Effects of pre-micelles of anionic surfactant SDS on the electron transfer reaction between methylene blue and ascorbic acid

Article abstract of DOI:10.1016/j.molliq.2016.10.034

The effect of pre-micellar cluster of the anionic surfactant sodium dodecyl sulphate (SDS) on the electron transfer reaction between methylene blue (MB) and ascorbic acid (AA) in dilute acid medium has been investigated in the temperature range 298–308?K. The reaction is first order each with respect to MB as well as AA. The reaction involves two parallel paths - one uncatalyzed and the other H⁺-catalyzed path resulting in the rate law, k_obs?=?(k₀?+?k₁ [H⁺]) [AA] [MB]. Iodide ion has been found to have a specific accelerating effect on the reaction rate. The reaction appears to take place between undissociated AA molecule and MB⁺/HMB^2?+ cation. Anionic surfactant SDS shows an inhibiting effect on the reaction rate in the pre-micellar region, ultimately leading to a limiting value. The inhibiting effect of SDS has been explained in terms of Piszkiewicz's co-operativity model. The co-operativity index (n) value varies from 1.45 to 1.76 in the studied temperature range. The values of K_S (the dissociation constant of the pre-micelle), n and surfactant concentration lead to the fact that the electrostatic binding in the pre-micelle is reasonably strong and most of the MB molecules remain in the pre-micelle cluster before the electron transfer with AA takes place. The formation of pre-micelle is exothermic in nature and thus favored at lower temperature.

Full text of DOI:10.1016/j.molliq.2016.10.034

Journal of Molecular Liquids 224 (2016) 472–479
Contents lists available at ScienceDirect
Journal of Molecular Liquids
journal homepage: www.elsevier.com/locate/molliq
Effects of pre-micelles of anionic surfactant SDS on the electron transfer
reaction between methylene blue and ascorbic acid
a,
a
b,
Pratik K. Sen ⁎, Piyali Mukherjee , Biswajit Pal ⁎
a
Department of Chemistry, Jadavpur University, Kolkata 700032, India
Department of Chemistry, St. Paul's C. M. College, 33/1 Raja Rammohan Roy Sarani, Kolkata 700009, India
b
a r t i c l e i n f o
a b s t r a c t
Article history:
Received 21 April 2016
Received in revised form 26 September 2016
Accepted 7 October 2016
Available online 11 October 2016
The effect of pre-micellar cluster of the anionic surfactant sodium dodecyl sulphate (SDS) on the electron transfer
reaction between methylene blue (MB) and ascorbic acid (AA) in dilute acid medium has been investigated in the
temperature range 298–308 K. The reaction is first order each with respect to MB as well as AA. The reaction in-
+
volves two parallel paths - one uncatalyzed and the other H -catalyzed path resulting in the rate law, kobs
=
+
0 1
(k + k [H ]) [AA] [MB]. Iodide ion has been found to have a specific accelerating effect on the reaction rate.
+
2+
The reaction appears to take place between undissociated AA molecule and MB /HMB cation. Anionic surfac-
tant SDS shows an inhibiting effect on the reaction rate in the pre-micellar region, ultimately leading to a limiting
value. The inhibiting effect of SDS has been explained in terms of Piszkiewicz's co-operativity model. The co-
Keywords:
Redox
Methylene blue
Ascorbic acid
Anionic surfactant
Pre-micelles
S
operativity index (n) value varies from 1.45 to 1.76 in the studied temperature range. The values of K (the dis-
sociation constant of the pre-micelle), n and surfactant concentration lead to the fact that the electrostatic bind-
ing in the pre-micelle is reasonably strong and most of the MB molecules remain in the pre-micelle cluster before
the electron transfer with AA takes place. The formation of pre-micelle is exothermic in nature and thus favored
at lower temperature.
©
2016 Elsevier B.V. All rights reserved.
1
. Introduction
induced formation of pre-micelles or from the complexation of the sub-
strate with monomeric surfactant or pre-micelles.
