Ethyl cyclohexanone-2-carboxylate in aqueous micellar solutions. 1: Ester hydrolysis in cationic and nonionic micelles

Article abstract of DOI:10.1021/jp010074c

Hydrolysis rates for the ester of ethyl cyclohexanone-2-carboxylate (ECHC) were measured in aqueous dioxane mixtures and in aqueous micellar solutions of cationic, nonionic, and mixed cationic-nonionic detergents. The following cationic surfactants were used: dodecyltrimethylammonium bromide (DTABr), tetradecyltrimethylammonium bromide (TTABr), tetradecyltrimethylammonium chloride (TTACl), tetradecyltrimethylammonium acetate (TTAAc), and cetyltrimethylammonium bromide (CTABr). Poly(ethylene oxide)-9-dodecyl ether (C₁₂E₉) and poly(ethylene oxide)-20-cetyl ether (C₁₆E₂₀) were used as the nonionic surfactants. The kinetic profiles of the reactions were investigated as a function of surfactant concentration. In aqueous strong mineral acid (HCl or HBr) solutions of the cationic surfactants DTABr, TTABr, TTACl, and CTABr, the reaction was greatly inhibited at surfactant concentrations above the critical micelle concentration; the same behavior was observed for the influence of the nonionic surfactants and for the mixed surfactants DTABr + C₁₂E₉ analyzed in an equimolar mixture of both surfactants. Micellar effects were analyzed by means of the pseudophase model. When the hydrolysis reaction was studied in aqueous buffered solutions of acetic acid-acetate, the shape of the observed rate constant versus surfactant concentration (of TTACl and TTABr) profiles depended on surfactant concentration. At low surfactant concentration, first-order rate constants pass through maxima with increasing surfactant concentration followed by a gradual but steady decrease in the rate as the surfactant concentration increases further. By contrast, the first-order rate constant increases throughout the whole concentration range of TTAAc used. Micellar effects observed for the reaction in aqueous buffered solutions were analyzed qualitatively or quantitatively by means of the pseudophase ion exchange (PPIE) model. In every experimental situation, the reaction at the micellar interface is not negligible; however, the second-order rate constant in this region is lower than that measured in water.

Full text of DOI:10.1021/jp010074c

J. Phys. Chem. B 2001, 105, 10287-10294
10287
Ethyl Cyclohexanone-2-carboxylate in Aqueous Micellar Solutions. 1. Ester Hydrolysis in
Cationic and Nonionic Micelles
Emilia Iglesias†
Departamento de Qu´ımica F´ısica e E.Q. I, Facultad de Ciencias, UniVersidad de La Corun˜a,
15071-La Corun˜a, Spain
ReceiVed: January 7, 2001; In Final Form: August 14, 2001
Hydrolysis rates for the ester of ethyl cyclohexanone-2-carboxylate (ECHC) were measured in aqueous dioxane
mixtures and in aqueous micellar solutions of cationic, nonionic, and mixed cationic-nonionic detergents.
The following cationic surfactants were used: dodecyltrimethylammonium bromide (DTABr), tetradecyl-
trimethylammonium bromide (TTABr), tetradecyltrimethylammonium chloride (TTACl), tetradecyltrimethyl-
ammonium acetate (TTAAc), and cetyltrimethylammonium bromide (CTABr). Poly(ethylene oxide)-9-dodecyl
ether (C12E9) and poly(ethylene oxide)-20-cetyl ether (C16E20) were used as the nonionic surfactants. The
kinetic profiles of the reactions were investigated as a function of surfactant concentration. In aqueous strong
mineral acid (HCl or HBr) solutions of the cationic surfactants DTABr, TTABr, TTACl, and CTABr, the
reaction was greatly inhibited at surfactant concentrations above the critical micelle concentration; the same
behavior was observed for the influence of the nonionic surfactants and for the mixed surfactants DTABr +
C12E9 analyzed in an equimolar mixture of both surfactants. Micellar effects were analyzed by means of the
pseudophase model. When the hydrolysis reaction was studied in aqueous buffered solutions of acetic acid-
acetate, the shape of the observed rate constant versus surfactant concentration (of TTACl and TTABr) profiles
depended on surfactant concentration. At low surfactant concentration, first-order rate constants pass through
maxima with increasing surfactant concentration followed by a gradual but steady decrease in the rate as the
surfactant concentration increases further. By contrast, the first-order rate constant increases throughout the
whole concentration range of TTAAc used. Micellar effects observed for the reaction in aqueous buffered
solutions were analyzed qualitatively or quantitatively by means of the pseudophase ion exchange (PPIE)
model. In every experimental situation, the reaction at the micellar interface is not negligible; however, the
second-order rate constant in this region is lower than that measured in water.
