Micellar medium effects on the hydrolysis of phenyl chloroformate in ionic, zwitterionic, nonionic, and mixed micellar solutions

Article abstract of DOI:10.1002/kin.10067

The spontaneous hydrolysis of phenyl chloroformate was studied in various anionic, nonionic, zwitterionic, and cationic aqueous micellar solutions, as well as in mixed anionic-nonionic micellar solutions. In all cases, an increase in the surfactant concen

Full text of DOI:10.1002/kin.10067

Micellar Medium Effects
on the Hydrolysis of Phenyl
Chloroformate in Ionic,
Zwitterionic, Nonionic, and
Mixed Micellar Solutions
´
˜
´
´
´
´
MARIA MUNOZ, AMALIA RODRIGUEZ, MARIA DEL MAR GRACIANI, MARIA LUISA MOYA
Departamento de Qu ı´ mica F ı´ sica, Universidad de Sevilla, C/ Profesor Garc ı´ a Gonz a´ lez s/n, 41012 Sevilla, Spain
Received 18 September 2001; accepted 4 March 2002
DOI 10.1002/kin.10067
ABSTRACT: The spontaneous hydrolysis of phenyl chloroformate was studied in various an-
ionic, nonionic, zwitterionic, and cationic aqueous micellar solutions, as well as in mixed
anionic–nonionic micellar solutions. In all cases, an increase in the surfactant concentration
results in a decrease in the reaction rate and micellar effects were quantitatively explained
in terms of distribution of the substrate between water and micelles and the first-order rate
constants in the aqueous and micellar pseudophases. A comparison of the kinetic data in non-
ionic micellar solutions to those in anionic and zwiterionic micellar solutions makes clear that
charge effects of micelles is not the only factor responsible for the variations in the reaction
rate. Depletion of water in the interfacial region and its different characteristics as compared
to bulk water, the presence of high ionic concentration in the Stern layer of ionic micelles,
and differences in the stabilization of the initial state and the transition state by hydrophobic
interactions with surfactant tails can also influence reactivity. The different deceleration of the
reaction observed in the various micellar solutions studied was discussed by considering these
factors. Synergism in mixed-micellar solutions is shown through the kinetic data obtained in
these media. �C 2002 Wiley Periodicals, Inc. Int J Chem Kinet 34: 445–451, 2002
INTRODUCTION
the interfacial region have major effects on reaction
rates, these being usually called micellar concentration
effects. The interfacial region differs from water as a
reaction medium [6] and this can affect rates of bi-
molecular reactions as well as rates of spontaneous
reactions through the micellar medium effects. Water-
catalyzed processes are interesting in order to study
the latter. The hydrolysis of phenyl chloroformate is
well studied [7–9]. Bunton et al. [8] investigated this
and other spontaneous hydrolyses in cetyltrimethylam-
monium bromide (CTAB), and sodium dodecyl sulfate
Organized assemblies affect the rates of chemical and
photochemical reactions and the position of chemi-
cal equilibria [1–5]. In the case of bimolecular re-
actions the concentration or depletion of reactants in
Correspondence to: Mar ´ı a Luisa Moy a´ : e-mail: moya@us.es.
Contract grant sponsor: D.G.Y.C.Y.T.
Contract grant number: PB98-1110.
Contract grant sponsor: Consejer ´ı a de Educaci o´ n y Ciencia de la
Junta de Andaluc ´ı a.
(SDS) aqueous micellar solutions. Particular attention
Contract grant number: FQM-274.
�
c 2002 Wiley Periodicals, Inc.
was paid to the relation between kinetic micellar effects

˜
MUNOZ ET AL.