The self-organized surfactant assemblies in aqueous medium have
L-ascorbic acid, a water soluble vitamin C, is considered to be an im-
been widely studied in the last few decades due to its enormous indus-
trial applications [1–3]. It is well known that such amphiphilic mole-
cules form aggregates above a definite concentration known as critical
micelle concentration and such aggregates may possess different struc-
tures/shapes. These aggregates or micelles impart a critical role in the
localization and/or dispersion of charges on reactants and their activat-
ed states, and thereby may influence the reaction rates, equilibrium and
concentration/depletion of reactants in the interfacial region [4]. Such
interactions of the micelles in biochemical systems may be compared
to the complex reactions occurring in biological assemblies [5]. That is
why in the last few decades a large amount of effort has been devoted
to understand the nature of interactions in the enhancement/inhibition
of reaction rates in the presence of micelles [6,7]. However, during re-
cent years, there have been a number of experimental findings [8–11]
as well as theoretical ones [12,13] which indicate the appearance of
clusters at surfactant concentrations below the actual CMC, a phenom-
enon referred to as premicellar aggregation. This type of pre-micellar
aggregation has been found to show accelerating/inhibiting effects on
rates of reactions [14–16]. These effects may arise either from substrate
portant micronutrient and plays significant role in physiological system
[
17]. It is a lactone with 2,3-endiol group. Owing to its anti-oxidant
properties it helps in reducing oxidative stress in biological system
17]. It acts as an enzyme cofactor for the biosynthesis of a number of
[
important biochemicals [18]. Thus, the oxidation studies of ascorbic
acid (AA) are important both biochemically and biomedically [19–21].
The oxidation of ascorbic acid is highly dependent on the pH of the re-
action medium. A number of studies have been made to investigate
the kinetics of oxidation of ascorbic acid by different oxidants [22–25].
There are also reports of the oxidation of ascorbic acid by organic dye-
stuff such as methylene blue [26–28]. Methylene blue is a commonly
used water soluble thiazine class of dye and finds use in biology as a
non-toxic stain. Under acidic conditions, it is easily reduced to the color-
less hydrogenated molecule leucomethylene blue by ascorbic acid
which undergoes oxidation [26] to dehydroascorbic acid (Scheme 1).
The above mentioned reports mainly discussed kinetic and mechanistic
aspects of the electron transfer reaction between ascorbic acid and
methylene blue. There is a report showing the determination of vitamin
C in pharmaceutical sample utilizing the micelle catalyzed methylene
blue reduction by ascorbic acid [29]. However no such literature report
is available involving the influence of the surfactant on the reduction of
methylene blue by ascorbic acid in the premicellar region.
⁎
Corresponding authors.
E-mail addresses: senpk@yahoo.com (P.K. Sen), palbiswajit@yahoo.com (B. Pal).
http://dx.doi.org/10.1016/j.molliq.2016.10.034
167-7322/© 2016 Elsevier B.V. All rights reserved.
0

P.K. Sen et al. / Journal of Molecular Liquids 224 (2016) 472–479
473
N
S
O
O
H
+
CH OH
2
(
H C) N
N(CH3)2
3
2
HO
OH
OH
Methylene Blue
Ascorbic Acid
H
N
O
O
O
H
+
CH OH
2
OH
(
H C) N
S
N(CH3)2
H
3
2
O
Dehydroascorbic Acid
Leucomethylene blue
Scheme 1. Schematic representation of the reaction of methylene blue with ascorbic acid.
Since methylene blue is positively charged and contains benzene rings
in its moiety, there is a possibility of formation of pre-micellar aggregates
with anionic surfactant monomers at surfactant concentrations approxi-
mately one tenth to one thousandth of the CMC value [14,15] through
electrostatic as well as hydrophobic interactions. Thus an attempt has
been made to study the influence of the premicelles formed between
the positively charged methylene blue and the anionic surfactant sodium
dodecyl sulphate (SDS) on the kinetics of reduction of methylene blue by
blue (AR, S. D. Fine Chem. Ltd., India) were prepared. Potassium chloride
(AR, E. Merck, Germany) was used to maintain the ionic strength of the
reaction medium. 1 M HCl (AR, Merck, India) was prepared after stan-
dardization with previously standardized sodium hydroxide solution.
A stock solution of 2 × 10 M SDS (high purity, SRL, India) was pre-
pared. All the solutions were prepared using doubly distilled water.
−
3
2.2. Kinetic measurements
L-ascorbic acid and to see whether the results can be analyzed in the light
of Piszkiewicz's co-operativity model [30] of pre-micellar catalyzed
reactions.