Introduction
lyzed in water or that the ketonization reaction is slower than
the ester hydrolysis. Therefore, the presence of micelles should
have no effect on its UV-vis absorption spectrum, in contrast
to what occurs with, e.g., benzoylacetone.10 However, this
compound is unstable in water, and spontaneous ester decom-
position occurs at moderate rates. The reaction is catalyzed by
strong mineral acids, even though the rate of the spontaneous
ester hydrolysis is comparable to the acid-catalyzed pathway
at moderate acid concentrations (e.g., 0.1 M). The experimen-
tally determined first-order rate constant is ko ) kow + kH[H+],
One of the most important properties of aqueous micellar
solutions is their ability to solubilize a wide variety of organic
solutes normally not dissolvable in water.1,2 This phenomenon
is readily noticeable through chemical reactivity studies. Aque-
ous micelles may increase rates of bimolecular reactions by
concentrating both reactants at the interface of micellar ag-
gregates, or they may decrease rates if both reactants are kept
apart due to their differing hydrophobic/hydrophilic characters
or electrostatic effects.3-7 In these studies, a micelle is treated
as a reaction medium distinct from bulk water, whose properties
gradually vary, from the purely nonpolar inner core (a hydro-
carbon-like solvent) out to the partially polar surface. Between
these two sites, a micelle offers a gamut of solubilization
environments. Reactant concentrations at micellar interfaces can
be determined experimentally or estimated, and second-order
rate constants can then be calculated. To do this, one must
usually define the partial molar volume of the interfacial reaction
region, V, in the micellar pseudophase where the reaction takes
place to refer to concentrations in the micellar phase.8,9
Previous studies demonstrate that aqueous micellar solutions
strongly alter the keto-enol equilibrium of â-diketones by
increasing the enol tautomer concentration,10-12 thus altering
their reactivity toward electrophiles.13,14
where ko ) 0.987 × 10-3 s-1 and kH ) 8.98 × 10-3 mol-1
w
dm3 s-1. The kinetic hydrogen isotope effects agree with the
generally accepted mechanism for acid-catalyzed hydrolysis of
esters: a rapid equilibrium step involving protonation of the
ester on the carbonyl group, and subsequent hydration of the
cation-like species thus formed to yield the corresponding
alcohol and caboxylic acid. In aqueous buffered solutions of
organic acids (acetic acid and its chloro derivatives), the reaction
proves to be strongly general base catalyzed; i.e., a new reaction
path through the basic form of the buffer (e.g., acetate ion)
arises. A good Brønsted plot is obtained with a slope of â )
0.5, indicating that the attack of the basic form of the buffer is
involved in the rate-controlling step of the reaction; that is, ester
hydrolysis goes through tetrahedral intermediates (Scheme 1).
In a recent study,15 we showed that the â-keto ester, ethyl
cyclohexanone-2-carboxylate (ECHC), is almost entirely eno-
In this study we report the results of the influence of cationic,
nonionic, and mixed cationic-nonionic surfactants, forming
micelles, on the ester hydrolysis reaction of ECHC. The presence
of either cationic or nonionic micelles with unreactive counte-
† E-mail: qfemilia@udc.es.
10.1021/jp010074c CCC: $20.00 © 2001 American Chemical Society
Published on Web 09/27/2001

10288 J. Phys. Chem. B, Vol. 105, No. 42, 2001
Iglesias
SCHEME 1: Carboxylic Esters: Mechanism of Acid- and Base-Catalyzed Hydrolysis
rions strongly inhibits the acid-catalyzed hydrolysis reaction.
By contrast, small concentrations of cationic micelles catalyzed
the reaction pathway promoted by the basic form of weak
organic acids; however, the presence of functionalized micelles
of tetradecyltrimethylammonium acetate strongly enhanced the
rate of ester hydrolysis conducted in aqueous buffered solutions
of acetic acid-acetate. The aim of this work is to cover the
constants (ko) were obtained by a nonlinear least-squares fit of
the experimental data (absorbance-time, A-t) by the first-order
rate equation expressed in eq 1
At ) A∞ + (A0 - A∞)e-k t
(1)
o
with A0, At, and A∞ being the absorbance readings at zero, t,
and infinite times, respectively. Values of A∞ were consistent
with the experimental ones. Absorbance changes (∆A ) A0 -
A∞) higher than 0.7 unit were monitored in each kinetic
experiment to ensure accurate values of ko. In the fitting process
of A-t data to eq 1, correlation coefficients higher than 0.9999
were obtained.
behavior of buffer solutions in the presence of surfactants,16
a
situation frequently avoided in studies with aqueous micellar
solutions.
Experimental Section
Ethyl cyclohexanone-2-carboxylate (ECHC), a Merck product
of maximum purity, was used as supplied. Surfactants of the
highest purity available, dodecyltrimethylammonium bromide
(DTABr), tetradecyltrimethylammonium bromide (TTABr),
cetyltrimethylammonium bromide (CTABr), poly(ethylene oxide)-
9-dodecyl ether (C12E9), and poly(ethylene oxide)-20-cetyl ether
(C16E20) were purchased from Sigma and were used without
further purification. Both tetradecyltrimethylammonium chloride
(TTACl) and tetradecyltrimethylammonium acetate (TTAAc)
were prepared through ion exchange of a solution of TTABr,
using an Amberlite IRA-400 (Cl) anionic exchange resin that,
in the case of TTAAc, was previously saturated with acetate
ions. The same kinetic results are obtained when working with
commercial samples of TTACl; the purity of TTAAc has been
probed by titration with HCl. All other reagents were supplied
by Merck and were used as received. Solutions were prepared
with doubly distilled water (first distilled over potassium
permanganate solution and then redistilled).
Results
The influences of several variables were investigated to
elucidate the mechanism of the ECHC hydrolysis reaction in
micelles.
Influence of the Solvent. The ester hydrolysis of ECHC was
studied in aqueous dioxane mixtures at [H+] ) 0.083 M (HCl).
The experimental data of A-t perfectly fit the first-order
integrated equation, and ko decreases as the percentage of
dioxane in the reaction mixture increases; e.g. at approximately
10 M dioxane the observed rate constant is nearly 34 times less
than the value determined in water. However, the ko versus
[dioxane] profile does not indicate any change in the reaction
mechanism. Typical results are shown in Figure 1. The reaction
was also studied in several solvents. The observed rate constants
measured in 83% v/v organic solvent at [H+] ) 0.083 M are
displayed in Table 1.