4
46
and the mechanism of the processes. Recently, Savelli
et al. [9] found that the inhibition of the phenyl chlo-
roformate hydrolysis by cationic and sulfobetaine mi-
celles increases with increasing head-group bulk and
affinity of the counterion for cationic micelles, these
effects indicate that depletion of water in the interfa-
cial region complements the charge effect in control-
ling reactivity in micelles. No kinetic data in nonionic
micellar solutions are shown. However, a way of inves-
tigating the importance of the contribution of charge–
charge interactions, independently of other micellar
medium effects, is to study the reaction in nonionic
micellar solutions [10]. With this in mind, the spon-
taneous hydrolysis of phenyl chloroformate was car-
ried out in the nonionic micellar solutions of dodecyl
tricosaoxyethylene glycol ether (Brij35) and the esters
of long-chain fatty acids and sorbitan polyethyleneg-
lycol Tween 20, Tween 40, and Tween 80. The re-
action was also studied in tetradecyltrimethylammo-
nium bromide (TTAB) micellar solutions as well as
in the zwitterionic micellar solutions of N-tetradecyl-
N,N-dimethyl-3-ammonio-1-propanesulfonate, SB3-
received. Sodium monomethylsulfate and tetramethy-
lamonium bromide were from Aldrich as were tertbutyl
alcohol, dioxane, and acetonitrile. HBr and HCl were
obtained from Fluka. Water was obtained from a Milli-
pore Milli-Q water system, its conductivity being less
�
6
� 1
than 10 s cm .
Kinetics
Hydrolysisofphenylchloroformatewasrecordedspec-
trophotometrically at 270 nm (appearance of phenol)
with a Unicam UV-2 and Unicam Helios-␥ spectropho-
tometers. The substrate was added to the reaction solu-
tions in 1-cm cuvettes as relatively concentrated solu-
tions in acetonitrile, so that the final reaction mixtures
contained 0.1% of acetonitrile and the final substrate
�
4
� 3
concentration was 10 mol dm . Phenyl chlorofor-
mate reacted readily with hydroxide ions. The reaction
of the substrate with basic impurities was suppressed
�
3
� 3
by the addition of HBr or HCl 1.5 � 10 mol dm
to the reaction medium. It was found that the pres-
ence of an HBr (or HCl) concentration within the range
�
3
� 2
� 3
14, and N-hexadecyl-N,N-dimethyl-3-ammonio-1-
10 –10 mol dm does not affect the rate constant.
The temperature for the kinetic runs was maintained
at 298.2 � 0.1 K by using a water-jacketed cell com-
partment. Rate constants were reproducible within a
precision of better than 4%.
propanesulfonate, SB3-16. The aim of this work was
to compare rate constants in the micellar pseudophases
of ionic, zwitterionic, and nonionic micellar solutions.
Taking into account the uncertainties regarding the ad-
justable parameters from the simulation of kinetic data
pointed out recently [11], to follow the same proce-
dure in obtaining km in all the micellar media was im-
portant. For this reason, and in order to carry out the
experiments under the same working conditions in all
the micellar media, the reaction was also studied in
CTAB, SDS, and SB3-14 micellar solutions in spite
of it being already investigated in these micellar solu-
tions. Finally, the hydrolysis of phenyl chloroformate
was studied in mixed nonionic–anionic micellar solu-
tions of SDS-Brij35 at different molar fractions of the
components. A previous structural study on these bi-
nary mixtures has shown an appreciable synergism at
high mole fractions of the nonionic component [12].
It seemed interesting to investigate if this synergism is
shown through the kinetic data.
RESULTS AND DISCUSSION
Figures 1 and 2 show the kinetic data for the hydroly-
sis of phenyl chloroformate in several aqueous micellar
Kinetic data in water-cosolvent mixtures as well as
in aqueous salt solutions were obtained as a helpful
tool in the discussion of the kinetic data in micellar
solutions. All experiments were done at 298.2 K.
EXPERIMENTAL
Materials
Figure 1 Influence of changes in the surfactant concentra-
tion on the observed rate constant for the spontaneous hy-
drolysis of phenyl chloroformate in various nonionic micel-
lar solutions. Solid lines are obtained from the fittings of the
experimental data by using Eq. (1). T = 298.2 K.