All the kinetic investigations were carried out under pseudo-first
order conditions with [AA] ≫ [MB]. A UV–Visible Spectrophotometer
(
Shimadzu, UV-1800) equipped with
a
Peltier controlled
thermostated cell compartment was used to monitor the progress
of the reaction. The course of the reaction was followed spectropho-
2
. Experimental
tometrically by measuring the absorbance (A
ture at 665 nm (λmax of MB) at known time intervals. The pseudo-
first order plots of ln A versus ‘time’ were found to be linear and
from such linear plots the pseudo-first order rate constants (kobs or
) were evaluated. Such values were reproducible to within ± 5%.
t
) of the reaction mix-
2
.1. Reagents
t
All the reagents used were of AR grade. Stock solutions of 0.01 M
−
4
ascorbic acid (extra pure AR, SRL, India) and 1.0 × 10 M methylene
k
ψ
-
-
-
-
0.2
0.4
0.6
0.8
-
4
3
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0.0
5 x 10
x 10
M
M
M
M
M
M
-
1
λmax = 665 nm
-3
2 x 10
-3
-3
-3
3
4
5
x 10
x 10
x 10
in absence of HCl
in presence of 0.3 M HCl
-1.0
-1.2
600
620
640
660
680
700
0
200
400
600
800
1000
1200
1400
Wavelength (nm)
Time (s)
Fig. 1. Absorbance spectra of methylene blue in water and HCl medium. [MB] = 1.0 ×
Fig. 2. Pseudo-first order plot of methylene blue-ascorbic acid reaction at 293 K. Plot of ln A
+
−
5
−5
10
M.
vs. time at different ascorbic acid concentrations. [MB] = 1.0 × 10 M, [H ] = 0.2 M.

474
P.K. Sen et al. / Journal of Molecular Liquids 224 (2016) 472–479
Table 1
Dependence of rate constants on the variation of ascorbic acid concentration.
+
−5
6
5
4
3
2
1
0
[H ] = 0.1 M, [MB] = 1.0 × 10 M.
2
93 K
298 K
03 K
308 K
Temperature (K)
103 [AA] (M)
103
k
obs (s−1
)
2
k )
(M−1 s−1
3
2
2
3
3
93
98
03
08
0.5
0.507
0.986
1.89
3.07
4.05
4.98
0.623
1.25
2.42
3.73
4.95
6.21
0.755
1.57
3.02
4.53
6.26
7.67
0.875
1.76
3.64
5.10
7.01
9.01
1.014
0.986
0.945
1.023
1.013
0.996
1.246
1.250
1.210
1.243
1.238
1.242
1.510
1.570
1.510
1.510
1.565
1.534
1.750
1.760
1.820
1.700
1.753
1.802
1
2
3
4
5
.0
.0
.0
.0
.0
0.5
1
2
3
4
5
.0
.0
.0
.0
.0
0.5
1
2
3
4
5
.0
.0
.0
.0
.0
0.0
0.1
0.2
0.3
0.4
+
[
H ] (M)
0.5
1
2
3
4
5
.0
.0
.0
.0
.0
Fig. 4. Dependence of pseudo-first order rate constant (k ) on acid concentration. Plot of
obs
kobs versus [H+] at four different temperatures. [MB] = 1.0 × 10 M, [AA] = 2.0 × 10
−5
−3
M.
concentration range studied. From the slopes of the plots of absorbance
at λmax) versus [MB], the molar extinction coefficient (ε) for neutral so-
(
5
−1
−1
3
. Results
.1. Absorbance spectra of methylene blue
The absorbance spectra of methylene blue solution in water and also
lution was calculated to be 1.06 × 10 M cm whereas the value was
1
5
−1
−1
.03 × 10 M cm for MB in presence of 0.3 M HCl medium.
3
3.2. Test for free radical
in 0.3 M HCl medium were recorded between 600 and 700 nm (Fig. 1).
The nature of the spectra remained unchanged in the two cases. Howev-
er, the absorbance was found to decrease slightly in presence of 0.3 M
HCl.
The reaction mixture containing ascorbic acid, methylene blue and
HCl was allowed to stand for an hour in presence of 10% (m/v) acrylam-
ide solution. No polymerization or turbidity was observed even after the
addition of ethanol. This observation indicated the absence of any free
radical formation i.e., the reaction possibly did not involve a one-elec-
tron transfer process [31].
