Kinetics of the MnO4- oxidation of anionic surfactant (sodiumdodecyl sulphate): Evidence for the formation of soluble colloidal MnO2

Article abstract of DOI:10.1246/bcsj.78.1218

A conventional spectrophotometric technique was used to study the oxidation of SDS by permanganate in a perchloric acid medium. It was observed that the reaction proceeded in two stages (fast first stage followed by a relatively slow second stage). Plots of log(absorbance) versus time deviate from linearity. The kinetic and spectroscopic data are consistent with the formation of soluble colloidal MnO₂. The first-order kinetics with respect to [SDS] at low concentrations shifted to second-order at higher concentrations. The kinetics of oxidation is first-order with respect to both [MnO₄^-] and [HClO₄]. The oxidation rate was decreased by the addition of P₂O₇^4- and Mn(II) ions. Second-step oxidation is not a true path for the oxidation of the Mn(IV)-SDS reaction. In the presence of Mn(II) (a reaction product), the MnO₄^- oxidation of SDS becomes more complicated, and an exact dependence on [Mn(II)] can not be estimated. Different activation parameters have been evaluated. Mechanisms consistent with the kinetic data have been proposed and discussed. The -O-SO₃^- group is responsible for the oxidative degradation of SDS by MnO₄^-.

Full text of DOI:10.1246/bcsj.78.1218

1218 Bull. Chem. Soc. Jpn., 78, 1218–1222 (2005)
Ó 2005 The Chemical Society of Japan
À
Kinetics of the MnO4 Oxidation of Anionic Surfactant (Sodiumdodecyl
Sulphate): Evidence for the Formation of Soluble Colloidal MnO2
ꢀ
Raju and Zaheer Khan
Department of Chemistry, Jamia Millia Islamia, (Central University), Jamia Nagar, New Delhi 110025, India
Received July 20, 2004; E-mail: drkhanchem@yahoo.co.in
A conventional spectrophotometric technique was used to study the oxidation of SDS by permanganate in a
perchloric acid medium. It was observed that the reaction proceeded in two stages (fast first stage followed by a rela-
tively slow second stage). Plots of log(absorbance) versus time deviate from linearity. The kinetic and spectroscopic
data are consistent with the formation of soluble colloidal MnO2. The first-order kinetics with respect to [SDS] at
low concentrations shifted to second-order at higher concentrations. The kinetics of oxidation is first-order with respect
to both [MnO4ꢁ] and [HClO4]. The oxidation rate was decreased by the addition of P2O74ꢁ and Mn(II) ions. Second-step
oxidation is not a true path for the oxidation of the Mn(IV)–SDS reaction. In the presence of Mn(II) (a reaction product),
ꢁ
the MnO4 oxidation of SDS becomes more complicated, and an exact dependence on [Mn(II)] can not be estimated.
Different activation parametꢁers have been evaluated. Mechanisms consistent with the kinetic data have been proposed
ꢁ
and discussed. The –O–SO3 group is responsible for the oxidative degradation of SDS by MnO4
.
(¼6:0 cm3, 11.63 mol dmꢁ3), and HgCl2 (¼1:2 g) in a volume of
In a previous publication we showed that a non-ionic, TX-
100, surfactant is unstable in aqueous solution, and acts as a
reductant for cerium(IV).1 At that time we pointed out that a
primary –OH group of a polyoxyethylene chain of TX-100
was responsible for the oxidative degradation of TX-100. In
the present study, we explored a related family of reactions
involving the anionic surfactant (sodiumdodecyl sulphate) and
ꢄ
50 cm3. The reaction mixture was heated to 40 C for 30 min. A
white precipitate of mercurous chloride appeared slowly, showing
free-radical intervention. No precipitate formation could bꢁe
observed under the experimental conditions with either MnO4
or SDS alone.
The kinetics measurement was carried out by mixing the solu-
tions in a three-necked reaction flask fitted with a condenser (to
arrest evaporation). The reaction flask was kept inꢄa thermostate
controlled at the desired temperature within ꢅ0:1 C. The prog-
ress of the reaction was followed spectrophotometrically by pipet-
ting aliquots of the reaction mixture at definite time intervals, and
the absorbance was measured at 525 nm (ꢀmax for MnO4ꢁ). The
concentrations of SDS were always much in excess over that of
permanganate. The pseudo-first-order rate constants were calculat-
ed from the slope of the conventional log(absorbance) versus t
(time) plots. The reactions were usually followed by the complete
potassium permanganate. In this paper, we wish to report on
ꢁ
the results of our work on the oxidation of SDS by MnO4
.