Phenyl chloroformate was obtained from Aldrich. All
the surfactants were obtained from Fluka and used as

MICELLAR MEDIUM EFFECTS ON THE HYDROLYSIS OF PHENYL CHLOROFORMATE
447
Scheme 1
refer to pseudo-first-order rate constants in the aqueous
and micellar pseudophases, respectively. Km is the
equilibrium binding constant of the phenyl chlorofor-
mate molecules to the micelles present in the reaction
medium and it is written in terms of the molarity of mi-
cellized surfactant. Following Scheme 2, the observed
rate constant can be written as [3,16,17]:
Figure 2 Influence of changes in the surfactant concen-
tration on the observed rate constant for the spontaneous
hydrolysis of phenyl chloroformate in various cationic and
sulfobetaine micellar solutions. Solid lines are obtained from
the fittings of the experimental data by using Eq. (1). T =
kW + km Km[Surfactantm]
2
98.2 K.
kobs =
(1)
1
+ Km[Surfactantm]
solutions. Figure 3 shows the kinetic data obtained for
this process in SDS/Brij35 mixed micellar solutions at
various molar fractions of the two surfactants. In all
the cases the process is inhibited by the presence of
micelles in the reaction medium. However, the hydrol-
ysis is not severely inhibited, therefore, it has to take
place in a relatively wet region of the micelle. Since the
hydrocarbon core of the micelles has shown to be dry
In Eq. (1), [Surfactantm] is the micellized surfactant
concentration. SolidlinesinFigs. 1and2wereobtained
from the fittings of the kinetic data by using Eq. (1),
the agreement between the theoretical and the experi-
mental data being good. Data in Fig. 3 were fitted by
using Eq. (1) and the agreement between the theoreti-
cal and the experimental rate constants were also good.
However, solid lines are not shown in order to make the
visualization of the experimental data easier. The cmc
[
13], the hydrolysis takes place in the Stern region.
The reaction under study is represented by Scheme 1
7–9,14,15] which corresponds to an addition–
� 5
values considered for the fittings were 5.8 � 10 mol
[
� 3
� 5
� 3
� 3
dm for Brij35 [18], 5 � 10 mol dm for Tween
elimination mechanism with the addition step being
rate determining. In the presence of micelles, the re-
action occurs as shown in Scheme 2, where kw and km
� 5
2
1
0 [19], 2.3 � 10 mol dm for Tween 40 [19],
� 5
� 3
� 3
� 4
.5 � 10 mol dm for Tween 80 [19], 2.8 � 10
� 5
� 3
mol dm for SB3-14 [20], 2.8 � 10 mol dm for
�
3
� 3
SB3-16 [20], 3 � 10 mol dm for TTAB [21], 7 �
�
3
� 3
� 3
� 3
1
0
mol dm for CTAB [21], 4 � 10 mol dm
4
�
� 3
for SDS [22], 1.45 � 10 mol dm for SDS/Brij35
�
4
(
XSDS = 0.67/XBrij35 = 0.33) [12], 1.0 � 10
mol
�
3
dm
[
for SDS/Brij35 (XSDS = 0.50/XBrij35 = 0.50)
�
5
� 3
12], and 8 � 10 mol dm for SDS/Brij35 (XSDS =
0
.33/XBrij35 = 0.67) [12]. The low acid concentration
present in the micellar media does not affect the cmc
of the nonionic, zwitterionic, and ionic/nonionic micel-
lar solutions, decreasing the cmc of both anionic and
cationic micellar solutions. Table I shows the values
Figure 3 Influence of changes in the total surfactant con-
centration on the observed rate constant for the spontaneous
hydrolysis of phenyl chloformate in SDS/Brij35 mixed micel-
lar solutions at various molar fractions of the two surfactants.
Scheme 2

˜
MUNOZ ET AL.
4
48
Table I Equilibrium Binding Constants and
Pseudo-First-Order Rate Constants for the Hydrolysis
of Phenyl Chloroformate in Several Micellar Solutions
arise from the interaction of the head group of SDS
with those of Brij35 and the activity of SDS in the
micelles is greatly reduced. The extent of transfer of
the components from solution into the micelles is not
at par with the stoichiometric composition, the mole
fractional concentration of SDS in the micelles being
drastically low. This low fraction of SDS in the mi-
celle could explain why the equilibrium binding con-
stant is practically the same in Brij35 micellar solutions
and in mixed micellar solutions with XSDS = 0.50 and
XSDS = 0.33.