For both the solutions of methylene blue the absorbance maxima
were found to be at 665 nm (λmax). The absorbance values of MB solu-
tions (in water as well as in HCl medium) at the λmax at different con-
−
5
−5
centrations of MB (0.4 × 10
to 1.2 × 10
M) were recorded and
3.3. Effect of variation of [MB] and [AA]
Beer's law was found to be valid for both the cases within the
The oxidation reaction was studied at various initial concentrations
−
5
−5
+
of MB ranging from 0.6 × 10
[
to 1.0 × 10
M, keeping [H ] and
−
3
AA] constant at 0.1 M and 2 × 10 M respectively. The plots of lnA ver-
1
0
8
6
4
2
0
sus ‘time’ gave straight lines which were parallel to each other indicat-
ing that kobs values remain more or less constant with change in [MB].
Thus it indicates that the reaction is first order with respect to MB.
The average value of kobs was found to be approximately (1.35 ±
2
2
3
3
93 K
98 K
03 K
08 K
−
3 −1
0
.11) × 10
s
at 298 K.
−
4
The effect of change in AA concentration (5.0 × 10
to 5.0 ×
10− M) was studied at four different temperatures (293, 298, 303
3
+
−5
and 308 K), keeping [MB] and [H ] fixed at 1.0 × 10 M and 0.1 M re-
spectively. The pseudo-first order plots of lnA versus ‘time’ were drawn
for different initial [AA] and at a constant temperature (Fig. 2) and from
the slopes of the plots, the pseudo-first order rate constants (kobs) were
determined. The rate was found to increase with increasing [AA] (Table
1). The plots of kobs versus [AA] at different temperatures (Fig. 3) gave
straight lines passing through the origin indicating that the reaction
Table 2
The values of empirical constants at different temperatures.
0
1
2
3
4
5
1
1
k )
(M−2 s−1
Temperature (K)
k
0
(M− s−1
)
3
1
0 [AA] (M)
2
2
3
3
93
98
03
08
0.131
0.265
0.425
0.604
3.726
4.546
5.621
6.459
Fig. 3. Dependence of pseudo-first order rate constant (kobs) on ascorbic acid
concentration. Plot of kobs versus [AA] at four different temperatures. [MB] = 1.0 ×
−
5
+
10
M, [H ] = 0.2 M.

P.K. Sen et al. / Journal of Molecular Liquids 224 (2016) 472–479
475
6.0
5.5
5.0
4.5
4.0
3.5
3.0
2.5
2.0
KCl
KBr
KI
0.00
0.05
0.10
0.15
0.20
0.25
0.30
[
KX] (M)
Fig. 7. Piszkiewicz's model. Plot of log{(k
w
− k
Ψ
)/(k
Ψ
− k
m
)} versus log[S] at three
different temperatures.
Fig. 5. Dependence of pseudo-first order rate constant (kobs) on concentration of KCl, KBr
and KI at 298 K. Plot of kobs versus [KX]. [MB] = 1.0 × 10− M, [AA] = 2.0 × 10 M,
5
−3
+
[H
] = 0.2 M.
the rate increased almost linearly with a positive intercept on the ordi-
nate and a positive slope (within the acidity range studied). Therefore,
the rate law corroborating the above fact would be.
was first order with respect to ascorbic acid. Thus the reaction appeared
to follow the rate law
1
d½MBꢀ
1
d½MBꢀ
À
Â
þÃÁ
−
¼ kobs ¼ k2½AAꢀ
ð1Þ
−
¼ kobs ¼ k þ k
0
1
H
½AAꢀ
ð2Þ
½
MBꢀ dt
½
MBꢀ dt
where k
.24, 1.54 and 1.78 M
2
is the second order rate constant and the values are 1.01,
where k and k
0
1
are the empirical constants and their values at four dif-
−
1 −1
1
s
at 293, 298, 303 and 308 K respectively.
ferent temperatures are presented in Table 2. The above relation indi-
cated that the reaction would occur through an acid independent path
as well as an acid dependent path.
+
3
.4. Effect of variation of [H ]
+
The effect of change in [H ] (0.048–0.4 M) was studied at four differ-
ent temperatures by the addition of HCl, keeping [MB] and [AA] fixed at
1
3.5. Effect of addition of alkali halides (KX)
−
5
−3
.0 × 10 M and 2.0 × 10 M respectively. The ionic strength of the
Since one of the reactants, viz., MB was a positively charged species,
the influence of different halides such as KCl, KBr and KI on the reaction
rate was studied in order to see whether the effect was simply due to
the change of ionic strength or there was any specific effect of the halide
medium was maintained constant by adding appropriate amount of
KCl. The rate was found to increase with increasing acid concentration.