A large number of books and review articles are available
concerning the role of surfactants in a variety of organic and
inorganic reactions.2–15 The common behavior of micelle-cat-
alyzed reactions is the solubilization of the substrate(s). How-
ever, despite extensive research on the effect of the surfactant,
the kinetics and mechanistic aspects of the reaction involving
micelle as a reductant have not been studied so far, except for a
recent report by a few researchers.1,16
ꢁ
disappearance of MnO4 color. However, a few runs were also
performed at 425 nm to confirm the autocatalytic nature of the
oxidation. Other details of the kinetic procedure are described
elsewhere.17–20
Experimental
Sodiumdodecyl sulphate (Merck, India, 99%), potassium per-
manganate (Merck, India, 99%), mercuric chloride (Merck, India,
99%), hydrochloric acid (Merck, India, 36%), 2,4-dinitrophenyl-
hydrazine (Qualigens, India, 98%), dimethylsulphoxide-d6 (Merck,
Results and Discussion
The rate of disappearance of permanganate ion is shown in
Fig. 1 as a log(absorbance)–time profile. The sigmoid nature of
these plots suggests the existence of an autocatalytic reaction
path. The extent of the induction period (non autocatalytic
reaction path) depends on the experimental conditions, i.e.,
[SDS], [HClO4], and temperature. Therefore, choice of the
best experimental conditions for the kinetic experiments is a
crucial problem that we address first. In order to examine the
effects of the variables, experiments were tried at 1:2 ꢃ 10ꢁ4
Germany, 99.8%), manganese(II) chloride (MnCl2 4H2O, Quali-
ꢂ
gens, India, 98%), and sodium pyrophosphate (Merck, India,
99%) were used as such without further purification. In order to
check the purity of SDS, the critical micelle concentration
(cmc) was determined conductometrically, and the obtained value
was found to be in good agreement with the reported cmc. The
[Hþ] of the medium was adjusted with the required [HClO4]
(Merck, India, 60% reagent). Solutions of the entire reagent werꢁe
prepared in double-distilled, deionized and boiled water. MnO4
solutions were prepared and tested by the vogel method.
to 2:4 ꢃ 10ꢁ4 mol dmꢁ3, 4:0 ꢃ 10ꢁ3 to 30:6 ꢃ 10ꢁ3 mol dmꢁ3
,
and 0.73 to 3.28 mol dmꢁ3 of permanganate, SDS, and HClO4,
respectively. HClO4 was used as an acidifying agent due to the
A permanganate solution (¼6:0 cm3, 1:0 ꢃ 10ꢁ3 mol dmꢁ3
)
was added to a mixture of SDS (¼5:0 cm3, 0.1 mol dmꢁ3), HClO4
Published on the web July 1, 2005; DOI 10.1246/bcsj.78.1218

Raju et al.
Bull. Chem. Soc. Jpn., 78, No. 7 (2005) 1219
log [SDS]
−3
[SDS]/10−2 moldm
Fig. 1. Plots of log(absorbance) versus time at 525 nm for
ꢁ
the MnO4 oxidation of SDS at different temperatures.
Reaction conditions: ½SDSꢆ ¼ 10:0 ꢃ 10ꢁ3 mol dmꢁ3
Fig. 2. Plot of kobs1 versus [SDS]. Inset: log–log plot be-
tween kobs1 and [SDS]. Reaction conditions: ½SDSꢆ ¼ 0:0
to 30:0 ꢃ 10ꢁ3 mol dmꢁ3, ½MnO4ꢁꢆ ¼ 1:2 ꢃ 10ꢁ4 mol
dmꢁ3, ½HClO4ꢆ ¼ 1:46 mol dmꢁ3, temperature ¼ 40 ꢄC.
,
½MnO4ꢁꢆ ¼ 1:2 ꢃ 10ꢁ4 mol dmꢁ3, ½HClO4ꢆ ¼ 1:46 mol
dmꢁ3
.
