Effect of microheterogeneous environments of CTAB, Triton X-100, and Tween 20 on the oxidative degradation of d-fructose by nanoparticles of MnO2

Article abstract of DOI:10.1002/kin.21239

The kinetics of the oxidative degradation of d-fructose by nanoparticles of MnO₂ has been studied in dilute sulfuric acid medium and also in the presence of surfactants of cetyl trimethyl ammonium bromide (CTAB), Triton X-100 (TX-100), and Tween 20. Amorphous nanoparticles of MnO₂ in the form of spherical particulates of size 50–200?nm, as detected by a transmission electron microscope, have been found to exist, supported on two-dimensional gum acacia sheets. The reaction is first order in MnO₂ but complex order with respect to fructose and H⁺. The reaction is inhibited due to adsorption of reaction products on the surface of MnO₂ nanoparticles. The reaction takes place through an intermediate complex formation between β-d-fructopyranose and protonated MnO₂. A one-step two-electron transfer reaction ultimately leads to the formation of an aldonic acid and formic acid. The entropy of activation plays the key role for the reaction in the absence of surfactants. In the surfactant-mediated reaction, partitioning of both the reactants takes place between the aqueous and micellar pseudophases and reaction occurs following Berezin's model. Binding of fructose with the surfactants in the Stern/palisade layer takes place through the ion–dipole interaction and H-bonding while protonated MnO₂ remains at the outer side of the Stern/palisade layer within the micelle. Both the enthalpy and entropy changes associated with the fructose–water interaction, fructose–micelle interaction, and micelle–water interaction finally control the fructose–micelle binding.

Full text of DOI:10.1002/kin.21239

Received: 14 May 2018
DOI: 10.1002/kin.21239
Revised: 1 November 2018
Accepted: 28 November 2018
A R T I C L E
Effect of microheterogeneous environments of CTAB, Triton
X-100, and Tween 20 on the oxidative degradation of D-fructose
by nanoparticles of MnO2
Jayanta K. Midya1
Dinesh C. Ghosh1
Biswajit Pal2
Pratik K. Sen1
1Department of Chemistry, Jadavpur
University, Kolkata, India
Abstract
The kinetics of the oxidative degradation of D-fructose by nanoparticles of MnO2 has
been studied in dilute sulfuric acid medium and also in the presence of surfactants
of cetyl trimethyl ammonium bromide (CTAB), Triton X-100 (TX-100), and Tween
20. Amorphous nanoparticles of MnO2 in the form of spherical particulates of size
50–200 nm, as detected by a transmission electron microscope, have been found to
exist, supported on two-dimensional gum acacia sheets. The reaction is first order in
2Department of Chemistry, St. Paul's C. M.
College, Kolkata, India
Correspondence
Pratik K. Sen, Department ofChemistry,
Jadavpur University, Kolkata700032, India.
Email: senpkju@gmail.com
+
MnO2 but complex order with respect to fructose and H . The reaction is inhibited due
to adsorption of reaction products on the surface of MnO2 nanoparticles. The reaction
takes place through an intermediate complex formation between ꢀ-D-fructopyranose
and protonated MnO2. A one-step two-electron transfer reaction ultimately leads to the
formation of an aldonic acid and formic acid. The entropy of activation plays the key
role for the reaction in the absence of surfactants. In the surfactant-mediated reaction,
partitioning of both the reactants takes place between the aqueous and micellar pseu-
dophases and reaction occurs following Berezin's model. Binding of fructose with the
surfactants in the Stern/palisade layer takes place through the ion–dipole interaction
and H-bonding while protonated MnO2 remains at the outer side of the Stern/palisade
layer within the micelle. Both the enthalpy and entropy changes associated with the
fructose–water interaction, fructose–micelle interaction, and micelle–water interac-
tion finally control the fructose–micelle binding.
K E Y W O R D S
Berezin model, fructose, kinetics, MnO2 nanoparticles, surfactant effects
properties owing to its having a larger specific surface. A num-
ber of kinetic studies14–16 on the oxidation of a few organic
1
INTRODUCTION
compounds by colloidal solution of MnO2 have already been
reported.
Over the past few decades, a great deal of work has been
undertaken to explore the mechanistic pathways for the oxi-
dation of both inorganic1–3 and organic4–11 substrates by per-
manganate in acidic and alkaline media. However, the +4 oxi-
dation state of manganese has received less attention until
now, due to its solubility problem There are reports12,13 on
the preparation of colloidal MnO2 by the reduction of KMnO4
with sodium thiosulfate under neutral conditions. Such form
of colloidal manganese dioxide may show better oxidizing
Carbohydrates are the structural backbone of the DNA,
RNA, and nucleic acids,17 and they serve as store and source
of energy in biological systems and play important roles in dif-
ferent nutrition processes.18 Monosaccharides are involved in
carbohydrate metabolism, and the mechanisms of their oxida-
tive degradations are of immense importance from biologi-
cal point of view.19–21 There are a number of reports on the
Int J Chem Kinet. 2018;1–12.
wileyonlinelibrary.com/journal/kin
© 2018 Wiley Periodicals, Inc.
1

