapted literature proce-
dure34 was used. To a solution of 5.8 mL (40 mmol) of anhydrous
diisopropylamine and 20 mL of sodium-dried tetrahydrofuran (THF),
14.4 mL of a 2.5 M solution of BuLi in hexane (36 mmol) was added
slowly at -78 °C. After stirring for 5 min, 4.0 g (28 mmol) of methyl
2,2-dichloroacetate were added and stirring continued for another 15
min. Next, 2.9 mL (28 mmol) of ethyl iodide was added. The mixture
was stirred for another 15 min and then allowed to reach room
temperature. The reaction mixture was poured out into a saturated NH4-
Cl solution, and 60 mL of ether was added. The ether layer was
separated from the aqueous layer and washed with water and brine.
The ether layer was dried over sodium sulfate and filtered, and ether
was removed by evaporation. Distillation in a Kugelrohr apparatus (120
°C, ca. 10 mmHg) gave 4.153 g (24 mmol, 60%) of product. 1H NMR
(CDCl3, ppm): δ 1.16 (3H, CH2CH3, t), 2.46 (2H, CH3CH2CCl2, m),
3.89 (3H, OCH3, s).13C NMR (CDCl3, ppm): δ 8.0, 37.1, 52.8.
(b) Methyl 2,2-Dichloropentanoate was synthesized analogously
using propyl iodide. 1H NMR (CDCl3, ppm): δ 1.00 (3H, CH2CH3, t),
1.59 (2H, CH3CH2CH2, sextet), 2.40 (2H, CH2CH2CCl2, m), 3.89 (3H,
(f) 2,2-Dichlorobutanoyl Chloride and 2,2-Dichloropropanoyl
chloride were synthesized analogously from 2,2-dichlorobutanoic acid
and 2,2-dichloropropanoic acid, respectively.
(1) 2,2-Dichlorobutanoyl Chloride. 1H NMR (CDCl3, ppm): δ 1.20
(3H, CH3CH2, t), 2.53 (2H, CH3CH2CCl2, q). 13C NMR (CDCl3,
ppm): δ 7.0, 36.0, 88.5, 165.5. IR (CCl4, cm-1): 1773, 1802.
1
(2) 2,2-Dichloropropanoyl Chloride. H NMR (CDCl3, ppm): δ
2.36 (3H, CH3CCl2, t). IR (CCl4, cm-1): 1778, 1795.
(g) p-Methoxyphenyl 2,2-Dichloroacetate (1a) was synthesized
according to a literature procedure.6
(h) p-Methoxyphenyl 2,2-Dichloropentanoate (1d). To 3 mL of
absolute ether, equimolar amounts (6 mmol) of 2,2-dichloropentanoyl
chloride and p-methoxyphenol and pyridine were added. The mixture
was stirred for 3 h at room temperature. Pyridine salts were filtered
off, and the solvent was evaporated. The crude ester was dissolved in
petroleum ether (40:60). On cooling, a two-phase system was formed.
The upper colorless layer was separated, and the solvent was removed
by evaporation, yielding the crude ester. The ester was further purified
by column chromatography over silica, using 1:1 CH2Cl2/n-hexane as
the eluent. 1H NMR (CDCl3, ppm): δ 1.07 (3H, CH3CH2, t), 1.82 (2H,
CH3CH2CH2, sextet), 2.53 (2H, CH2CH2CCl2, m), 3.81 (3H, CH3O,
s), 7.00 (4H, phenyl, AB system).
(i) p-Methoxyphenyl 2,2-Dichlorobutanoate (1c) and p-Methox-
yphenyl 2,2-Dichloropropanoate (1b) were synthesized analogously.
1
(1) p-Methoxyphenyl 2,2-dichloropropanoate. H NMR (CDCl3,
ppm): δ 2.43 (3H, CH3CCl2, s), 3.85 (3H, CH3O, s), 7.00 (4H, phenyl,
AB-system).
1
(2) p-Methoxyphenyl 2,2-Dichlorobutanoate. H NMR (CDCl3,
ppm): δ 1.36 (3H, CH3CH2, t), 2.58 (2H, CH3CH2Cl2, q), 3.81 (3H,
CH3O, s), 7.00 (4H, phenyl, AB-system).
Acknowledgment. We thank M. de Rapper for excellent
technical support.
(33) Urry, W. H.; Eiszner, J. R.; Wilt, J. W. J. Am. Chem. Soc. 1957,
79, 918.
JA010617W
(34) Villieras, J.; Disnar, J. R.; Perriot, P.; Normant, J.-F. Synthesis 1975,
524.
(35) Benedetti, M.; Forti, L.; Ghelfi, F.; Pagnoni, U. M.; Ronzoni, R.
Tetrahedron 1997, 53, 14031.