Table 1. Pseudo-First-Oꢁrder Rate Constants for the Oxida-
tion of SDS by MnO4 at 313 K
[MnO4]
ꢃ 104
[SDS]
ꢃ 103
[HClO4] kobs1 ꢃ 104 kobs2 ꢃ 104
/mol dmꢁ3
1.46
/sꢁ1
/sꢁ1
/mol dmꢁ3 /mol dmꢁ3
1.2
1.4
1.6
1.8
2.0
2.4
1.2
10.0
3.4
3.0
2.3
1.5
1.5
1.5
1.5
2.3
3.4
5.3
6.1
8.4
10.7
13.8
1.5
3.4
5.3
6.6
7.6
9.2
1.9
1.5
1.5
1.5
0.7
0.7
−3
[HCIO4]/moldm
T−1/10−3 K−1
4.0
8.0
1.46
not observed
not observed
Fig. 3. A-Plot of kobs1 versus [HClO4]. Reaction condi-
tions: ½SDSꢆ ¼ 10:0 ꢃ 10ꢁ3 mol dmꢁ3, ½MnO4ꢁꢆ ¼ 1:2 ꢃ
10ꢁ4 mol dmꢁ3ꢁ, temperature at 40 ꢄC. B-Arrhenius plot
for the MnO4 oxidation of SDS. Reaction conditions
were the same as in Fig. 1.
10.0
14.0
18.0
22.0
26.0
30.0
10.0
1.9
4.2
5.3
7.6
7.6
9.2
ꢁ
HSO4 and SO42ꢁ). To maintain the concentration of the
Hþ ion constant, HCl cannot be used as a source of Hþ ion.
The effect of the Hþ ion concentration was investigated by
keeping the [MnO4ꢁ] and [SDS] constants and varying [Hþ]
of the solution. There was an increase in the rate of the reac-
tion with an increase in [HClO4] (Table 1). The values of
kobs were plotted against [HClO4], which gave a straight line
passing through the origin with a slope of 1.0 (Fig. 3A). The
reaction was, therefore, first-order with respect to [HClO4].
The reactions were also studied in the temperature range 30–
1.2
0.73
1.46
1.82
2.55
2.92
3.28
not observed
1.9
3.8
5.3
6.1
6.9
non-complexing nature of ClO4ꢁ. The rate constants of the
non-autocatalytic route were determined by measuring the
slopes of the lines extended over the induction period. The
oxidation of SDS proceeded in two stages, i.e., initial fast stage
followed by a relatively slower step.
At constant [Hþ], the pseudo-first-order rate constants
increased with increases in [SDS] from 4.0 to 30:0 ꢃ 10ꢁ3
mol dmꢁ3 (Table 1). The plot of kobs1 versus [SDS] was a
curve passing through the origin; Figure 2 shows the relevant
curve. The reaction order in SDS changes from first towards
second as the [SDS] is increased. Generally, H2SO4 and
HClO4 are used as sources of hydrogen ion in the redox reac-
tion of permanganate. H2SO4 contains at least two catalyzing
species, namely, Hþ ion and HSO4ꢁ, which complicate the
mechanism. The investigations were made in the presence of
HClO4, because the perchlorate ion is a weak ligand (unlike
ꢄ
60 C (Table 2). The plot of log kobs1 versus 1=T was linear
(Fig. 3B). The value of Ea (activation energy) was calculated
from the slope of Fig. 3B, plot, and is recorded in Table 2
along with other parameters. A useful explanation of the acti-
vation of enthalpy and the activation of entropy is not possible,
because the kobs1 does not represent a single elementary step; it
is a complex function of ionization, the true rate, and binding
constants. However, the large negative values of the entropy of
the activation show that the transition state is well-structured
and highly solvated. Thus, the oxidation of SDS is an entro-
py-controlled rather than an enthalpy-controlled one. The
order with respect to [MnO4ꢁ] was determined by studying the
reaction at different initial [MnO4ꢁ] values. The plots of
log(absorbance) versus time were linear, indicating a first-order
dependence on [MnO4ꢁ] in the initial stages of the reactions.