2
MIDYA ET AL.
oxidative degradation of sugars, methylated sugars, and sugar
phosphates by different oxidizing agents.22–25
There is growing interest in recent times on the influence
of different surfactants on the kinetics and mechanism of dif-
ferent reactions.26–33 Reaction rates are often modified owing
to hydrophobic and/or electrostatic interactions between the
reactant species and surfactant monomers or micelles. The
oxidation of D-glucose by MnO2 nanoparticles34 in the
absence and presence of different surfactants has very recently
been studied. D-Glucose is an aldohexose (–CHO group in
the C-1 position), whereas D-fructose is a ketohexose (–C=O
group in the C-2 position). Thus the reaction centers as well as
their environments are different in the two substrates. Hence,
for a continuation of this study, an attempt has been made to
investigate the detailed kinetics of the oxidative degradation
of D-fructose by the nanoparticles of MnO2 in the absence
and presence of a cationic surfactant cetyl trimethyl ammo-
nium bromide (CTAB) and two nonionic surfactants Triton
X-100 (TX-100) and Tween 20. An attempt has been made
to account for the differences in the rate and thermodynamic
parameters for the two related monosaccharides on the basis
of the proposed mechanism
F I G U R E 1 TEM image of nanoparticles of MnO2 embedded in
the gum acacia cluster
2
MATERIALS AND METHODS
2.1 Materials
D(-)fructose (extrapure; SRL, Mumbai, India), potassium per-
manganate (extrapure AR; SRL, Mumbai, India), manganous
sulfate monohydrate (extrapure AR; SRL, Mumbai, India),
surfactants CTAB (extrapure AR; SRL, India), TX-100 (SRL,
Mumbai, India), Tween 20 or polysorbate 20 (extrapure; SRL,
Mumbai, India), and gum acacia (Loba Chemie, Mumbai,
India) were used without purification. Water from Millipore
synergy was used in all the experiments.
F I G U R E 2 TEM image of a cluster of nanoparticles of MnO2 in
the presence of fructose
2.2 Preparation of nanoparticles of MnO2
and their characterization
microscope (TEM) (model: Tecnai G2 30ST, FEI, Hillsboro,
OR, USA). The solution appeared to be homogeneous and
transparent, although it contained aggregates of nanoparticles
of size 50–200 nm. A bright-field TEM image (Figure 1)
of such a cluster of nanoparticles of MnO2 in the absence
of fructose shows that the cluster is almost spherical in
nature34 and such nanoparticles were adhered on the surface
of two-dimensional gum acacia sheets attaining the shape
of a crumbled paper ball. The amorphous nature of the
nanoparticles was confirmed by the absence of diffraction
spots in the SAED (Selected Area Electron Diffraction)
pattern. In the presence of fructose, the cluster gives a dense
appearance (Figure 2). Therefore, it may be concluded that
the oxidation of fructose by MnO2 possibly takes place on
the active surface of the nanoparticles of MnO2 adhered to
the surface of gum acacia sheets.
The colloidal solution of MnO2 was prepared via a com-
proportionation reaction between KMnO4 and MnSO4 as
described in an earlier communication.34 The absorption
spectra of the colloidal solution of MnO2 showed a large
absorption band covering the UV–visible region of the
spectrum with a flat maximum in the region 336–337 nm.34
The kinetics of the reaction was followed by noting the
absorbance at 390 nm.34–36. The solution was found to obey
Beer's law at this wavelength and in the studied concentration
−4
range of MnO2 ((0.5–3.0) × 10 mol L−1). All other chem-
ical species in the reaction mixture showed no absorption at
this wavelength. The nature of the reactant particles in the
oxidant solution alone as well as in the reaction mixture was
then characterized with the help of a transmission electron

MIDYA ET AL.
3
ride solution. Appearance of a bright yellow color indicated
the formation of an aldonic acid in the oxidation product.11
The presence of carboxylate anion in the reaction product was
also reported in the electro-oxidation of fructose on chemi-
cally prepared MnO2 on Pt support.37 The presence of formic
acid in the reaction product was confirmed by the following
experiment. Three milliliters of the reaction mixture (after
completion of the reaction) was mixed with 1.5 mL of 10%
HgCl2 solution and 2 mL of acetic acid–sodium acetate buffer
(pH ∼ 4.2) on a watch glass and evaporated to dryness on a
water bath. The white residue on the watch glass, after cool-
ing, was then treated with a drop of ammonia solution when
it turned black.38 To see whether any intermediate lactone
was formed, the reaction mixture was treated with alkaline
solution of hydroxylamine and then 2% FeCl3 and 1% HCl
solution were added, but no blue coloration was obtained.23
Possibly the intermediate lactone, which was formed in an
intermediate step, was hydrolyzed to aldonic acid in a fast
step.37
F I G U R E 3 Surface tension measurements of CTAB solution in
the absence and presence of fructose and H2SO4 . Plots of surface
−1
tension versus log [CTAB]: A, water; B, 0.1 mol L fructose ; C,
−1
−1
mixed solution (0.1 mol L fructose and 0.005 mol L H2SO4)
[Color figure can be viewed at wileyonlinelibrary.com]
3
RESULTS AND DISCUSSION
2.3 Kinetic procedure
3.1 CMC of different surfactants
Appropriate quantities of H2SO4, water, and fructose solu-
tion (and surfactant, if any) were taken in a glass cuvette
(1.0 cm path length) and mixed thoroughly. It was then placed
in the Peltier controlled thermostated cell holder of a UV–
visible Shimadzu 1800 spectrophotometer. After temperature
equilibration, a requisite volume of colloidal MnO2 solution
was added to start the reaction. The decreasing absorbance
of the reaction mixture at 390 nm with time indicated the
advancement of the reaction. Reagents other than MnO2 did
With the help of tensiometric experiments, the CMC values
of cationic surfactant CTAB and the nonionic surfactants TX-
100 and Tween 20 in aqueous medium at 298 K were deter-
mined as 1.08 × 10−3, 0.199 × 10−3, and 0.050 × 10 mol
−3
L
−1, respectively. One representative plot for CTAB is shown
in Figure 3. With CTAB, in the presence of fructose solution
(0.1 mol L−1), the nature of the plot before CMC is deviated
from linearity. Possibly CTAB monomers tend to aggregate
with fructose molecules beyond a critical aggregation con-
′
not absorb at this wavelength. The rate constants (k and k )
centration (CAC) of 0.271 × 10 mol L−1, and this became
−3
of the reaction were determined from the kinetic plots as
described in the Results and Discussion section.
saturated or complete at the [CTAB] of 0.610 × 10−3 mol L
−1
known as the saturation concentration (C ). Beyond C , the
surfactant monomers began to form normaSl micelles whiSle the
2.4 Tensiometric measurements of critical
micelle concentration
−3
−1
CMC value increased to 1.156 × 10 mol L compared to
the value in pure water (Figure 3). This type of aggregation
between substrate and surfactant monomer possibly indicated
a weak interaction between fructose molecules and CTAB
Surface tension of the three different surfactants at different
concentrations was measured with the help of a Jencon (India)
Du Nuoy tensiometer with a platinum ring for the measure-
ment. The critical micelle concentration (CMC) for different
surfactants in pure water as well as in the presence of fructose
and H2SO4 was determined from the point of intersection in
the plots of “surface tension” versus log[surfactant] values.
Figure 3 shows three representative plots for CTAB.
monomers.34 In the presence of a mixed solution (0.1 mol L
−1
−1
fructose and 0.005 mol L H2SO4), the plot again showed
some curvature but the CMC value again decreased to 0.955
× 10−3 mol L−1 (Figure 3), which was slightly lower than that
of CTAB in pure water. For TX-100, the nature of the ten-
−1
siometric plots in the presence of 0.1 mol L fructose and
in the presence of a mixed solution remained unchanged, but
2.5 Product analysis
−3
−1
the CMC value first increased to 0.212 × 10 mol L and
then it again decreased to 0.169 × 10−3 mol L−1, respectively.
For Tween 20, the nature of the tensiometric plots remained
unchanged, but the CMC value first decreased to 0.034 ×
10−3 mol L−1 in the presence of 0.1 mol L−1 fructose and then
After completion of the reaction of fructose with MnO2, the
reaction mixture was acidified with 4 N H2SO4 and then
treated with the violet colored solution obtained by the addi-
tion of a few drops of phenol to a few milliliters of ferric chlo-