Km
km
s
(mol 1 dm3)
�
(10
+3 � 1
)
Surfactant
Brij35
Tween 20
Tween 40
Tween 80
TTAB
CTAB
SB3-14
SB3-16
209 � 15
195 � 10
500 � 53
776 � 54
75 � 16
64 � 7
2.7 � 0.2
2.8 � 0.2
3.0 � 0.3
3.4 � 0.2
3.2 � 0.4
3.3 � 0.3
8.0 � 0.1
8.2 � 0.2
0.22 � 0.2
0.86 � 0.5
2.0 � 0.4
2.1 � 0.4
340 � 31
328 � 34
61 � 3
Now the rate constants for the hydrolysis of phenyl
chloroformate in the micellar pseudophases are consid-
ered. Table I shows that the hydrolysis is retarded by
the presence of micelles, the reaction being slower in
anionic than in cationic micellar solutions. km is sim-
ilar in all the nonionic micellar solutions and close
to those obtained in cationic micellar solutions. Fi-
nally, the hydrolysis of phenyl chloroformate is faster
in the SB3-14 and in SB3-16 sulfobetaine micellar so-
lutions than in the rest of the disperse media studied.
This result was also found by Brinchi et al. in SB3-
14 micellar solutions [9], their km value is in good
agreement with that found in this work. In the mixed
SDS/Brij35 micellar solutions, and according to the
trend found for the Km values, the rate of hydrolysis
increases by decreasing the SDS molar fraction from 1
to 0.67. However, for XSDS = 0.50, XSDS = 0.33, and
XSDS = 0.0 km is practically the same. That is, syner-
gism is shown through the experimental kinetic data. It
would have also been interesting to study CTAB/rij35
(or TTAB/Brij35) mixed micellar solutions. However,
we did not find structural information about these mix-
tures in the literature.
Scheme 1 shows that the addition–elimination
mechanism proposed leads to a polar transition state.
Transfer of this polar transition state from the bulk
aqueous phase to the micelles will be affected by the
different characteristics of the surface region of the mi-
celles. The fractional charge at the reaction center in the
hydrolysis of the phenyl chloroformate was estimated
as in the work of Bunton et al. [25] where the Eyring
equation was used to analyze micellar rate effects.Thus
[25,26]:
SDS
SDS/Brij35(XSDS = 0.67)
SDS/Brij35(XSDS = 0.50)
SDS/Brij35(XSDS = 0.33)
104 � 15
216 � 37
223 � 25
�
3
� 31
at 298.2 K, [HBr or HCl] = 10� 3 mol
k
w
3
= 14.1 � 10
s
�
dm , T = 298.2 K.
of km and Km obtained from the fittings. The values
of these magnitudes for SB3-14 solutions agree with
those previously found [9]. In the case of SDS and
CTAB micellar solutions, the km values found in this
work are smaller than those in the literature [9]. This
could be due to the different acid concentration used in
Ref. [9].
Table I shows that the equilibrium binding constants
are similar, and low, for all the ionic micellar solutions
studied, either anionic or cationic. Low Km values were
also found by El Seoud et al. [15] for 4-nitrophenyl
chloroformate in anionic and cationic aqueous micellar
solutions. This organic substrate resides in the surface
regions of the micelles, as is the case of the phenyl
chloroformate molecules [9]. As shown in Table I, Km
values are higher in sulfobetaine than in cationic micel-
lar solutions. The same result has been found for other
organic substrates [23,24] and this could be related
to the higher hydrophobic character of the micellar
pseudophase of sulfobetaine micellar solutions in re-
spect to that of TTAB or CTAB ones. In nonionic micel-
lar solutions, Km increases when the hydrophobic char-
acter of the micellar pseudophase increases [10a,12].