Pure ECHC was dissolved in dry dioxane (spectrophotometric
grade). The hydrolysis reaction of ECHC was initiated with the
addition of 10 or 15 µL of a solution of ECHC in dioxane into
a thermostated cell containing the acid and the surfactant
solutions, to reach a final reaction volume of 3.0 mL. The ECHC
concentration in the reaction medium was 6.0 × 10-5 M.
The kinetic measurements were recorded with a Kontron-
Uvikon (Model 942) double-beam spectrophotometer, provided
with multiple cell carriers thermostated by circulating water.
The pH was measured with a Crison pH meter equipped with
a GK2401B combined glass electrode and calibrated using
commercial buffers at pH 4.01 and 7.01. All experiments were
performed at 25 °C.
The kinetic runs were performed under pseudo-first-order
conditions, with the acid concentration greatly exceeding the
ethyl cyclohexanone-2-carboxylate concentration. In each kinetic
experiment the integrated method was applied. The decreasing
absorbance due to the reactant ECHC was monitored at 256
nm15 during the course of the reaction. Pseudo-first-order rate
Influence of Cationic Surfactants. The influence of cationic
surfactants on the ester hydrolysis reaction of ECHC was
investigated by using strong mineral acids: HBr with TTABr
and DTABr surfactants, and HCl with TTACl and CTABr
surfactants. (Aqueous solutions of CTABr precipitate in the
presence of HBr at 25 °C). Whatever the cationic surfactant,
the observed rate constant decreased as the surfactant concentra-
tion increased and reached a saturation limit at high [surfactant].
Figure 2 shows typical results obtained from working with
TTACl and TTABr surfactants.
Influence of Nonionic and Mixed Cationic-Nonionic
Surfactants. The acid-catalyzed hydrolysis of the ester function
of ECHC was studied in aqueous micellar solutions of the
nonionic surfactants C12E9 and C16E20 and of the mixed
cationic-nonionic surfactants DTABr-C12E9 at equal molar
ratio (i.e., r ) [C12E9]/[DTABr] ) 1). We used HCl as the strong
mineral acid. The presence of micelles reduced ko values, with
the inhibition effect depending on the nature of the micelles
and on their concentration. Figure 3 shows representative results.

ECHC in Aqueous Micellar Solutions. 1
J. Phys. Chem. B, Vol. 105, No. 42, 2001 10289
Figure 2. Influence of cationic surfactants TTACl (A) and TTABr
(B) on the pseudo-first-order rate constant obtained in the kinetic study
of the acid-catalyzed hydrolysis of ECHC (6.0 × 10-5 M) at [H+] )
0.050 M by using HCl (A) or HBr (B).
Figure 1. (A) Variation of the observed rate constant, ko, obtained in
the acid-catalyzed hydrolysis of ethyl cyclohexanone-2-carboxylate
([ECHC] ) 6.0 × 10-5 M) at [HCl] ) 0.083 M as a function of
[dioxane]. (B) Plot of ln(ko) versus 1/ꢀ (ꢀ, relative permittivity of the
solvent); (2) values obtained in 83% indicated solvent-water mixture
(DMSO, dimethyl sulfoxide; MeOH, methanol; PrOH, propanol).
micellization process, the electrical conductivity of aqueous
solutions of TTAAc was measured at surfactant concentrations
below and above the critical micelle concentration (cmc). The
results are shown in Figure 6. The micellization process was
investigated in pure water as well as in an aqueous buffer
solution of acetic acid-acetate at pH 4.40, with the aim of
reproducing the kinetic conditions.
Data in Figure 6A are representative of a micellization
process: two linear regions of κ vs [surfactant] were obtained.
The least-squares fitting of both linear portions afforded the
following results: (i) above the cmc, κ/mS‚cm-1 ) (0.216 (
0.002) + (26.64 ( 0.08)[TTAAc], r ) 0.99995; (ii) below the
cmc, κ/mS‚cm-1 ) (0.0024 ( 0.0004) + (63.3 ( 0.2)[TTAAc],
r ) 0.99996. The intersection point was used to obtain the cmc,
and the ratio of slopes above (m2) and below (m1) the cmc
provides the micelle ionization degree, R ()m2/m1).17 The
obtained results were cmc ) 5.8 × 10-3 M and R ) 0.42.
TABLE 1: Observed Rate Constants Obtained in 83% v/v
Organic Solvent/Water at [HCl] ) 0.083 M in the Acid
Hydrolysis of ECHC (6.3 × 10-5 M)
solvent
ꢀ
ko/10-5 s-1
dioxane
2.2
20.5
32.7
36.0
46.5
4.95
3.05
6.50
14.7
6.25
propanol
methanol
acetonitrile
dimethyl sulfoxide
a From ref 20.
Reaction in Acetic Acid-Acetate Buffers. Figure 4 shows
typical results of the influence of cationic surfactants on the
ester hydrolysis reaction carried out in buffer solutions of acetic
acid-acetate at pH 4.1 and 4.6 at total [buffer] ) 0.067 and
0.15 M. At low [TTAX] catalysis was observed, and ko goes
through maxima and decreases at high [TTAX]. The maximum
catalytic effect, measured as the ratio komax/ko, is higher with
TTACl than with TTABr; it decreased as the [buffer] increased,
and the [TTAX] required to reach the komax was somewhat higher
with TTACl than with TTABr. The results reported in Table 2
summarize these observations. In C12E9 micelles, ko decreases
for the increasing nonionic surfactant concentration, Figure 5.