4
MIDYA ET AL.
constants involved in the reaction steps as well as the concen-
trations of the reactants and k ′ is a rate constant related to the
adsorption of the products and the term (k ′x) gives the rate of
inhibition.
As explained in an earlier communication,34 the relation
between At/A and time will be
0
ꢍ′
ꢍ + ꢍ′ ꢍ + ꢍ′
ꢍ
ꢁꢂ
ꢁ0
′
ꢂ
ꢅ− ꢍ+ꢍ
(
(3)
)
=
+
Therefore, a plot of At /A versus time was expected to pro-
0
duce an exponential graph and the experimental results were
analyzed with the help of a nonlinear fitting and the values of
k and k ′ were evaluated.
3.3 Dependence of the reaction rate on
[MnO2] and [D-fructose]
F I G U R E 4 Plots of At/A versus time for the oxidation of
0
+
fructose by MnO2 at 293 K. [H ] = 0.01 mol L−1, [MnO2] = 1.0 ×
−3
10 mol L−1, 102 [Fruc] (mol L−1): a, 2.5 ; b, 5.0; c, 7.5; d, 10.0; e,
The dependence of the rate constant k on the MnO2 concen-
12.5; f, 17.5; g, 25.0 [Color figure can be viewed at
wileyonlinelibrary.com]
tration was determined by varying the initial [MnO2] ((0.5–
−4
+
3.0) × 10 mol L−1) at 298 K while [D-fructose] and [H ]
were kept constant at 0.1 and 0.01 mol L−1, respectively. The
rate constant (103k s−1) were found to be 5.53, 3.91, 3.64,
3.42, 3.15, and 2.64 at MnO2 concentrations of 0.5, 1.0, 1.5,
further decreased to 0.031 × 10−3 mol L−1 in the presence of
a mixed solution.
−4
2.0, 2.5, and 3.0 × 10 mol L−1, respectively. Thus increas-
ing [MnO2] showed an inhibiting effect on the reaction rate.
The flocculation of the nanoparticles at higher concentrations
might decrease the effective surface area, thereby resulting in
a decrease in the value of k.36
3.2 Evaluation of the rate constant from
time-dependent absorption studies
The experiments were carried out under the pseudo–first-
order condition such that [MnO2] ≪ [fructose]. The plots of
log At versus time at different initial fructose concentrations
were nonlinear. Instead, a plot of At/A (where At and A rep-
+
At a fixed temperature and constant [MnO2] and [H ],
an increase in the concentration of fructose ((2.5–25.0) ×
10−2 mol L−1) increased both the values of k and k ′ (Table 1).
Such studies were carried out at different acidities (0.01–
0.04 mol L−1) and different temperatures (293–308 K). It may
be noted here that k is directly related to the actual reaction
steps associated with the oxidation of fructose. Hence, k will
be dependent on the concentration of fructose. At a particular
acidity and temperature, a plot of k versus [Fruc] was found
to be nonlinear. On the other hand, a linear plot with a posi-
tive slope and positive intercept was obtained when k−1 was
0
0
resent the absorbances at time t and at the start, respectively)
versus time (Figure 4) was found to obey the following empir-
ical equation:
ꢁꢂ∕ꢁ0 = ꢃ + ꢄ ꢅ−ꢆꢂ
(1)
The variable parameters a, b, and z are dependent on the con-
centrations of reactants and different rate constants and equi-
librium constants. It may be suggested that fructose molecules
get adsorbed on the surface of the nanoparticles prior to oxi-
dation by the oxidant. The adsorbed products may hinder the
oxidation by the oxidant before getting completely desorbed.
Based on this idea ,the rate of the reaction may be expressed
in the form
−1
plotted against [Fruc] (Figure 5). At a particular tempera-
ture, the slopes of such double reciprocal plots at four differ-
+ −1
ent acidities were plotted against [H ] when a straight line
was obtained (Figure 6). The values of k , K , and K (Equa-
1
2
tion 10) at four different temperatures wdere determined from
such straight-line plots (Figure 6) at those temperatures, and
such values were found to increase with an increase in tem-
perature.
ꢈ (ꢉ − ꢊ)
ꢇ = −
= ꢇꢋ − ꢇꢌ = ꢍ (ꢉ − ꢊ) − ꢍ′ꢊ
(2)
ꢈꢂ
where vf and vi are the rates of forward reaction and inhi-
bition, respectively, c is the initial concentration of MnO2
nanoparticles, and x gives the concentration of MnO2 that has
been reduced at time t. Here, k is a pseudo–first-order rate
constant that involves different rate constants and equilibrium
3.4 Dependence of the reaction rate on [H+]
+
A number of experiments for the variation of [H ] (0.01–
0.04 mol L−1) were performed where at a constant tempera-
+
ture, [MnO2] and [H ], the fructose concentration was altered