The plot of kobs versus [H ] (Fig. 4) at each temperature showed that
+
3.0
2.5
2.0
1.5
1.0
0.5
0.0
298 K
303 K
308 K
-4.8
-4.4
-4.0
-3.6
-3.2
-2.8
log [SDS]
Ψ Ψ
Fig. 6. Dependence of observed rate constant (k ) on SDS concentration at three different temperatures. Plot of k
versus log [SDS]. [MB] = 1.0 × 10− M, [AA] = 2.0 × 10−3 M, [H+] = 0.1
5
M.

476
P.K. Sen et al. / Journal of Molecular Liquids 224 (2016) 472–479
k
procedure was applied using Olson-Simonson model [32]. However,
+
MB + H A
+
MB + A
H
2
2
the fitting was not found to be reasonably good. Hence further attempt
was made to see whether Piszkiewicz co-operativity model [30] could
be fitted to explain this specific effect in the pre-micellar region. There-
slow
Scheme 2. Uncatalyzed reaction path.
fore, a plot of log [(k
w
ψ ψ m
− k ) / (k − k )] versus log[S] (Fig. 7) was drawn
at each temperature. Here k
w
is the pseudo-first order rate constant in
K
+
MB + H
+
2
+
the absence of SDS and ‘S′ represents the surfactant monomer. It was
found that the plot at each temperature was a straight line with a posi-
tive slope and a positive intercept (Fig. 7).
HMB
fast
/
k
4. Discussion
2
HMB + H A
+
+
MB + H + A
+
H
2
2
slow
The reaction involving oxidation of ascorbic acid by methylene blue
was studied in the HCl medium in the concentration range ((0.048–0.4)
M). Ascorbic Acid is known to be a weak acid and undergoes the follow-
ing ionization equilibria [26]:
Scheme 3. H+-catalyzed reaction path.
ꢀ
−5at 298 Kꢁ
H2A⇌HA− þ Hþ K1 ¼ 5:01 ꢁ 10
ð3Þ
ð4Þ
+
ions on the reactant species, MB . The reaction was studied at various
initial concentrations of KCl, KBr and KI (0.0–0.3 M), keeping [MB],
þ Hþꢀ
K2 ¼ 1:51 ꢁ 10−12at 298 K
ꢁ
+
−5
−3
−
2−
[
AA] and [H ] fixed at 1.0 × 10 M, 2.0 × 10 M and 0.1 M respective-
HA ⇌A
ly. It was found that the rate increased slowly with increasing salt con-
centration in the case of KCl and KBr and the effects were very much
similar. However, the behavior was slightly different with KI. In case
of KI, the increase was slow initially, but at higher concentrations the in-
crease of rate was much more rapid (Fig. 5). At the same ionic concen-
In the present investigation, the acid concentration was compara-
+
−
3
tively high. Thus at [H ] = 0.1 M, [H
therefore, evident that under the condition of kinetic investigations,
ascorbic acid remains almost fully in the undissociated form H A,
while the concentration of HA and A are negligible.
Since the reaction showed a linear dependence on [H ], it may be
suggested that the reaction occurs between undissociated ascorbic
2
A]:[HA ] ≈ 2.0 × 10 :1. It is,
2
−
−
−
tration, the effect of I was much greater than that of Cl or Br . This
indicated that possibly I had a specific interaction effect on the reac-
−
2−
−
+
tion rate [15].
2
acid (H A) and unprotonated as well as protonated methylene blue
+
2+
(
MB and HMB
respectively). As evident from the empirical rate
3
.6. Effect of variation of surfactant concentration
law Eq. (2) the reaction involves two parallel paths - one uncatalyzed
path and the other H - catalyzed path, the reaction may be assumed
to take place according to the following schemes where H
dized to dehydroascorbic acid (A) and methylene blue is reduced to
leucomethylene blue (H
According to the above Schemes 2 and 3, the rate equation comes
out to be
+
Since the oxidant methylene blue was a cationic substrate, an at-
2
A gets oxi-
tempt was made to investigate whether there was any effect of an an-
ionic surfactant SDS monomers in the pre-micellar region through co-
operative binding between the two oppositely charged species. The re-
action between methylene blue and ascorbic acid was studied in the
presence of SDS at concentrations (0.24–10) × 10 M, lower than its
CMC (8.0 × 10
+
2
MB ):
−
4
−
3
M) at three different temperatures keeping [MB],
Â
Ã
0
þ
+
−5
−3
d½MBꢀ k þ k K H
[AA] and [H ] fixed at 1.0 × 10 M, 2.0 × 10 M and 0.1 M respective-
−
¼
Â
à ½AAꢀ½MBꢀ
ð5Þ
þ
−
5
dt
1 þ K H
ly. The concentration of SDS was varied from 2.4 × 10
to 1.0 ×
−
3
1
0
M. It was observed that in the pre-micellar region, there was an
initial rapid decrease of pseudo-first order rate constant (k ) which ul-
timately reached a low limiting value (k ) at higher SDS concentration
but below its CMC. Plots of k versus log [SDS] at three different temper-
or
ψ
m
Â
Ã
0
þ
1
d½MBꢀ
k þ k K H
ψ
−
¼ kobs
¼
Â
à ½AAꢀ
ð6Þ
þ
½
MBꢀ dt
1 þ K H
atures are presented in Fig. 6. It may be noted that the concentration
range of SDS was too small to give an effect of changing ionic strength.