All these observations suggest an ionic exchange between
acetate ions and Br- or Cl- ions.
Hydrolysis in Functionalized Micelles. The study of the
reaction in functionalized micelles of TTAAc should provide
valuable information on ion exchange between acetate ions and
surfactant counterions.
This unusual surfactant was prepared through an ion-exchange
resin saturated with sodium acetate (NaAcO). To test the
The electrical conductivity measurements performed in a
buffer of acetic acid-acetate at pH 4.40 and [buffer]t ) 0.10
M are typical of a normal micellization process, Figure 6B. The
break point corresponding to the cmc is not well-defined due
to the higher conductivity of the solvent. However, from the
straight line above the cmc, κ/mS‚cm-1 ) (2.787 ( 0.005) +
(25.58 ( 0.02)[TTAAc] (r ) 0.9998), one may note that the
slope value is quite close to that determined in water. This allows
us to conclude that the presence of the buffer does not modify
the characteristics of TTAAc micelles.
The influence of TTAAc micelles on the hydrolysis reaction
was studied in 0.10 M acetic acid-acetate buffer at pH 4.40.
Figure 7 shows the variation of ko as a function of [TTAAc],
and no maxima is found. Under these experimental conditions

10290 J. Phys. Chem. B, Vol. 105, No. 42, 2001
Iglesias
Figure 3. (A) Influence of the nonionic surfactant, C12E9, on ko at
[H+] ) 0.050 M (HCl). The curve fits eq 2; the inset shows the
linearization of the data according to eq 3; for parameters, see Table
3. (B) Variation of ko as a function of (1) [DTABr] at [HBr] ) 0.050
M; (2) [DTABr] + [C12E9], equimolar mixture of both surfactants at
[HCl] ) 0.050 M; and (b) the same as in (A).
Figure 4. (A) Influence of (O) [TTACl] and of (3) [TTABr]on the
pseudo-first-order rate constant of the base-catalyzed hydrolysis of
ECHC in aqueous buffered solutions of 0.067 M acetic acid-acetate
at pH 4.60. (B) Variation of ko as a function of (O) [TTACl] at [buffer]
) 0.067 M and pH 4.57; (4) [TTABr] at [buffer] ) 0.07 M, pH 4.57;
and (3) [TTABr] at [buffer] ) 0.15 M, pH 4.11. Dotted line to guide
the eye.
there is no possibility of ion exchange between acetate ions and
micellar counterions; thus, ko increases with [surfactant] through-
out the range concentration. The typical observed behavior of
functionalized micelles, or micelles with reactive counterions,
is an increase of ko along with the increasing surfactant
concentration, leveling off at high [surfactant].
micellized surfactant, Dn, forming ECHC complexed to a
micelle, ECHCm, with an association constant Ks. Reaction may
occur in the aqueous or micellar pseudophases, with first-order
rate constants kw and kom, respectively. The rate constant in water
w
is the sum of rate constants via the uncatalyzed (ko ) kH O
-
2
[H2O]) or acid-catalyzed (kH) reaction; i.e. kw ) kow + kH[H+].
Still, as the [H+] at the surface of a cationic micelle must be
negligible, the rate constant for the reaction in the micellar phase
only accounts for the reaction through the attack of water
molecules, that is, spontaneous hydrolysis.
Discussion
The ester hydrolysis reaction is both acid and base catalyzed.
In acid medium the reaction mechanism involves protonation
of the carbonyl group in a rapid and reversible reaction step
followed by attack of water. This mechanism is consistent with
the decrease in ko as both water concentration and solvent
polarity decrease. The transition state is a cation-like intermedi-
ate; consequently a reduction of the relative permittivity of the
solvent causes destabilization of the transition state due to the
reduction of solvating power. The data in Figure 1B show the
typical ln(ko) versus 1/ꢀ profile expected for a bimolecular
reaction between an ion and a polar molecule18 yielding a
transition state more polar than the reactants, so that the rate
constant increases with the increase in the dielectric constant
of the solvent. Values of ꢀ for aqueous dioxane mixtures were
determined following the work of Anderson.19 With the purpose
of including also in Figure 1B the ko values determined in 83%
solvent-water, the ꢀ values corresponding to these mixtures
were taken as for pure solvents.20
The first-order rate constant for the overall reaction, ko, is
given by
m
kw + ko Ks[Dn]
ko )
(2)
1 + Ks[Dn]
If kom were negligible, then plots of 1/ko against [Dn] would be
straight lines. Nevertheless, curves such as those shown in the
inset of Figure 3B for the case of DTABr are observed in every
case. This means that the reaction in the micellar phase is not
negligible. Therefore, the observed rate constants measured as
a function of surfactant concentration can be fitted by using
either eq 2 or its linear form given in eq 3.