MIDYA ET AL.
5
TA B L E 1 Dependence of reaction rate constants on the fructose concentration at different temperatures
293 K
103k (s−1
2.14
298 K
103k (s−1
2.98
303 K
103k (s−1
3.61
308 K
103k (s−1
4.41
102[Fructose] (mol L−1
)
)
103k′ (s−1
1.28
)
)
103k′ (s−1
2.01
)
)
103k′ (s−1
2.27
)
)
103k′ (s−1
2.4
)
2.5
5.0
3.22
1.73
4.61
2.23
5.48
2.46
7.03
4.1
7.5
3.93
1.74
5.33
2.68
6.10
2.69
7.72
4.91
5.29
6.32
6.52
7.01
10.0
12.5
17.5
4.49
1.95
6.44
3.34
6.82
4.01
8.51
5.12
2.33
7.04
3.62
8.02
4.36
9.81
5.62
2.63
8.05
3.86
9.24
4.91
11.30
12.91
25.0
7.11
3.21
8.72
4.52
10.47
5.32
−4
+
−1
[MnO2] = 1.0 × 10 mol L−1; [H ] = 0.04 mol L
.
F I G U R E 7 Variation of rate constant (k) with acid concentration.
F I G U R E 5 Variation of rate constant (k) with the fructose
+
−1
Plots of k -1 versus [H ] at different temperatures. [MnO2] = 1.0 ×
−1
concentration. Plots of k −1 versus [Fruc] at different temperatures.
−4
−1
−1
10 mol L and [Fruc] = 0.1 mol L [Color figure can be viewed at
+
−1
−4
−1
[H ] = 0.02 mol L and [MnO2] = 1.0 × 10 mol L [Color figure
wileyonlinelibrary.com]
can be viewed at wileyonlinelibrary.com]
SCHEME 1 Oxidative degradation of fructose occurring through
different reaction steps
−2
+
in the range (2.5–25.0) × 10 mol L−1. An increase in [H ]
enhanced the reaction rate. At a particular temperature and a
fixed [Fruc], a plot of k –1 versus [H ] produced a straight
+ −1
+
F I G U R E 6 Dependence of slope on [H ]. Plots of slope ({1+
line (Figure 7). Different rate constants and equilibrium con-
stants (Scheme 1) were then evaluated from the slopes and
intercepts of these plots.
+
+
+
−1
K [H ]}/{k K K [H ]}) versus [H ] at four different temperatures
1
d
2
[Color figure ca1n be viewed at wileyonlinelibrary.com]

6
MIDYA ET AL.
TA B L E 2 Equilibrium constants and activation parameters for the oxidation of glucose and fructose
Substrate
Glucose
K1 (298 K)
108
K2 (298 K)
15.6
ꢀH≠ (kJ mol−1
)
ꢀS≠ (J K−1 mol−1
–223 18
)
ꢀG≠ (kJ mol−1
87.7 3.2
)
Reference
34
20.7 2.2
Fructose
141
37.9
17
3
–230 18
85.5 2.9
Present work
3.5 Dependence on temperature
An increase in temperature increased the values of k , K , and
d
1
K . The enthalpy of activation (ΔH#) and the entropy of acti-
2
vation (ΔS#) for the rate-limiting step (Scheme 1) were cal-
culated from the linear plot of ln(k /T) versus 1/T using the
d
Eyring equation:
(
)
[
(
)
]
ΔS# Δꢐ#
ꢍd
ꢎ
ꢍB
ℎ
ln
= ln
+
(4)
ꢏ
ꢏꢎ
where k and h are Boltzmann constant and Planck constant,
B
respectively, and the values are presented in Table 2. Also the
0
0
0
enthalpy changes (ΔH , ΔH ) and entropy changes (ΔS ,
1
2
1
0
ΔS ) for the protonation of MnO2 and for the intermediate
2
complex formation step (Scheme 1), respectively, were deter-
mined from the linear plots of ln K versus 1/T and ln K ver-
1
2
F I G U R E 8 Variation of pseudo–first-order rate constant (kΨ)
sus 1/T using the following relations:
with [CTAB] and temperature. [H ] = 0.001 mol L−1, [MnO2] = 1.0 ×
+
−4
−1
10 mol L−1, and [Fruc] = 0.1 mol L [Color figure can be viewed at
Δꢐ0 1
ln ꢑ =
+ constant
(5)
wileyonlinelibrary.com]
ꢏ
ꢎ
Δꢐ0 − Δꢒ0
Δꢒ0 = −ꢏꢎ ln ꢑ and Δꢓ0 =
(6)
alytic part just before the CMC may be due to the reorganiza-
tion of the surfactant monomers to form micellar aggregates
and consequent interaction with the reactant molecules.
ꢎ
The values of ΔH 0 and ΔH 0 are 13 4 and 28 3 kJ mol
−1
1
2
and of ΔS 0 and ΔS 0 are 86 15 and 123 9 J K−1 mol
,
−1
2
respective1ly.
3.7 Mechanism of oxidaion in the absence of
surfactants
3.6 Reaction in the presence of surfactants
+
The reaction rate increases with increasing [H ], indicating
that either fructose or MnO2 may be protonated. But at this
The reaction was further studied in the presence of two non-
ionic surfactants TX-100 and Tween 20 and a cationic surfac-
tant CTAB. It may be noted that the stability of the colloidal
solution of MnO2 increases with decreasing acidity. More-
over, addition of electrolytes has been found to affect the size
and shape of the micelles as well as the CMC. For these rea-
sons, the kinetic study in the presence of surfactants was car-
ried out at comparatively much lower acidities.34,36,39 In case
of CTAB, in the premicellar region there was first a negligi-
ble increase and then the rate constant (kꢔ ) decreased until
near the CMC there appeared again a hump and then the rate
constant decreased gradually in the postmicellar region (Fig-
ure 8). Thus it appears that inhibition occurs both before and
after CMC for CTAB. In case of each of TX-100 and Tween
20, there was at first a decrease in the value of kꢔ in the pre-
micellar region, showing small inhibition, and then the value
of kꢔ increased passing through a maximum at the CMC and
then falls gradually, indicating clear inhibition in the postmi-
cellar region. For both the nonionic surfactants, the small cat-
+
low concentration of H , protonation of fructose is expected
to be negligible.10 Thus it is quite plausible that in a fast
preequilibrium, the protonation of MnO2 takes place. Again,
on the basis of the Lineweaver–Burk type linear plot of 1/k
against 1/[Fruc], it may be proposed that fructose and proto-
nated MnO2 react in a fast step to form an intermediate com-
plex followed by its decomposition in a slow rate-determining
step34 to form products (Scheme 1).
Based on the above scheme of reaction, the forward reac-
tion rate for the reduction of MnO2 is given by
[
]
ꢍdꢑ1ꢑ2 H+ [Fruc]
[
]
ꢇꢋ =
MnO2
(7)
1 + ꢑ1 [H+] + ꢑ1ꢑ2 [H+] [Fruc]
Comparing Equations 2 and 7, the pseudo–first-order rate
constant is as follows:
[
]
ꢍdꢑ1ꢑ2 H+ [Fruc]
1 + ꢑ1 [H+] + ꢑ1ꢑ2 [H+] [Fruc]
ꢍ =
(8)