Therefore, the decrease in rate must be owing to some specific effect
of the anionic surfactant on the cationic substrate MB. To explain this
specific effect in the pre-micellar region, a non-linear curve fitting
+
Methylene blue exists as a positively charged species MB and the
positive charge on the sulfur atom is delocalized through conjugation
involving nitrogen atoms (Scheme 4).
Scheme 4. Delocalization of positive charge through conjugation.

P.K. Sen et al. / Journal of Molecular Liquids 224 (2016) 472–479
477
-
-
-
-
3.8
4.0
4.2
4.4
versus [AA] (Fig. 3) at four different temperatures are straight lines pass-
(
b) Catalyzed path
ing through the origin, thus corroborating the rate law Eq. (7). Also the
+
linear plots of kobs versus [H ] at different temperatures (Fig. 4) having
positive intercepts and positive slopes are in agreement with Eq. (7).
It may be noted that the empirical constant k
rate constant (k) for the slowest step in the uncatalyzed path whereas
is a composite rate constant involving the rate constant (k') for the
slowest step and the equilibrium constant (K) for the protonation of
0
is identical with the
k
1
3.24
3.28
3.32
3.36
3.40
3
-1
1
0 / T (K )
+
0 1
MB in the catalyzed path. Plots of ln(k /T) versus 1/T and ln(k /T) ver-
≠
sus 1/T (Fig. 8) were drawn to find out the enthalpies of activation (ΔH )
≠
and entropies of activation (ΔS ) for both the uncatalyzed path and the
-6.4
-6.8
-7.2
-7.6
≠
(
a) Uncatalyzed path
acid catalyzed path using the Eyring equation. The values of ΔH and
≠
ΔS for the slowest step of the uncatalyzed path were found to be
−
1
−1
−1
(
73.6 ± 7.7) kJ mol and (−10.9 ± 25) JK mol respectively. For
≠
≠
the acid catalyzed path, the values of ΔH and ΔS were (25.5 ± 1.5)
kJ mol and (−148 ± 5) JK mol respectively, which are the acti-
−1
−1
−1
vation parameters for the overall reaction. The activation parameters for
3
.24
3.28
3.32
3.36
3.40
the two paths could not be compared as the composite rate constant k
1
3
-1
1
0 / T (K )
is not a true rate constant for the rate determining step of the acid cat-
alyzed path.
Fig. 8. Dependence of k
/T) versus (1/T).
0
and k
1
on temperature. Plots of (a) ln (k
0
/T) versus (1/T) and (b) ln
In the rate determining steps (Schemes 2 and 3) the reaction takes
(k
1
+
2+
place between a positively charged species (MB or HMB ) and a
polar molecule (AA). The influence of KCl, KBr, and KI on the reaction
rate (Fig. 5) indicates that there is a specific interaction of the anions
Since nitrogen centers are partially positively charged, the possibility
of their further protonation will be low and hence K is expected to be
small. In the present experiment, [H ] was maintained in the range
0.048–0.4) M and therefore, K[H ] ≪ 1. Thus Eq. (6) becomes.