ꢁ
1220 Bull. Chem. Soc. Jpn., 78, No. 7 (2005)
Oxidative Degradation of Sodiumdodecyl Sulphate by MnO4
Table 2. Pseudo-First-Order Rate Constants and Activation
Parameters (Ea, ꢀH#, ꢀS#) For the Oꢁxidation of SDS
(¼14:0 ꢃ 10ꢁ3 mol dmꢁ3
)
by MnO4
)
(¼1:2 ꢃ 10ꢁ4
mol dmꢁ3), [HClO4] (¼1:46 mol dmꢁ3
Temperature
/K
kobs1 ꢃ 104
kobs2 ꢃ 104
/sꢁ1
/sꢁ1
303
313
323
333
2.3
1.5
4.2
7.3
9.2
5.3
10.7
20.7
Activation Parameters:
Ea (kJ molꢁ1
Fig. 4. Plot of absorbance versus time at 525 nm. Inset: plot
of absorbanceꢁat 525 nm versus absorbance at 420 nm
for the MnO4 oxidation of SDS. Reaction conditions:
½SDSꢆ ¼ 10:0 ꢃ 10ꢁ3 mol dmꢁ3, ½MnO4ꢁꢆ ¼ 1:2 ꢃ 10ꢁ4
)
57
54
ꢁ122
80
77
ꢁ88
ꢀH# (kJ molꢁ1
ꢀS# (J Kꢁ1 molꢁ1
)
)
mol dmꢁ3
,
½HClO4ꢆ ¼ 1:46 mol dmꢁ3
, temperature =
ꢄ
40 C.
On the other hand, the pseudo-first-order rate constants de-
creased with an increase in [MnO4ꢁ], which may be due to
the possible flocculation of the colloidal MnO2 particles (inter-
mediate). This type of behavior (dependent of pseudo-first-
order rate constant on the initial concentration of the reactant
in defect) has been observed in many permanganate reactions,
and especially for those reactions, which has an autocatalytic
reaction path.21–23
It has been established that the oxidation of organic sub-
strates by permanganate is characterized by two principal
processes:24
Fig. 5. Plots of absorbance versus time for the formation of
intermediate (colloidal MnO2) during the oxidation of
[SDS] (¼ 10:0 ꢃ 10ꢁ3 ꢄmol dmꢁ3) by [MnO4ꢁ] (¼1:2 ꢃ
10ꢁ4 mol dmꢁ3) at 40 C and ½HClO4ꢆ ¼ 0:73 ( ), 1.46
Mn(VII) þ organic reductant(s) ꢁ! Intermediate; ð1Þ
Intermediate ꢁ! Mn(II);
ð2Þ
where Int. = one (or several) reaction intermediate(s).
(
), and 2.92 ( ) mol dmꢁ3
.
Therefore, in order to obtain insight into the autocatalytic
reaction path and to confirm the formation of an intermediate
(Fig. 4), the oxidation of SDS was also studied by monitoring
the absorbance of the reꢁaction mixture at 420 nm, where the
contribution from MnO4 is negligible. A plot of the absorb-
ance at 525 nm versus absorbance at 420 nm was linear
(Fig. 4, inset), indicating that product formation occurred at
the same rates as the reaction of permanganate ion.25 The spe-
cies, which absorbs light at 420 and 525 nm, is a soluble spe-
cies of Mn(IV). In addition, some experiments were also car-
ried out at 420 and 525 nm. These results are shown graphical-
ly in Fig. 5 as an absorbance time profile at different [HClO4]
values. At 420 nm, the absorbance increases until it reaches a
maximum, and then decreases with time. This behavior indi-
cates the formation and disappearance of intermediate (water
soluble colloidal MnO2) during the course of the reaction.
The formation of an intermediate was very sensitive to
[HClO4]. Surprisingly, as the HClO4 increased from 0.73 to
3.28 mol dmꢁ3, the formation of intermediate was not ob-
was a Mn(IV) species (Mn(IV) is commonly involved in the
permanganate oxidation of organic reductants). The spectra
of Mn(IV) depend on the experimental conditions as well as
the nature of the reducing agent. The spectra of this intermedi-
ate showed a maximum absorption at 390 nm, and the absorb-
ance uniformly decreased with the wavelength over the whole
visible region.27,28 In addition, some experiments were also
carried out at 420 nm. On the other hand, the plot of log ab-
sorbance versus log wavelength is also linear,29 suggesting that
the formation of water-soluble colloidal species of Mn(IV) as
an intermediate product during the SDS oxidation of MnO4
in the presence of HClO4.
The kinetic parameters experimentally found in this work
are consistent with the following tentative mechanism for the
formation of Mn(IV) as an intermediate.