MIDYA ET AL.
7
SCHEME 2 Mechanistic path of the redox reaction showing different steps
The relation may be rearranged to
D-Fructose is known to exist in an equilibrium mixture of
pyranoid (58%) and furanoid (42%) forms.40 Of these two
forms, only the pyranose form is known to participate in the
oxidation process.41 Again between the two conformers, the
ꢀ-D-fructopyranose is reported to be much more reactive than
the ꢕ-anomer.41 Therefore, it may be suggested that the equa-
torial –OH of ꢀ-D-fructopyranose combines with the proto-
nated form of MnO2 to form an intermediate complex, which
disproportionates in a rate-limiting two-electron transfer step
(Scheme 2) to produce a lactone and a hydroxymethyl carbo-
cation. The lactone is hydrolyzed in a fast step to the aldonic
acid (five membered). The unstable carbocation, on reaction
with water, is fast converted to formaldehyde, which is oxi-
dized to formic acid.42,43
[
]
1 + ꢑ1 H+
1
ꢍ
1
=
+
(9)
ꢍd ꢍdꢑ1ꢑ2 [H+] [Fruc]
Thus Equation 9 corroborates the double reciprocal linear
plots of k–1 versus [Fruc]−1 at different temperatures and dif-
ferent acidities (Figure 5). The values of k at different tem-
peratures were evaluated from the intercedpts of such plots.
+
From four sets of similar experiments at four different [H ],
+
+
the slopes ({1+ K [H ]}/{k K K [H ]}) of such straight
1
d
1 2
+
lines were evaluated and plotted against 1/ [H ] (Figure 6).
As expected from the relation, a straight-line plot resulted at
each temperature. From the slope and intercept of these plots
(Figure 6), the values of K and K at four different tempera-
1
2
A similar type of two-electron transfer oxidation of fructose
was observed in the reaction with N-bromoacetamide42 in an
acid medium. The fact that one-electron transfer step did not
intervene in the present reaction is further supported by the
absence of polymerization within the reaction mixture in the
presence of 15%(v/v) acrylonitrile.23,32 It may be noted that
Mn(III) shows an absorption peak at 470 nm.36 Keeping this
in mind, time-resolved spectra of the reaction mixture were
taken in the range 350–550 nm. But no increase in absorbance
at 470 nm was observed during the reaction, indicating that
Mn(III) was not formed as an intermediate. This further con-
firmed that one-electron transfer reaction did not take place.
The enthalpy of activation for the rate-determining step was
found to be low, which may be ascribed to the enhanced oxi-
dizing power of MnO2 owing to its nanoscale structures. Such
low values of the enthalpy of activation have earlier been
reported.34,44 On the contrary, for the present reaction the
entropy of activation was found to be highly negative (–230
tures were determined by using the earlier obtained value of
k at those temperatures. The values of k at a particular tem-
pderature obtained at four different aciditides were found to be
in agreement with each other. Thus at a temperature of 298 K,
+
the straight-line plots of 1/k versus 1/[Fruc] for [H ] at 0.01,
−1
0.02, 0.03, and 0.04 mol L yielded the values of k to be
d
5.05, 4.90, 5.36, and 5.46, respectively. Equation 8 may also
be rearranged as
{(
) (
)}
ꢍ−1
=
1+ꢑ2 [Fruc] ∕ ꢍdꢑ2 [Fruc]
{
(
)} [
]
−1
+
1∕ ꢍdꢑ1ꢑ2 [Fruc]
H+
(10)
The above relation justifies the linear plot of k−1 versus
+ −1
[H ] (Figure 7). Moreover, the intercept from such plots
were found to be 279, 225, 185, and 163 at 293, 298, 303,
and 308 K, respectively, whereas the calculated values of
{(1 + K [Fruc]) / (k
K
[Fruc])} were 292, 246, 203,
and 178. 2The values werde found to agree to a good extent.
2
−1
18 J K mol−1), which might be due to the formation of