−
moderating the reaction rate. The influence is highest in case of I . On
+
the average, there is an increase in rate with increase in salt concentra-
tions. This may be accounted for by the fact that the transition state pro-
duced from the charged and polar reactants must be highly polar and
hence will be more stabilized by the presence of ion atmospheres than
the reactants themselves. Therefore, an increase in ionic strength will
favor the transition state and thereby increases the reaction rate [33,
+
(
1
d½MBꢀ
À
0
Â
þÃÁ
−
¼ kobs ¼ k þ k K H ½AAꢀ
ð7Þ
½
MBꢀ dt
−
Therefore, comparing with the empirical Eq. (2), we have,
34]. Of the three halide ions, I ion may have some specific effect on
−
the reaction rate in that the large I ion, being much more polarisable,
0
k0 ¼ kandk1 ¼ k K
ð8Þ
acting as a bridging ligand may facilitate the electron transfer between
+
2+
MB /HMB and AA molecule.
Thus the proposed scheme of reaction corroborates the empirical
rate law Eq. (2). Eq. (7) shows that the reaction is first order with re-
spect to each of methylene blue and ascorbic acid. The plots of kobs
Acceleration or inhibition of reaction rates often found in the pres-
ence of surfactants may occur owing to either the change in the reaction
environments or to the modification of the reaction scheme [30,35,36].
N
O
-O
O
O
(H C) N
S
N(CH3)2
3
2
Na+
S
+
SDS monomer
O
O
H
CH OH
OH
2
O
HO
H
N
O
O
H
OH
CH OH
OH
2
N
S
+ O
(
H C) N
3
S
N(CH3)2
H
2
(
H C) N
N(CH3)2
-O
3
2
O-
O
O
O
+
S O Na
Na+
S
O
O
Pre-micelle aggregate
Scheme 5. Representation of the overall reaction through premicellar aggregation.

478
P.K. Sen et al. / Journal of Molecular Liquids 224 (2016) 472–479
nS
+
MB
S MB
n
The k
ence of surfactant (i.e., constant value of rate constant at higher surfac-
tant concentration but less than its CMC value) in the plot of k versus
log [SDS] (Fig. 6). The linear plots of log{(k − k ) / (k − k )} versus
log [S] (Fig. 7) with positive slope and intercept at three different tem-
peratures (298, 303 and 308 K) are shown in Fig. 7. The values of co-
m
was taken as the limiting or minimum rate constant in the pres-
K
S
Ψ
w
Ψ
Ψ
m
k
m
S MB
n
Product
S
operativity index (n) and K (the dissociation constant of the pre-mi-
celle, S MB) at three different temperatures have been evaluated from
n
k
w
the slopes and intercepts of these plots (Fig. 7) and are shown in Table 3.
The small values of n (1.45 to 1.76) are in agreement with earlier ob-
servations [14–16,30]. Such values are much less than the aggregation
number of the normal micelles. Moreover, these pre-micelles are aggre-
gates between the substrate molecule and a small number of surfactant
monomers while the normal micelles are aggregates among a large
MB
Product
Scheme 6. Reaction steps according to Piszkiewicz's model.
In the presence of surfactants, partitioning of the reactants often occur
between the bulk aqueous phase and the micellar pseudo-phase
which may either increase or decrease the local concentration of the re-
actants, thereby resulting in an enhancement or diminution of the reac-
tion rate [36]. A change in the reaction rate in the presence of
surfactants may sometimes be accounted for by a change in the dielec-
tric constant or the pH of the medium at the micellar surface [37,38].