In regard to the mixtures SDS/Brij35, Km increases in
the mixed micellar solutions with XSDS = 0.67 in re-
spect to pure SDS. However, the equilibrium binding
constant has the same value in mixed micellar solutions
with XSDS = 0.50 and XSDS = 0.33 than in Brij35 pure
micellar solutions. This is in agreement with the fact
that the synergism in these mixed micellar solutions
is appreciable at high mole fraction of the nonionic
surfactant, which declines with increasing fraction of
SDS in the mixture. The overall nonideality seems to
6=
K
km = kw
(2)
Km
where the symbols have the same meaning as in Eq. (1)
6
and K is the (hypothetical) equilibrium constant for
transfer of the transition state between water and mi-
celles, which gives
�
��
�
+
�
6
k
k
K
K
K
m
m
+
=
(3)
�
+
6
m
m
K�

MICELLAR MEDIUM EFFECTS ON THE HYDROLYSIS OF PHENYL CHLOROFORMATE
449
10 3
�
s
� 1
, respectively. These kinetics were followed,
where + and � refer to reactions in cationic and anionic
micellar solutions, respectively. To a first approxima-
tion we assume that the equilibrium association con-
stant Km is the same in cationic and anionic micellar
�
3
as in micellar solutions, in the presence of 10 mol
�
3
dm of HCl to suppress the possible reaction of the
substrate with basic impurities. In all cases, the reaction
rate was slower than in water and isodielectric mixtures
result in different decelerations of the process, as was
expected [14]. Since the interfacial region, where the
reaction occurs, is “less polar” and less aqueous than
water [34], k is expected to be smaller than k , in any
6
solutions (see Table I). If, on the other hand, K is fac-
torized into electrostatic and apolar terms and the apo-
lar term is supposed to be the same in cationic and an-
+
�
ionic micellar solutions, it is possible to relate km /km
to the micellar surface potentials, 9 and the apparent
charge on the organic reaction center in the transition
state. On these assumptions [25]:
m
w
micellar solution. However, these micellar effects are
difficult to rationalize when comparing various poly-
disperse media since the hydrophobicity of the micel-
lar surface is notoriously difficult to define (besides,
they depend on the average solubilization site of the
organic substrate). The ET(30) solvatochromic probe
is regarded as one of the most useful polarity indica-
tors [35]. In previous works [9a,12,6a], the use of this
dye allowed the estimation of the ET parameter and
the effective relative permittivities ε for several micel-
lar solutions. It was shown that ET as well as ε values
are lower for nonionic and sulfobetaine micelles than
for ionic ones. Therefore, “polarity effects,” in terms
of solvatochromic comparison, would result in a larger
inhibition of the reaction in nonionic and sulfobetaine
micelles than in ionic ones.
Other factor to consider is the high ionic strength
present in the Stern layer of ionic micelles, either an-
ionic or cationic [3a,b,7,14]. As a model system for
the Stern region of the micelle, concentrated aqueous
electrolyte solution of a model salt have been used [6c].
The salts were chosen on the basis of their resemblance
to the micellar head group. For both CTAB and TTAB,
tetraethyl trimethylammonium bromide (TMAB) was
chosen. For SDS, sodium monomethylsulfate (NMS)
was employed. It was found that in salt aqueous so-
lutions the presence of TMAB and NMS up to 3.5
�
�
�
�
+
6
k
k
K
19(�Q)
m
+
log
= log
= �
(4)
�
6
58
m
K�
where 19 is the sum of the absolute values of 9 for
cationic and anionic micellar solutions in mV and �Q
is the apparent fractional charge in the transition state.
Considering values of 140 mV for cationic micelles
and � 130 mV for anionic micelles [28–30] we ob-
tain �Q = � 0.25. Of course, this estimated value has
to be considered with care since the treatment used to
obtain it is oversimplified with a lot of assumptions.