8
MIDYA ET AL.
a cyclic complex (Scheme 2). At this point, it will be worth-
while to compare the rates of oxidation of glucose34 and fruc-
tose by the same oxidant as well as their different activation
parameters and equilibrium constants. The rate of oxidative
degradation of fructose is much faster than that of glucose.
However, their enthalpies and entropies of activation are very
near to each other and their Gibbs energies of activation (ΔG#)
are almost similar (Table 2). Thus the faster rate of fructose
oxidation cannot be explained simply on the basis of acti-
vation parameters. But if one examines the values of equi-
librium constants (K ) for complex formation (Scheme 1), it
becomes evident tha2t the value is much higher for fructose
than that for glucose (Table 2). The much greater value of K
for fructose may be ascribed to the fact that the environment2s
at the reaction centers are different in the two cases. In case
of glucose,34 the intermediate complex is formed through the
–OH group associated with C-1 carbon and thus the H atom
in the axial position of C-1 is incapable of forming any H-
bond with MnO2. But in fructose the reaction center is C-
2, the –OH group and in the axial position –CH2OH group
(C-1) are capable of forming H-bonding with residual MnO2
and thus the intermediate complex formation is favored in
case of fructose, resulting in a higher concentration of the
complex (Scheme 1), thereby making the reaction rate faster
than that of glucose. In addition, the enthalpy of activa-
tion for fructose degradation is slightly less than that of
glucose.
SCHEME 3 Berezin model showing reaction occurring in both
aqueous and micellar pseudophases
dilution effect). Berezin et al.52 proposed that a normal reac-
tion might occur in the aqueous pseudophase and simultane-
ously both the reactants get solubilized in the micellar pseu-
dophase and then react (Scheme 3).
According to Berezin's model52 (Scheme 3), the pseudo–
first-order rate constant (kꢔ ) at surfactant concentrations
above CMC will be
(
)
2
ꢍW + ꢍMꢑFꢑO Csurf − CMC
ꢍꢔ
=
[
(
)] [
(
)]
1 + ꢑF ꢖsurf − CMC 1 + ꢑO ꢖsurf − CMC
(11)
where C
is the surfactant concentration, k and k are
surf
w
the pseudo–first-order rate constants in aqueous and miMcellar
pseudophases, respectively, and K and K are the binding
constants of fructose and oxidant FMnO2, respectively, with
the surfactant. Considering the weak interactions between
the micelles and the reactants and also the very low con-
centrations of the surfactants present, Equation 11 may
be simplified to Equation 12 as described in our earlier
works:32,34
O
3.8 The role of microheterogeneous
environments in affecting the reaction rate
In the presence of surfactants, hydrophobic and electrostatic
interactions occur between the reactants and the surfactant
monomers/aggregates, which may either favor or restrict the
movement of the attacking reagent and thus the rate of the
reaction is either enhanced or retarded.26,45–49 For all the
three surfactants, there is an initial decrease of the reaction
rate with an increase in [surfactant]. There may be the ion–
dipole interaction between the positively charged head group
of CTAB and the oxygen atom of the polar –OH groups of
the fructose molecule. In case of TX-100 and Tween 20,
there may be interaction between the –OH groups of fruc-
tose and the polar oxyethylene groups and the –OH groups
of the surfactant monomers through H-bonding. Such inter-
actions may lead to the formation of small aggregates, called
premicelles,45–51 which may either sterically or entropically
inhibit the approach of the oxidant, thus decreasing the reac-
tion rate. It is evident from Figure 8 that there is either a hump
or a maximum near the CMC when both the reactants become
associated with the micelle and come in the vicinity of each
other, and beyond this point kꢔ decreases slowly with further
increase of [surfactant], possibly due to lowering of concen-
trations of both the reactants within the micelle (known as
(
)
−1
ꢍꢔ = ꢍW−1 + ꢑF + ꢑO ꢍW−1(ꢖsurf − CMC) (12)
−1
The relation (12) justifies the linear plots of kꢔ
versus
(C − CMC) at four different temperatures for all the three
surf
surfactants (Figure 9). The values of kW at four different tem-
peratures were obtained from the intercepts of these straight
lines. The values were found to corroborate with the exper-
imental values of the pseudo–first-order rate constants (k) at
the respective temperatures in aqueous medium where no sur-
factant was present. Moreover, at these four temperatures, the
values of (K + K ) were evaluated (Table 3) from the slopes
F
O
and intercepts of these straight lines in Figure 9. The values
of standard enthalpy changes (ΔH0) associated with the two
equilibria (Scheme 3) are −42.2 1.5, −19.9 3.5 and −27.8
0.6 kJ mol−1, respectively, for CTAB, TX-100 and Tween
20. The standard entropy change (ΔS0) values are −87 1,
−28 12, and −38 2 J K−1 mol−1 for CTAB, TX-100, and
Tween 20, respectively.

MIDYA ET AL.
9
TA B L E 3 Values of K + K for different surfactants at different
F
O
temperatures
KF + KO
CTAB
890
Temperature (K)
TX-100
125
Tween 20
999
293
298
303
308
681
117
820
493
102
677
388
84
575
ion–dipole interactions and possibly resides in the Stern layer
(Scheme 4). In case of TX-100 and Tween 20, the interaction
between substrate and surfactant molecules occur through H-
bonding. As explained in case of glucose34, K will predomi-
nate in the value of (K + K ) and hence the eFnthalpy change
F
O
(ΔH0) will actually stand for the fructose–micelle binding. It
is plausible that the fructose–micelle interaction is stronger
than fructose–water interaction and consequently when fruc-
tose is transferred from the aqueous pseudophase to the micel-
lar pseudophase, the resultant ΔH0 values are negative.34
One −OH group and nine to ten polyoxyethylene groups are
present in the molecule of TX-100, and these groups are
hydrophilic in nature. When fructose binds with TX-100, the
structure formation leads to a decrease in entropy, whereas
release of water molecules surrounding fructose results in an
increase in entropy. Possibly the decrease in entropy is greater
−1
F I G U R E 9 Plots of kΨ versus (C – CMC) in the post-CMC
surf
region for CTAB. [H ] = 0.001 mol L−1, [MnO2] = 1.0 × 10−4 mol l
,
+
−1
−1
and [Fruc] = 0.1 mol L [Color figure can be viewed at
wileyonlinelibrary.com]
Similar to our earlier discussion,34 it may be assumed that
protonated MnO2, through ion–ion or ion–dipole interaction,
will reside at the outer region of the Stern layer, while the fruc-
tose molecule, having a number of –OH groups in its struc-
ture, will be interacting with the CTAB molecules through
SCHEME 4 Possible interaction between
reactants and CTAB micelles involving
intermolecular H-bonding [Color figure can be
viewed at wileyonlinelibrary.com]