1
1
1
Ks
1
[Dn]
)
1 +
‚
(3)
m
(
)
kw - ko
kw - ko
Reaction rates in the presence of cationic micelles were
analyzed quantitatively in terms of the pseudophase model
(Scheme 2), where the substrate ECHC in water associates with
Equation 3 is a type of Lineweaver-Burk representation, and
this data analysis method has the slight disadvantage of leading

ECHC in Aqueous Micellar Solutions. 1
J. Phys. Chem. B, Vol. 105, No. 42, 2001 10291
TABLE 2: Experimental Conditions and Values of ko and
max
ko
Obtained in the Study of the Influence of TTABr and
TTACl Surfactant Concentration in the Ester Hydrolysis
Reaction of ECHC Performed in Acetic Acid-Acetate
Buffers
[buffer]t/
M
ko/10-3 komax/10-3
a
surfactant
pH
s-1
s-1
komax/ko [Dn]max
TTABr
TTABr
TTABr
TTACl
TTACl
TTACl
0.067
0.15
0.067
0.067
0.067
0.070
4.11
4.11
4.60
4.11
4.60
4.57
4.33
8.32
8.10
4.35
7.80
8.00
9.81
15.86
16.03
13.19
20.92
19.90
2.3
1.9
2.0
3.0
2.7
2.5
4.6 × 10-3
4.6 × 10-3
4.6 × 10-3
5.8 × 10-3
5.8 × 10-3
5.8 × 10-3
max
a Micelle concentration ()[surfactant]t - cmc) at which ko
observed.
is
Figure 6. Electrical conductivity measurements of aqueous solutions
of tetradecyltrimethylammonium acetate, TTAAC, as a function of
[TTAAc] measured (A) in water and (B) in a buffer solution of 0.10
M acetic acid-acetate at pH 4.40.
Figure 5. Influence of nonionic surfactant (C12E9) concentration on
the overall rate constant, ko, measured in the base-catalyzed hydrolysis
of ECHC (6.0 × 10-5 M) in aqueous buffered solutions of 0.067 M
acetic acid-acetate at pH 4.57. The curve fits eq 2; for parameters,
see Table 3. The inset shows the linearization according to eq 3.
constant of ECHC enol to micelles is the phase hydrophobicity,
which increases along with the hydrocarbon chain length of the
surfactant. In line with this, Figure 8 shows the correlation
between ln(Ks) and No, the number of carbon atoms in the
hydrocarbon chain of the surfactant. Association to the nonionic
surfactants turns out to be higher than that to the homologous
cationic surfactant. This effect can be explained by considering
the solute hydrogen-bond acidity, which favors incorporation
into the micelles and seems to be more important in nonionic
micelles.
to a plot in which most of the points tend toward one end of
the line. Hence, we used the nonlinear regression analysis of
the experimental data, ko - [surfactant]m according to eq 2.
Curves in Figures 2 and 3 show the corresponding fits; the
Figure 3A inset shows an example of the linear representation
according to eq 3. In each case, satisfactory agreement between
the experimental data and the theoretical treatment was attained.
Aqueous micellar solutions of pure nonionic surfactants or
mixed cationic and nonionic surfactants behave as pure cationic
micellar solutions. Because of the high hydrophilic character
of H+, together with the low polarity of the micellar interface,
the rate of the acid-catalyzed reaction in this region is negligible.
Therefore, the quantitative analysis undertaken with cationic
micelles works here also.
The parameters determined, along with the experimental
conditions, are listed in Table 3.The fact that the reaction in
the micellar phase is not zero indicates that ECHC is dissolved
not in the micellar core but in the micellar interface, where the
presence of water molecules is very important.21,22 This notion
applies to either cationic or nonionic micelles, which indicates
that the higher solubilization of ECHC in the micellar phase is
due to hydrophobic effects. The cyclic structure of the enol of
a â-keto ester (like ECHC), stabilized by intramolecular
H-bonding, strongly favors the intermolecular forces between
solute and micellar interface as compared with solute and water,
thus resulting in a modification of the interfacial partition
coefficient with respect to the hypothetic noncyclic structure.23
The main factor that could control the value of the association
The enols of both benzoylacetone (BZA) and ECHC have
quite similar structures, and both substrates form complexes with
cationic or nonionic micelles yielding quite similar values for
the association constants, Ks.10 A reason for this may emerge if
it is the cyclic structure of the enol that controls the higher
solubility of this tautomer in a less polar medium, such as the
micellar interface. However, greater differences are found in
the complexes of both compounds with â-cyclodextrin (â-
CD): the enol of ECHC forms inclusion complexes15 with â-CD
that are nearly 3 times more stable than those formed by the
enol of BZA.24 Interactions between the secondary -OH groups
in â-CD and the -OEt moiety in the ECHC enol could account
for these differences.
Values of kom listed in Table 3 account for the rates of ester
hydrolysis in the micellar phase under the attack of H2O, but
with the water concentration referring to the total solution
m
w
volume. A comparison between ko and ko ()0.987 × 10-3
s-1) indicates that the former rate constant is more than 20 times
lower than the value of the latter. It should also be noted that
the kom values reported in Table 3 compare quite well with the
ko values in Table 1; i.e. the rate constants observed in the
micellar interface are similar to rates measured in solvents with

10292 J. Phys. Chem. B, Vol. 105, No. 42, 2001
Iglesias
conditions of acetic acid-acetate buffer solutions of total buffer
concentration equal to 0.067 or 0.15 M and at pH values of
4.11 or 4.57, the expression of ko is that which is reported in
eq 4, where [AcO-] ) [buffer]tKa/(Ka + [H+]), with Ka being
the acidity constant of acetic acid, pKa ) 4.77.25
ko ) kow + kB[AcO-]
(4)
In the presence of cationic micelles of TTACl or TTABr, ko
goes through maxima as the surfactant concentration increases;
i.e., ko increases at low [surfactant] and then decreases at high
[surfactant] (see Figure 4). These ko-[surfactant] profiles may
be understood in terms of the pseudophase ion exchange model
(PPIE),26 according to which the surfaces of ionic micelles are
viewed as ion exchangers and competition between reactive and
inert counterions is characterized by an empirical exchange
constant, KI. Here, acetate ions are reactive counterions that
exchange with Cl- or Br- ions when the reaction is carried out
in the presence of TTACl or TTABr, respectively. The exchange
constant between Cl- ions (or Br- ions) at the micellar surface
and acetate ions in water (X-m + AcO-w h X-w + AcO-m, X
AcO
AcO
) Cl or Br) has been determined as KCl
)0.5 (or KBr
)0.098).27 Hence, low [TTAX] means low [Cl-] (or [Br-]) and
competition for the micellar surface is favorable to acetate ions.