This is to be noted that the present oxidation reaction was per-
formed at SDS concentrations (0.24–10) × 10 M, which is sufficiently
lower than its normal CMC. Hence the formation of micelles is not pos-
sible at this low concentration and so the inhibition of reaction rate is
definitely not due to the effect of micellar aggregates. Moreover, the in-
hibition effect of SDS monomers (in the pre-micellar region) on the re-
action rate cannot be explained as an effect of increase of ionic strength,
because the change in ionic strength is small and the effect is opposite to
that obtained by addition of halides. Therefore, one plausible explana-
tion may be that a few anionic SDS monomers (having negatively
charged head groups) bind with the positively charged methylene
blue species (both protonated and unprotonated forms) through elec-
trostatic as well as hydrophobic interactions in a co-operative manner
thus making the approach of the ascorbic acid molecule more difficult
number of surfactant monomers only. The dissociation constant (K
of the pre-micelle is very small (Table 3). The small values of K were
also reported in a number of investigations [14–16]. The small K
S
)
S
S
value indicates that the dissociation of the pre-micellar complex back
to its free components (Scheme 6) is negligible. It is quite plausible
that before the formation of pre-micelle, the reactant MB and the SDS
monomer are in hydrated forms. But in the pre-micellar assembly
they interact more strongly by means of electrostatic as well as hydro-
phobic interactions which are possibly stronger than the hydration in-
teractions of the individual molecules of MB and SDS. This may
−
4
account for the small value of K
value of K does not always reflect the stronger pre-micellar binding be-
tween the reactant and the surfactant monomer. Apart from the K
value, the extent of binding also depends on the co-operativity index
(n) and the SDS concentration (Scheme 6). The ratio of [S MB]/[MB]
S
. However, it must be noted that the
S
S
n
has been evaluated to be 154.8, 81.4 and 63.5 (Table 3) at the tempera-
tures of 298 K, 303 K and 308 K respectively. The ratio indicates that the
binding in the pre-micelle is reasonably strong and most of the MB mol-
ecules remain associated in the pre-micelle. The above results show that
the extent of association decreases with increase in temperature. This is
possibly due to the fact that both MB and SDS remain in hydrated forms
in the aqueous medium. When they move towards the formation of the
pre-micelles, water structures are destroyed with an absorption of ener-
gy (ΔH positive) and at the same time MB and SDS monomers unite to-
gether through electrostatic and hydrophobic interactions with a
(
Scheme 5). As a result the overall reaction rate rapidly decreases.
When almost all the oxidant species become bound to the SDS mono-
mers, the rate reaches a limiting low value.
4
.1. Piszkiewicz's model
There are reports of different models [16,30,39] which can explain
Table 3
the inhibition of reaction rates in presence of surfactant in the pre-mi-
cellar region. In our present investigation, the rate constant values at dif-
ferent SDS concentrations at a constant temperature have been
analyzed using earlier reported Piszkiewicz's model [30] of pre-micellar
catalyzed reactions. In this model, a small number (n) of SDS monomers
−
3
Influence of surfactant concentration on the reaction rate. [AA] = 2.0 × 10 M, [MB] =
.0 × 10 M and [H+] = 0.1 M.
−5
1
Temperature 104 [SDS] 104
k
)
Cooperativity
index (n)
108
Ψ
−
1
(K)
(M)
0.24
(s
K
S
n
[S MB]/[MB]
298
303
308
20.4
16.6
7.85
1.65
1.52
2.16
1.76
1.5
25.1
21.5
13.5
4.59
5.15
3.5
2.82
2.35
28.3
23.3
11.2
4.94
5.79
4.30
3.24
2.53
1.76
1.71
1.45
3.39 154.8
9.11 81.4
70.3 63.5
(
(
S) aggregates or associates with the reactant (MB) to form pre-micelles
MB). These pre-micellar aggregates react at a different rate along
0
0
2
.56
.96
.0
S
n
with the normal reaction occurring in the aqueous phase (Scheme 6).
Here, K is the dissociation constant of the pre-micelle (S MB) [30]
back to its free components, k is the rate constant in the aqueous
phase and k is the limiting value of the rate constant in the presence
of SDS. According to the Scheme 6, the rate law expressing the pseu-
S
n
4.0
6
8
1
.0
.0
0.0
w
m
0.24
0.56
0.96
ψ
do-first order rate constant (k ) as a function of SDS concentration be-
comes,
2.0
n
4.0
6.0
km½Sꢀ þ kwKS
kψ ¼
ð9Þ
n
KS þ ½Sꢀ
8
1
.0
0.0
0.24
Rearrangement of the above equation in a suitable form gives,
0
0
.56
.96
kw−kψ
kψ−km
log
¼ n log½Sꢀ− logKS
ð10Þ
2.0
4.0
6
8
1
.0
.0
0.0
w Ψ Ψ
The above relation demands a linear plot of log{(k − k ) / (k −
k
m
)} versus log[S] with the slope value of n (the co-operativity index).

P.K. Sen et al. / Journal of Molecular Liquids 224 (2016) 472–479
479
release of energy (ΔH negative). The latter is greater than the former so
that the net enthalpy change during the pre-micelle formation is nega-
tive or, in other words, the dissociation of the pre-micelle to its compo-
nents, namely, MB and SDS, involves an absorption of energy. That is
why the dissociation of the premicelle is favored at higher temperature
program (Phase II) to the Department of Chemistry, Jadavpur University
are gratefully acknowledged.
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Acknowledgement
Financial assistances from DST, Govt. of India, through the PURSE
program (Phase II) and from UGC, New Delhi, India, through the CAS