10
MIDYA ET AL.
than the increase in entropy, the resultant ΔS0 becomes nega-
tive. On the other hand, Tween 20 contains three –OH groups
and 20 polyoxyethylene groups in its molecule and hence
its interaction with fructose will be much greater than the
interaction of TX-100 with fructose. This factor accounts for
the higher negative values of both ΔH0 and ΔS0 in case of
Tween 20. This stronger interaction with Tween 20 is justi-
fied by the higher value of (K + K ) for Tween 20 over TX-
REFERENCES
1. Higginson WCE, Sutton D. The oxidation of hydrazine in aqueous
solution. Part II. The use of 15 N as a tracer in the oxidation of
hydrazine. J Chem Soc. 1953:1402–1406.
2. Kirschenbum LJ, Sutter JR. Kinetic studies of permanganate oxi-
dation reactions. I. Reaction with iodide ion. J Phys Chem.
1966;70:3863–3866.
3. Zahonyi-Budo E, Simandi LI. Induced oxidation of phospho-
rus(III) by a short-lived manganate(V) intermediate in the per-
manganate oxidation of arsenite(III). Inorg Chim Acta. 1992;191:
1–2.
F
O
100 (Table 3). Therefore, it may be concluded that both the
enthalpy changes and the entropy changes associated with the
fructose–water interaction, fructose–micelle interaction, and
micelle–water interaction finally control the binding of fruc-
tose with the micelle.
4. Sen Gupta KK, Sen PK, Mukhopadhyay G. Evidence of esterifica-
tion during the oxidation of some aromatic aldehydes by perman-
ganate. Transit Met Chem. 1993;18:369–371.
5. Sen PK, Sanyal A, Sen Gupta KK. Evidence of protonation during
the oxidation of some aryl alcohols by permanganate in perchloric
acid medium and mechanism of the oxidation processes. Int J Chem
Kinet. 1995;27:379–389.
4
CONCLUSIONS
Nanoparticles of MnO2 are found to be present in the colloidal
solution of MnO2 and these nanoparticles remain adhered to
two dimensional gum acacia sheets as spherical particulates
of size 50–200 nm. The oxidation of fructose takes place on
the active surface of the nanoparticles of MnO2 adhered on to
the surface of gum acacia sheets and occurs through an inter-
mediate complex formation between ꢀ-D-fructopyranose and
protonated MnO2. The complex disproportionates in a two-
electron transfer step to form a lactone and a hydroxymethyl
carbocation, which are ultimately converted to an aldonic
acid and formic acid, respectively. The reaction appears to
be controlled by the entropy of activation. Intermediate com-
plex formation between carbohydrate and protonated MnO2
is more favored in case of fructose than glucose leading to
a higher rate of oxidation for fructose than glucose. For the
reaction in the presence of different microenvironments, ion-
dipole interactions play a major role in the binding of fruc-
tose with CTAB, whereas H-bonding or dipole–dipole inter-
actions are mainly responsible for Tween 20 and TX-100.
Both the enthalpy and entropy changes involved in different
interactions ultimately control the binding of fructose with the
micelle.
6. Sen PK, Mukhopadhyay G, Sen Gupta KK. Kinetics and mechanism
of the oxidation of alkyl and aryl methyl ketones by permanganate
ion in aqueous ethanoic acid. Transit Met Chem. 1998;23:577–
581.
7. Sen PK, Mukhopadhyay G, Gani N, Sen Gupta KK. Relative
reactivities of 1,2-dicarbonyl compounds towards permanganate in
ethanoic acid medium. Mechanism of the oxidation processes. Tran-
sit Met Chem. 2001;26:451–456.
8. Sen PK, Samaddar PR, Das K. Reductivity of methanol and ethanol
towards permanganate in absence and presence of Tween 20 in per-
chloric acid medium. Mechanism of the oxidation processes. Transit
Met Chem. 2005;30:261–267.
9. Sen PK, Samaddar PR, Das K. Kinetics and mechanism of the
oxidation of propan-1-ol and propan-2-ol by permanganate in the
absence and presence of Tween 20 in perchloric acid medium. Tran-
sit Met Chem. 2005;30:907–914.
10. Andrabi SMZ, Malik MA, Khan Z. Permanganate partitioning in
cationic micelles of cetyltrimethylammonium bromide: a kinetic
study of D-fructose oxidation. Colloids Surf A: Physicochem Eng
Asp. 2007;299:58–64.
11. Malik MA, Al-Thabaiti SA, Khan Z. Kinetics of oxidation of
D-glucose by permanganate in aqueous solution of cetyltrimethy-
lammonium bromide. Colloids Surf A: Physicochem Eng Asp.
2009;337:9–14.
12. Perez-Benito JF, Arias C. A kinetic study of the reaction between
soluble (colloidal) manganese dioxide and formic acid. J Colloid
Interface Sci. 1992;149:92–97.
ACKNOWLEDGMENTS
Financial assistances from Department of Science and Tech-
nology, Government of India, through the PURSE Pro-
gramme (Phase II) and from University Grants Commission,
New Delhi, India, through the CAS programme (Phase II) to
the Department of Chemistry, Jadavpur University, Kolkata,
India, are gratefully acknowledged.
13. Perez-Benito JF, Arias C, Amat E. A kinetic study of the reduction
of colloidal manganese dioxide by oxalic acid. J Colloid Interface
Sci. 1996;177:288–297.
14. Kabir-ud-Din, Fatima W, Khan Z. Effect of surfactants on the oxi-
dation of oxalic acid by soluble colloidal MnO2. Colloids Surf A:
Physicochem Eng Asp. 2004;234:159–164.
15. Akram M, Altaf M, Kabir-ud-Din. Water-soluble colloidal man-
ganese dioxide as an oxidant for L-tyrosin in the absence
and presence of non-ionic surfactant TX-100. Colloids Surf, B.
2009;73:308–314.
ORCID
Pratik K. Sen
https://orcid.org/0000-0002-5863-0959