This fact, together with the high concentration of ECHC enol
in the micellar pseudophase, causes a catalysis of the reaction;
however, a further increase in [surfactant] also increases the
Cl- (or Br-) ions, with the concomitant reduction of acetate
ions at the micellar interface. Furthermore, the increase of
micelles reduces the ECHC enol concentration at the micellar
pseudophase. On the other hand, since the exchange constant
for Cl- at the micellar surface by Br- in the aqueous phase is
Figure 7. (A) Experimental values (b) of ko measured as a function
of [TTAAc] in the base-catalyzed hydrolysis of ECHC studied in
aqueous buffered solutions of 0.10 M acetic acid-acetate at pH 4.40;
(O) corrected ko values to refer them at pH 4.40. The curve fits eq 5;
for parameters and data, see Tables 3 and 4, respectively. (B) Plot of
kmod {)ko(1 + Ks[TTAAc]m) - kwcor} against the micellized TTAAc
concentration.
Br
KCl ) 5.1, the catalysis is higher with TTACl micelles than
with TTABr micelles.
SCHEME 2: Mechanism of Acid-Catalyzed Hydrolysis
of ECHC in Cationic Micelles
Further evidence of the exchange process is shown by the
data in Table 4. We can see that increasing the [TTACl] causes
a decrease in the pH of the bulk aqueous solution because the
ionization equilibrium of acetic acid is enhanced as more acetate
ions are taken up by the micellar surface, according to the Le
Chaˆtelier principle. In fact, the presence of C12E9 micelles
inhibits the reaction throughout the surfactant concentration
range; see Figure 5 (no ionic exchange is possible). The enol
of ECHC binds strongly to C12E9 micelles, but acetate anions
are not concentrated at the micellar surface; therefore, the
separation of both reactants results in a reduction of the reaction
rate. By fitting the experimental data by eq 2, the following
results are obtained: kw ) (8.0 ( 0.2) × 10-3 s-1; Ks ) 312 (
a water content of approximately 10 M. To compare reactivities
in micelles with those in water, we have to assume a volume
for the micellar phase. With this volume fraction estimated as
V ) 0.15 dm3 mol-1 8
, then Vko ) kH O represents the
m
m
2
bimolecular rate constant between ECHC and H2O at the
micellar interface. This rate constant is 4 times lower than the
m
3 mol-1 dm3, and ko ) (1.5 ( 0.1) × 10-4 s-1; see Table 3.
m
Notice that kom obtained here is approximately 5 times the ko
bimolecular rate constant measured in water (kH O ) 1.78 ×
2
10-5 mol-1 dm3 s-1). The difference probably comes from the
lower hydration power in the micellar interface, which makes
the transition state more unstable, even though the water
nucleophilicity could increase, making the H2O attack on the
ester more effective.
value obtained for the effect of the same surfactant in the
presence of HCl (acid-catalyzed hydrolysis). It is unlikely that
acetate ions would reside at the interface of C12E9 micelles. The
effect could be due to a different stabilization of the transitition
state because, in 83% v/v acetonitrile-water, e.g., the observed
ko has a value quite similar to that of the kom obtained here and
also to the kom value obtained in anionic micelles (see part 228).
When the reaction is performed in aqueous buffered solutions
of acetic acid-acetate, the observed rate constant is given by
ko ) ko + kH[H+] + kB[AcO-], where kB represents the
w
The previous results call for more work with functionalized
micelles of TTAAc. One might expect the reaction to occur in
both aqueous and micellar pseudophases, but with different rate
constants in each medium (see Scheme 3). The spontaneous
decomposition in water represents around 10% of the ko value
under the experimental conditions used in this work. However,
the acetate ion concentration at the interface of TTAAc micelles
catalytic rate constant by acetate ions, AcO-. At the same time,
by considering the values of these rate constants obtained in
w
water,15 namely, ko ) 0.987 × 10-3 s-1; kH ) 8.97 × 10-3
mol-1 dm3 s-1, and kB ) 0.269 mol-1 dm3 s-1 (for acetate ions),
one can easily see that, e.g., at pH >3, the second term in the
previous sum is negligible. Then, under the experimental

ECHC in Aqueous Micellar Solutions. 1
J. Phys. Chem. B, Vol. 105, No. 42, 2001 10293
TABLE 3: Experimental Conditions and Parameters Obtained in the Study of the Influence of Cationic and Nonionic Micelles
on the Acid-Catalyzed ([Acid] ) 0.050 M) and Base-Catalyzed (Acetic Acid-Acetate) Hydrolysis of ECHC
surfactant
acid
cmc/10-3
M
kw/10-3 s-1
Ks/mol-1 dm-3
kom/10-4 s-1
cc
DTABr
TTABr
TTACl
CTABr
C12E9
C16E20
C12E9 + DTABr
C12E9
TTAAc
HBr
HBr
HCl
HCl
HCl
HCl
11.4
0.85
1.0
0.065
0.23
0.008
0.6
1.45 ( 0.01
1.445 ( 0.01
1.44 ( 0.01
1.415 ( 0.013
1.395 ( 0.01
1.395 ( 0.002
1.405 ( 0.015
8.05 ( 0.02
9.5 ( 0.6
219 ( 6
312 ( 7
331 ( 6
446 ( 4
310 ( 5
480 ( 3.5
222 ( 2
312 ( 3
305 ( 9
0.395 ( 0.04
0.34 ( 0.06
0.43 ( 0.05
0.49 ( 0.03
0.305 ( 0.016
0.324 ( 0.015
0.285 ( 0.03
1.53 ( 0.11
376 ( 14
0.9997
0.9995
0.9998
0.9998
0.99995
0.9998
0.9998
0.99995
0.9995
HCl
buffera
bufferb
0.18
2.6 ( 0.7
a 0.067 M acetic acid-acetate at pH 4.57. b 0.10 M acetic acid-acetate at pH 4.40.