In Scheme 1, reaction (5) represents the formation of a com-
plex (C) with permanganate and protonated SDS. By analogy
with previous results,24,25 we assume that C decomposes by a
one-stop, three-electron oxidation–reduction mechanism di-
rectly to manganese(IV) and dodecanal (D). The oxidation
product of SDS, dodecanal, was characterized as follows. In
a typical experiment SDS (¼10:0 ꢃ 10ꢁ2 mol dmꢁ3), HClO4
(¼2:14 mol dmꢁ3), and potassium permanganate (¼5:8 ꢃ
10ꢁ4 mol dmꢁ3) were allowed to react at 30 ꢄC. After comple-
tion of the reaction (complete disappearance of purple color),
a saturated solution of 2,4-dinitrophenylhydrazine in 2 M HCl
was added to the reaction mixtures and the dinitrophenylhydra-
ꢁ
ꢁ
served. In the case of MnO4 as an oxidizing agent, the pos-
sible intermediates are Mn(VI), Mn(V), Mn(IV), and Mn(III).
The presence of Mn(VI) and Mn(V) is ruled out by the fact
that they are highly unstable in an acidic medium.25 As far
as the formation of Mn(III) is concerned, some experiments
were also performed where the 470 nm (characteristic of
Mn(III))26 was used to monitor the formation of Mn(III) during
the course of the reaction, but failed to detect any build up of
Mn(III). Thus, we may safely conclude that the intermediate

Raju et al.
Bull. Chem. Soc. Jpn., 78, No. 7 (2005) 1221
of 0.0, 2.0, 4.0, 8.0, 12.0, and 14.0 mol dmꢁ3, respectively.
This indicates that Mn(III) is the principal species responsible
for the oxidation of SDS in an autocatalytic reaction path
(although our attempts failed to observe the appearance of this
ion, vide supra). The decrease in the reaction rate with an in-
crease in [P2O74ꢁ] may be due to the one-step one-electron
oxidation–reduction while passing through Mn(III) as an inter-
mediate. The proposed mechanism is shown in Scheme 2 for
the reduction of Mn(IV).
fast
-
+ H +
HMnO4
(3)
(4)
MnO4
O
O
Ka
CH3 (CH2)10 CH2 O-S-O -
H+
CH3 (CH2)10CH2O-S-OH
+
O
O
( SDSH )
( SDS )
O
Kc
SDSH + HMnO 4
CH3 (CH2)10 CH2O-S-O-MnO 3 + H2O
(5)
(6)
O
( C )
The rate law, which can be deduced from the proposed
mechanism (Scheme 1), is
k
2-
C
CH3 (CH2)10 CHO + Mn(IV) + SO4
( D )
ꢁd½MnO4ꢁꢆ kKcKa½SDSꢆT½Hþꢆ½MnO4ꢁꢆ
NO2
NO2
¼
;
ð11Þ
dt
ð1 þ Ka½HþꢆÞ
D + NH2NH
NO2
CH3(CH2)10CH=N-NH
( P )
(7)
NO2
and
kKcKa½SDSꢆT½Hþꢆ
ð1 þ Ka½HþꢆÞ
kobs1
¼
:
ð12Þ
Scheme 1.
The inequality 1 ꢇ Ka½Hþꢆ is evident, and the rate equation 12
Kc 1
reduces to
SDSH + Mn(IV)
(8)
SDSH Mn(IV)
( C 1 )
kobs1 ¼ kKcKa½SDSꢆT½Hþꢆ;
which is in complete according to the observations, and thus
supports the mechanism. The derived rate-law is only known
to apply at lower [SDS], and could be different for the higher
[SDS] values.
ð13Þ
O
k 1
.
(9)
C1
Mn(III) + CH 3 (CH2)10 CH2O-S-O
O
( Radical )
fast
Radical
+ Mn(III)
The following rate equation has been derived based on the
proposed mechanism (Scheme 2):
D
+
Mn(II)
(10)
Scheme 2.
ꢁd½Mn(IV)ꢆ k1Kc1Ka½SDSꢆT½Hþꢆ½Mn(IV)ꢆ
¼
;
ð14Þ
ð15Þ
zone was filtered, washed, and dried. The product was identi-
fied as dodecanal 2,4-dinitrophenylhydrazone (P) by compar-
ing the H NMR spectrum in (DMSO-d6) with that of an au-
dt
ð1 þ Ka½HþꢆÞ
k1Kc1Ka½SDSꢆT½Hþꢆ
1
kobs2
¼
:
ð1 þ Ka½HþꢆÞ
thentic sample of SDS. The 1H NMR spectra were run at room
temperature (27 ꢄC) in a solution of DMSO on a Bruker DPX-
300 spectrospin instrument operating at 300.13 MHz.