Nonetheless, the result indicates that a negative frac-
tional charge is developed at the reaction center. This is
in agreement with the mechanism proposed for the pro-
cess under study in aqueous solutions and it points out
that the mechanism is the same in water and in aqueous
micellar solutions. Considering charge–charge interac-
tions, electrostatic interactions between the negatively
charged carbonyl oxygen of the transition state shown
in Scheme 1 and the ionic surfactant head groups are at-
tractive for cationic micelles and repulsive for anionic
micelles. Therefore, one would expect km to be higher
in cationic than in anionic micellar solutions, as is
found. However, on this basis, the expected trend would
be km(cationic) > km(sulfobetaine) > km(nonionic)
�
3
� 3
mol dm and 4.0 mol dm , respectively, decreases
�
3
� 1
slightly the reaction rate (from 14.1 � 10
s
in pure
�
� 1
3
� 1
� 3
>
km(anionic), since the interfacial electrical poten-
tial is lower in sulfobetaine than in cationic micelles
31] and surface of nonionic micelles has no charge.
However, the experimental trend is km(sulfobetaine)
water to 12.8 � 10
s
in TMAB 3.5 mol dm and
in NMS 4.0 mol dm ). At high
�
3
� 3
to 12.1 � 10
s
[
surfactant concentrations, when most of the organic
substrate molecules are incorporated into the micelles,
the decrease in the observed rate constant for SDS and
CTAB and TTAB micellar solutions is much larger than
that observed for the aqueous salt solutions used as
models. Since the calculated salt concentration in the
micellar Stern region of SDS and CTAB solutions are
>
km(cationic) � km(nonionic) > km(anionic) (see
Table I). This indicates that there are other factors oper-
ating on reactivity and these factors will be considered
next.
Rate constants for spontaneous hydrolyses typi-
cally decrease on addition of organic solvents to water
�
3
2.7–5.4 and 2.9–5.7 mol dm , respectively [6c], the
salt effects in the Stern region of ionic micelles can-
not be the main factor accounting for the inhibition of
the reaction in these systems. Nonetheless, they could
explain the low km values found in cationic and an-
ionic micellar solutions as compared to sulfobetaine
and nonionic micellar solutions.
[32,33]. The hydrolysis under study was carried out in
three water-cosolvent mixtures of the same dielectric
constant, ε = 60 at 298.2 K: H2O-tertbutanol (water
78.5 wt%), H2O-dioxane (water 80 wt%), and H2O-
acetonitrile (water 58.9 wt%). The observed rate con-
�
3
� 1
� 3 � 1
s
stants were 4.7 � 10
s
, 9.1 � 10
, and 1.6 �

˜
MUNOZ ET AL.
4
50
The differences in water availability between bulk
Nonetheless, since the salt effects do not seem to be
important (from the salt effects found in concentrated
aqueous solutions of model salts) km would be expected
to be smaller in sulfobetaine than in cationic micellar
solutions, in contrast with the observations. Maybe in
these micellar systems reaction-specific requirements
on water configuration, water activity, mobility of wa-
ter molecules or stabilization of the transition state and
the initial state by hydrophobic interactions with the
surfactant tails are operative, these being responsible
for the large km values found in SB3-14 and SB3-16
micellar solutions. One can see that the difficulty in
discussing the micellar medium effects stems from the
complexity of the characterization of the interfacial re-
gion of the different micellar solutions as a medium for
kinetics.
water and water as bound to the micelles and the differ-
ences in stabilization of the initial state and the transi-
tion state by hydrophobic interactions with surfactant
tails are other factors to take into consideration. In re-
gard to the former, water concentration in the Stern
region of various micelles have been estimated. For
�
3
CTAB a value of 45 mol dm was calculated [6b,34a],
�
3
whereas for SDS a value of 33 mol dm was found
[
36]. For the nonionic dodecyl heptaoxyethylene gly-
�
3
col ether, C2E7, water concentration was 48 mol dm
[
37]. This would indicate that large concentrations of
waterareavailableintheSternregionofmicelles. How-
ever, not only water concentration can influence on
the spontaneous hydrolysis studied, but also the wa-
ter activity, the reduced mobility of water molecules
and reaction-specific requirements on water configura-
tion can be operative. The contribution of these factors
cannot be estimated for the various micellar solutions
usedasreactionmediumandneithercanthedifferences
in stabilization of the initial state and the transition state
by hydrophobic interactions with the surfactant tails.
Therefore, one can consider their possible influence on
the inhibition of the reaction, but without estimating
the importance of their contributions to the micellar
effects observed.