SCHEME 3: Mechanism of Base-Catalyzed Hydrolysis
of ECHC in Cationic Micelles
forms micelles, with the monomer concentration remaining
constant and equal to the cmc. Meanwhile, the simple
pseudophase model assumes that the micellar neutralized degree
of ionic micelles, â ()1 - R), remains constant upon increasing
surfactant concentration. In this study, the surfactant counterions
are AcO- ions, and the bulk water phase is an aqueous buffered
solution of acetic acid-acetate: [buffer]t ) 0.104 M, pH 4.40.
Table 4 reports values of ko measured at each TTAAc
concentration, along with the measured pH value of the sample
mixture. Evidently, the [H+] decreases from 4.0 × 10-5 M to
Figure 8. Correlation of values for the association equilibrium constant,
Ks, of ECHC and the number of carbon atoms in the hydrocarbon chain
of the surfactant.
TABLE 4: Pseudo-First-Order Rate Constant Obtained as a
Function of Surfactant Concentration in the Study of the
Hydrolysis Reaction of ECHC in Aqueous Buffered
Solutions of Acetic Acid-Acetate
1.0 × 10-5 M as the [TTAAc] increases from 0.83 × 10-3
M
to 0.208 M; nevertheless, the ionized acetate ions (that come
from the surfactant) increase from 0.83 × 10-3 M to 0.086 M
()cmc + R[TTAAc]m; R ) 0.42). These results suggest that
only a small fraction of the ionized acetate ions is responsible
for the pH variation. This assumption is quite reasonable if we
assume that â is constant. On that basis, a correction of ko values
referred to the same initial pH ()4.40) yields the open points
in Figure 7. A quantitative analysis of these data indicates that
TTACl; [buffer]t
) 0.070 M
TTAAc; [buffer]t
) 0.10 M
[TTACl]/M
ko/10-3 s-1 pH
[TTAAc]/M
ko/10-3 s-1 pH
1.1 × 10-3
2.2 × 10-3
3.3 × 10-3
4.4 × 10-3
6.6 × 10-3
8.8 × 10-3
0.0132
7.98
8.10
4.57 0.833 × 10-3
4.57 1.25 × 10-3
4.56 1.67 × 10-3
4.55 2.50 × 10-3
4.54 3.33 × 10-3
4.53 5.00 × 10-3
4.51 6.67 × 10-3
4.49 0.010
9.42
9.53
9.61
4.40
4.41
4.42
4.41
4.42
4.42
4.43
4.43
4.44
4.45
4.48
4.51
4.59
4.64
4.73
4.81
4.91
5.00
12.13
16.36
19.36
19.88
19.34
18.64
17.23
14.75
13.20
11.36
9.92
10.0
14.0
21.1
25.0
28.7
30.2
31.8
34.1
35.3
38.1
38.9
41.5
43.0
44.5
46.3
cor
w
the term denoted by kw ()ko + kB[AcO-]w) in eq 5 is now
constant. On the other hand, [AcO-]m ) â[TTAAc]m, and
0.0165
kB mAc ) âk2Bm/V, where k2Bm is the second-order rate constant
m
0.022
4.48 0.0125
4.47 0.0167
4.46 0.025
0.033
for the reaction referred to the volume fraction, V, of the micellar
phase where the reaction takes place. The values of the
parameters that best fit the experimental results for the reaction
0.044
0.066
4.46 0.033
0.088
4.45 0.050
0.132
7.81
4.45 0.067
cor
0.165
6.77
4.45 0.100
hydrolysis of the ester of ECHC in TTAAc micelles are kw
0.220
5.60
4.44 0.125
) (9.5 ( 0.6) × 10-3 s-1, âk2Bm/V ) (3.76 ( 0.02) × 10-2
s-1, and Ks ) 305 ( 9 mol-1 dm3; the last two parameters
were treated as adjustable parameters. The solid line in Figure
7 represents the values of ko calculated with these parameters
in eq 5.
0.291
4.73
4.44 0.167
0.208
is much higher than that in the aqueous phase. This fact, together
with a much faster reaction hydrolysis via the general base
catalysis than via the water attack, makes it more reasonable to
neglect spontaneous decomposition at the micellar phase of the
TTAAc micelles. Thus, the observed first-order rate constant
is written as in eq 5.
The fit of theory and experiment is entirely favorable; in
addition, Ks takes a value comparable to that obtained for ECHC
association to TTACl or TTABr micelles because the nature of
the counterion should not strongly affect micelle hydrophobicity.