MIDYA ET AL.
11
16. Khan Z, Raju, Akram M, Kabir-ud-Din. Oxidation of lactic acid
by water-soluble (colloidal) manganese dioxide. Int J Chem Kinet.
2004;36:359–366.
34. Sen PK, Midya JK, Bysakh S, Pal B. Kinetic and mechanistic stud-
ies on the oxidation of D-glucose by MnO2 nanoparticles. Effect
of microheterogeneous environments of CTAB, Triton X-100 and
Tween 20. Mol Catal. 2017;440:75–86.
17. Lodish H, Berk A, Zipursky SL. Molecular Cell Biology. 4th ed.
New York, NY: W. H. Freeman; 2000.
35. Akram M, Altaf M, Kabir-ud-Din. Kinetics of colloidal MnO2
reduction by L-arginine in absence and presence of surfactants. Col-
loid J. 2011;73:149–157.
18. Jequier E. Carbohydrates as a source of energy. Am J Clin Nutr.
1994;59:6825–6855.
36. Akram M, Altaf M, Kabir-ud-Din. Oxidative degradation of dipep-
tide (glycyl-glycin) by water-soluble colloidal manganese dioxide
in the aqueous and micellar media. Colloids Surf, B. 2011;82:217–
223.
19. Gupta M, Saha SK, Banerjee P. Kinetics and mechanism of the
reduction of dodecatungstocobaltate (III) by D-fructose, D-glucose,
and D-mannose: comparison between keto- and aldohexoses. J
Chem Soc, Perkin Trans. 1978;2:1781–1785.
37. Das D, Sen PK, Das K. Carbohydrate electro-oxidation on chemi-
20. Signorella S, Civllo L, Lafarga R, Sala LF. The role of acetate in
the oxidation of D-galactose by Cu(II) in acetate-acetic acid buffers.
New J Chem. 1996;20:989–991.
cally prepared MnO2. J Electroanal Chem. 2007;611:19–25.
38. Feigl F. Spot Tests in Organic Analysis. New York, NY: Elsevier;
1975:452.
21. Ribeiro ACF, Lobo VMM, Valente AJM, et al. Association between
ammonium monovanadate and ꢀ-cyclodextrin as seen by NMR and
transport techniques. Polyhedron. 2006;25:3581–3587.
39. Andrabi SMZ, Malik MA, Khan Z. Permanganate partitioning in
cationic micelles of cetyltrimethylammonium bromide: a kinetic
study of D-fructose oxidation. Colloids Surf A: Physicochem Eng
Asp. 2007;299:58–64.
22. Olavi P, Virtanen I, Kurkisuo S. Oxidation of D-fructose with
vanadium(V): a kinetic approach. Carbohydr Res. 1985;138:215–
223.
40. Kabir-ud-Din, Ali MS, Khan Z. Effect of surfactant micelles on the
kinetics of oxidation of D-fructose by cerium(IV) in sulfuric acid
medium. Int J Chem Kinet. 2006;38:18–25.
23. Sen Gupta KK, Begum BA, Pal B. Kinetic behavior and rela-
tive reactivities of some aldoses, amino sugars and methylated
sugars towards platinum(IV) in alkaline medium. Carbohydr Res.
1998;309:303–310.
41. Kabir-ud-Din, Morshed AMA, Khan Z. Influence of sodium
dodecyl sulfate/Triton X-100 micelles on the oxidation of D-
fructose by chromic acid in presence of HClO4. Carbohydr Res.
2002;337:1573–1583.
24. Sen Gupta KK, Begum BA. Kinetics and mechanism of the
oxidation of some aldoses, amino sugars and methylated sug-
ars by tris(pyridine-2-carboxylato)manganese(III) in weakly acidic
medium. Carbohydr Res. 1999;315:70–75.
42. Singh AK, Srivastava J, Rahmani S, Singh V. Pd(II)-catalysed
and Hg(II)-co-catalysed oxidation of D-glucose and D-fructose
by N-bromoacetamide in the presence of perchloric acid: a
kinetic and mechanistic study. Carbohydr Res. 2006;341:397–
409.
25. Sen Gupta KK, Pal B, Begum BA. Reactivity of some sugars and
sugar phosphates towards gold(III) in sodium acetate-acetic acid
buffer medium. Carbohydr Res. 2001;330:115–123.
43. Azmat R, Naz R, Qamar N, Malik I. Kinetics and mechanisms of
oxidation of D-fructose and D-lactose by permanganate ion in acidic
medium. Nat Sci. 2012;4:466–478.
26. Fendler JH, Fendler EJ. Catalysis in Micellar and Macromolecular
Systems. New York, NY: Academic Press; 1975.
27. Bunton CA. Reactivity in aqueous association colloids. Descrip-
tive utility of the pseudophase model. J Mol Liq. 1997;72:231–
249.
44. Qamruzzaman, Nasar A. Degradation of tricyclazole by colloidal
manganese dioxide in the absence and presence of surfactants. J
Ind Eng Chem. 2014;20:897–902.
28. Cheong MY, Ariffin A, Khan MN. A comparative analysis of pseu-
dophase ion-exchange (PIE) model and Berezin pseudophase (BPP)
model: analysis of kinetic data for ionic micellar-mediated semi-
ionic bimolecular reaction. Bull Korean Chem Soc. 2007;28:1135–
1140.
45. Brinchi L, Profio PD, Germani R, Savelli G, Tugliani M, Bun-
ton CA. Hydrolysis of dinitroalkoxyphenyl phosphates in aque-
ous cationic micelles. Acceleration by premicelles. Langmuir.
2000;16:10101–10105.
46. Sen PK, Chatterjee P, Pal B. Evidence of co-operativity in the pre-
micellar region in the hydrolytic cleavage of phenyl salicylate in the
presence of cationic surfactants of CTAB, TTAB and CPC. J Mol
Catal A: Chem. 2015;396:23–30.
29. Holmberg K. Organic reactions in microemulsions. Eur J Org
Chem. 2007;5:731–742.
30. Mao S, Chen Z, An X, Shen W. Study of the alkaline fading of phe-
nolphthalein in microemulsions. J Phys Chem A. 2011;115:5560–
5567.
47. Sen PK, Talukder S, Pal B. Specific interactions of anions and pre-
micelles in the alkaline fading of crystal violet carbocation. Colloids
Surf A: Physicochem Eng Asp. 2015;467:259–269.
31. Sen PK, Gani N, Pal B. Effects of non-ionic micellar aggregates on
the electron transfer reaction between L-glutamic acid and gold(III)
complexes. Int J Chem Kinet. 2012;44:482–493.
48. Sen PK, Mukherjee P, Pal B. Effects of pre-micelles of anionic sur-
factant SDS on the electron transfer reaction between methylene
blue and ascorbic acid. J Mol Liq. 2016;224:472–479.
32. Sen PK, Gani N, Pal B. Effects of microheterogeneous environ-
ments of SDS, TX-100 and Tween 20 on the electron transfer reac-
−
−
49. Maiti K, Sen PK, Pal B. influence of premicelles and micellar aggre-
gates of ionic and nonionic surfactants in the oxidative decarboxy-
lation of L-lycine by gold(III) complexes. J Mol Liq. 2018;251:238–
248.
tion between L-leucine and AuCl4 /AuCl3(OH) . Ind Eng Chem
Res. 2013;52:2803–2813.
33. Samiey B, Dalvand Z. Study of fuchsin acid fading in micellar
media. Int J Chem Kinet. 2014;46:651–661.

12
MIDYA ET AL.
50. Pisckiewicz D. Positive cooperativity in micelle-catalysed reac-
tions. J Am Chem Soc. 1977;99:1550–1557.
How to cite this article: Midya JK, Ghosh DC,
51. Pisckiewicz D. Cooperativity in bimolecular micelle-catalysed
reactions. Inhibition of catalysis by high concentrations of deter-
gent. J Am Chem Soc. 1977;99:7695–7697.
Pal B, Sen PK. Effect of microheterogeneous envi-
ronments of CTAB, Triton X-100, and Tween 20 on
the oxidative degradation of D-fructose by nanopar-
ticles of MnO2. Int J Chem Kinet. 2018;1–12.
https://doi.org/10.1002/kin.21239
52. Berezin IV, Martinek K, Yatsmiriskii AK. Physicochemical foun-
dations of micellar catalysis. Russ Chem Rev (Eng Transl).
1973;42:787–802.