From the above discussion we concluded that the
micellar effects on the spontaneous hydrolysis of
phenyl chloroformate in cationic, zwitterionic, non-
ionic, and anionic micellar solutions are the result
of several factors operating on reactivity. The exper-
imental trend was km(sulfobetaine) > km(cationic) �
km(nonionic) > km(anionic). The large inhibition of
the reaction found in anionic micellar solutions can
be due to the repulsive charge–charge interactions, to
the presence of a high ionic concentration in the Stern
layer of the anionic micelles as well as to the low po-
larity and low water content in the interfacial region of
these micelles. In this particular case, all the consid-
ered factors influencing on reactivity make the process
slower in the micellar pseudophase than in the aque-
ous phase. In the case of nonionic micellar solutions,
the low polarity and low water content in the interfa-
cial region work decelerating the reaction. No “salt ef-
fects” or charge–charge interactions are operative and,
therefore one would expect the process to be faster in
nonionic than in anionic micelles, as is found. In both
cationic and sulfobetaine micellar solutions, charge–
charge interactions favor the reaction as compared to
anionic and nonionic micellar solutions, although they
are more favorable for cationic than for sulfobetaine
micelles. However, the presence of a high ionic con-
centration in the Stern layer of cationic micelles would
retard the reaction in respect to sulfobetaine micelles.
BIBLIOGRAPHY
1. Kunitake, T.; Shinkai, T. Adv Phys Org Chem 1980, 17,
435.
2
. Fendler, J. H. Membrane Mimetic Chemistry; Wiley:
New York, 1981.
3
. a) Bunton, C. A.; Savelli, G. Adv Phys Org Chem
1
986, 22, 213; b) Bunton, C. A.; Nome, F.; Quina,
F. H.; Romsted, L. S. Acc Chem Res 1991, 24, 357;
c) Romsted, L. S.; Bunton, C. A.; Yao, J. Curr Opin
Colloid Interface Sci 1997, 2, 622.
4
5
. El Seoud, O. A. Adv Colloid Interface Sci 1989, 30, 1.
. Tascioglu, S. Tetrahedron 1996, 34, 11113.
6. a) Zachariasse, K. A.; Phuc, N. Y.; Kozankiewicz, B. J
Phys Chem 1981, 85, 2676; b) Chandhuri, A.; Loughlin,
J. A.; Romsted, L. S.; Yao, J. J Am Chem Soc 1993,115,
8
351; c) Buurma, N. J.; Herranz, A.; Engberts, J. B. F.
N. J Chem Soc, Perkin Trans 2 1999, 113.
7
. a) Kevill, D. N.; D’Souza, M. M. J Chem Soc, Perkin
Trans 2 1997, 1721; b) Queen, A. Can J Chem 1967, 45,
1
619; c) Butler, A. R.; Roberston, I. H.; Bacaloglu, R. J.
J Chem Soc, Perkin Trans 2 1974, 1733.
. Al Lohedan, H.; Bunton, C. A.; Mhala, M. M. J Am
Chem Soc 1982, 104, 6654.
. Brinchi, L.; Di Profio, P.; Micheli, F.; Savelli, G.;
Bunton, C. A. Eur J Org Chem 2001, 1115.
8
9
10. a) Rodr ´ı guez, A.; Mu n˜ oz, M.; Graciani, M. M.;
Fern a´ ndez, G.; Moy a´ , M. L. New J Chem 2001, 25,
1
084; b) El Seoud, O. A.; Ruasse, M.; Possidonio, S.
J Phys Org Chem 2001, 14, 256.
1
1. Brinchi, L.;DiProfio, P.;Germani, R.;Savelli, G.;Gillitt,
N. D.; Bunton, C. A. J Colloid Interface Sci 2001, 236,
8
5.
2. Ghosh, S.; Moulik, S. P. J Colloid Interface Sci 1998,
08, 357.
3. a) Gruen, D. W. R. Prog Colloid Polym Sci 1985, 70, 6;
b) B o¨ cker, J.; Bopp, P. J Phys Chem 1994, 98, 712.