From the value calculated for âk2Bm/V, and assuming V ) 0.15
dm3 mol-1, one determines k2Bm ) (9.72 ( 0.05) × 10-3 mol-1
dm3 s-1 for the bimolecular reaction between acetate ions and
ECHC at the micellar interface. This rate constant is lower than
the corresponding one for water (kB ) 0.269 mol-1 dm3 s-1) as
(kow + kB[AcO-]w) + kBmKsmAc[TTAAc]m
ko )
(5)
1 + Ks[TTAAC]m
According to the pseudophase model, micelle formation
begins at the cmc, and it is assumed that all additional surfactant

10294 J. Phys. Chem. B, Vol. 105, No. 42, 2001
Iglesias
a consequence of the lower polarity of the micellar interface
and of the lower water content, which decreases the hydration
degree of the leaving group in the transition state, resulting in
a less stable activation complex. Therefore, the lower water
content, together with the low polarity of the micellar interface,
accounts for the estimated reduction in k2Bm, even though the
observed effect is a catalysis due to concentration effects.
Further evidence of the validity of the quantitative treatment
References and Notes
(1) Fendler, J. H.; Fendler, E. J. Catalysis in Micellar and Macromo-
lecular Systems; Academic Press: New York, 1975; Chapter 3.
(2) Tanford, C. The Hydrophobic Effect: Formation of Micelles and
Biological Membranes, 2nd ed.; Krieger Publishing Company, Malabar:
FL, 1991.
(3) Bunton, C. A.; Savelli, G. AdV. Phys. Org. Chem. 1986, 22, 231.
(4) Bunton, C. A.; Nome, F. J.; Quina, F. H.; Romsted, L. S. Acc.
Chem. Res. 1991, 24, 357.
(5) Bunton, C. A. In Kinetics and Catalysis in Microheterogeneous
Systems; Surfactant Science Series 38; Gra¨tzel, M., Kalyanasundaram, K.,
Eds.; Marcel Dekker: New York, 1991; p 13.
(6) Quina, F. H.; Chaimovich, H. J. Phys. Chem. 1979, 83, 1844.
(7) Romsted, L. S. In Surfactants in Solution; Lindman, B., Mittal, K.
L., Eds.; Plenum Press: New York, 1984; Vol. 2, p 1015.
(8) Martinek, K.; Yatsimirski, A. K.; Levashov, A. V.; Berezin, I. V.,
In Micellization, Solubilization, and Microemulsions; Mittal, K. L., Ed.;
Plenum Press: New York, 1977; pp 489-508.
(9) Bunton, C. A.; Carrasco, N.; Huang, S. K.; Paik, C. H.; Romsted,
L. S. J. Am. Chem. Soc. 1978, 100, 5420. Biresaw, G.; Bunton, C. A.;
Quand, C.; Yang, Z.-Y. J. Am. Chem. Soc. 1984, 106, 7178.
is obtained by plotting a modification of ko, komod, defined as
mod
ko
) ko(1 + Ks[TTAAc]m) - kwcor, against [TTAAc]m. To
calculate values of komod, we assume that Ks ) 320 mol-1 dm3,
i.e., the average value obtained with TTACl and TTABr
micelles, and cmc ) 3.0 × 10-3 M, i.e., the lowest [TTAAc]
at which micellar effects have been observed. This graph can
be shown in Figure 7B and features a straight line with a zero
intercept at the origin, as eq 5 predicts. The slope of the
m
corresponding straight line ()kB mAcKs) has been determined
as 12.7 ( 0.1 mol-1 dm3 s-1 (r ) 0.9993), in good agreement
with the results obtained by nonlinear regression analysis.
(10) Iglesias, E. J. Phys. Chem. 1996, 100, 12592.
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York, 1997; Chapter 6.
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2nd ed.; VCH Verlagsgesellschaft: Weinheim, 1988.
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Conclusions
Both cationic and nonionic micelles decrease the rate of the
acid-catalyzed hydrolysis of the ester function of ethyl cyclo-
hexanone-2-carboxylate. This â-keto ester associates strongly
to cationic or nonionic micelles by hydrophobic interactions due
to the greater stability of the enol tautomer in a less polar
solvent, such as the micellar interface. The micelle-solubilized
substrate self-decomposes while the substrate in the bulk water
phase also hydrolyzes via the acid-catalyzed pathway. The
bimolecular rate constant for the H2O attack to the ester residing
on the micellar surface is more than 5 times the value found in
water. The base-catalyzed hydrolysis goes through maxima as
the TTACl (or TTABr) concentration increases. However, the
addition of TTAAc enhances ko values throughout the concen-
tration range. The behavior observed is the consequence of the
ion exchange between acetate ions and Cl- or Br- ions at the
micellar surface. However, the second-order rate constant of
hydrolysis catalyzed by acetate ions at the micellar interface is
lower than the rate constant in water. The lower polarity or lower
water content of the micellar interface explains the difference
in the rate constants, even though the observed effect is a
catalysis as a consequence of concentration effects in the small
volume of the micelle.
Acknowledgment. Financial support from the Direccio´n
General de Investigacio´n (Ministerio de Ciencia y Tecnolog´ıa)
of Spain (Project BQU2000-0239-C02) is gratefully acknowl-
edged.
(27) Bartet, D.; Gamboa, C.; Sepu´lveda, L. J. Phys. Chem. 1980, 84,
272.
(28) Iglesias, E. J. Phys. Chem. B 2001, 105, 10295.