1
1
2
14. Bunton, C. A. In Advances in Chemistry Series 215:
Nucleophilicity; Harris, J. M.; Mc Manus, S. P. (Eds.);

MICELLAR MEDIUM EFFECTS ON THE HYDROLYSIS OF PHENYL CHLOROFORMATE
451
American Chemical Society: Washington, DC, 1987,
Ch. 12.
27. L o´ pez-Cornejo, P.; Jimenez, R.; Moy a´ , M. L.; S a´ nchez,
F.; Burgess, J. Langmuir 1996, 12, 4981.
1
1
5. Possidonio, S.; Siviero, F.; El Seoud, E. O. J Phys Org
Chem 1999, 12, 325.
6. Menger, F. M.; Portnoy, C. E. J Am Chem Soc 1967, 89,
28. Fernandez, M. S.; Fromherz, P. J Phys Chem 1977, 81,
1755.
29. a) Drummond, C. J.; Grieser, F.; Healy, T. W. J Chem
Soc, Faraday Trans 1 1989, 85, 537; b) Almgrem, M.;
Rydholm, R. J Phys Chem 1979, 83, 360.
4
698.
1
1
7. Bunton, C. A. J Mol Liq 1997, 72, 231.
8. Mu n˜ oz, M.; Rodr ´ı guez, A.; Graciani, M. M.; Ortega,
F.; Vazquez, M.; Moy a´ , M. L. Langmuir 1999, 15,
30. Gunnarsson, G.; J o¨ nsson, B.; Wennerstr o¨ m, H. J Phys
Chem 1980, 84, 3114.
7
876.
31. Chorro, N.; Kamenka, K.; Faucompr e´ , B.; Partyka, S.;
Lindheimer, N.; Zana, R. Colloid Surf A 1996, 110, 249.
32. Ingold, C. K. Structure and Mechanism in Organic
Chemistry, 2nd ed.; Cornell University Press: Ithaca,
New York, 1969, Ch. 7.
1
2
2
9. Bhattacharya, S. C.; Das, H. T.; Moulik, S. P. J Pho-
tochem Photobiol A 1993, 71, 257.
0. Rodr ´ı guez, A.; Graciani, M. M.; Guinda, A.; Mu n˜ oz, M.;
Moy a´ , M. L. Langmuir 2000, 16, 3182.
1. Mu n˜ oz, M.; Rodr ´ı guez, A.; Graciani, M. M.; Mozo,
J. D.; Moy a´ , M. L. Langmuir 1998, 14, 3524.
33. Kevill, D. N.; Anderson, S. W. J Org Chem 1991, 56,
1845.
2
2
2. Bunton, C. A.; Wolf, B. J Am Chem Soc 1973, 95, 3742.
3. Weers, J. G.; Rathman, J. F.; Axe, F. V.; Crichlow, C. A.;
Foland, L. D.; Scheving, D. R.; Wiersema, D. J.; Zielske,
A. G. Langmuir 1991, 7, 854.
34. a) Chauduri, A.; Romsted, L. S. J Am Chem Soc 1991,
113, 5052; b) Soldi, V.; Keper, J.; Romsted, L. S. Lang-
muir 2000, 16, 59.
35. Reichardt, C. Solvents and Solvent Efects in Organic
Chemistry, 2nd ed.; VCH: New York, 1988.
36. Angeli, A. D.; Cipiani, A.; Germani, R.; Savelli, G.;
Cerichelli, G.; Bunton, C. A. J Colloid Interface Sci
1988, 121, 42.
2
2
2
4. Graciani, M. M.; Rodr ´ı guez, A.; Mu n˜ oz, M.; Moy a´ ,
M. L. J Colloid Interface Sci 2001, 235, 260.
5. Bunton, C. A.; Gillitt, N. G.; Mhala, M. M.; Moffatt,
J. R.; Yatsimirsky, A. K. Langmuir 2000, 16, 8595.
6. Hall, D. G. J Phys Chem 1987, 91, 4287.
37. Romsted, L. S.; Yao, L. Langmuir 1996, 12, 425.