In order to establish the reduction of Mn(IV) (intermediate
product) and to confirm the formation of Mn(III) as an inter-
mediate, a series of kinetic experiments were carried out in
the presence of P2O74ꢁ (a complexing agent). It was found that
Interestingly, the autocatalytic reaction path (kobs2) is not a
true value of the oxidation of SDS. It may be a mixture of
the rates of SDS and manganese(II); the exact dependence of
the kobs2 on [SDS] and [Hþ] can not be estimated.
In order to further explain the autocatalytic path, the oxida-
tion kinetics of SDS was carried out in presence of Mn(II) (a
reaction product). The kinetic curve between log(absorbance)
and time is linear. It was observed that the extent of the noncat-
alytic reaction pathway disappeared completely. These results
are summarized in Table 3. The decrease in the oxidation rate
with increasing [Mn(II)] indicates that Mn(IV) is involved in
4ꢁ
the oxidation rate of SDS decreased upon adding P2O7 ions.
For example, under the conditions [SDS] (¼10:0 ꢃ 10ꢁ3
ꢁ
mol dmꢁ3), [MnO4
]
(¼1:2 ꢃ 10ꢁ4 mol dmꢁ3), [HClO4]
(¼1:46 mol dmꢁ3), temperature (¼40 ꢄC), 104 kobs values were
4ꢁ
found to be 3.4, 2.3, 1.5, 1.5, 1.5, and 1.5 sꢁ1 at 104 [P2O7
]
ꢁ
Table 3. Pseudo-First-Order Rate Constants for the Oxidation of SDS by MnO4 in Presence
of Mn(II) at 313 K, [HClO4] (¼1:46 mol dmꢁ3
)
[SDS] ꢃ 103
[MnO4ꢁ] ꢃ 104
[Mn(II)] ꢃ 104
kobs1 ꢃ 104
kobs2 ꢃ 104
/mol dmꢁ3
/mol dmꢁ3
/mol dmꢁ3
/sꢁ1
/sꢁ1
10.0
1.2
1.0
2.0
4.0
8.0
12.0
0.0
1.9
not observed
not observed
not observed
not observed
3.4
2.0
1.9
1.5
1.5
1.5
1.9
4.2
10.0
14.0
0.0
1.2
1.2
5.3
0.28
0.28
0.28
1.0
2.0
4.0
not observed
not observed
not observed

ꢁ
Oxidative Degradation of Sodiumdodecyl Sulphate by MnO4
1222 Bull. Chem. Soc. Jpn., 78, No. 7 (2005)
2
New York (1982).
3 J. H. Fendler and E. J. Fendler, ‘‘Catalysis in Micellar and
Macromolecular System,’’ Academic, New York (1975).
J. H. Fendler, ‘‘Membrane Mimetic Chemistry,’’ Wiley,
-
MnO4
+SDS
+Mn(II)
+Mn(II)
Mn(IV)
Mn(III) / Mn(IV)
+SDS
4
C. A. Bunton, F. Nome, F. H. Quina, and L. S. Romsted,
Acc. Chem. Res., 24, 357 (1991).
‘‘Physics of Amphiphiles: Micelles, Vesicles, and Micro-
+SDS
5
Mn(II)
emulsions,’’ ed by J. Degiorgio and M. Corti, North-Holland,
Amsterdam (1985).
( Product )
6
L. S. Romsted, ‘‘Surfactants in Solution,’’ ed by K. L.
Mittal and B. Lindman, Plenum, New York (1984), Vol. 2.
C. A. Bunton and G. Savelli, Adv. Phys. Org. Chem., 22,
213 (1986).
Scheme 3.
7
the autocatalytic reaction path. In order to confirm the observa-
tions of previous investigatꢁors,30–32 the kinetics of Mn(II) oxi-
dation to Mn(III) by MnO4 as studied spectrophotometrically
at 525 nm (Table 3). Based on their results, we conclude that in
presence of exꢁternally added Mn(II), the path of oxidation of
SDS by MnO4 may become more complicateꢁd (Scheme 3).
Table 1 clearly demonstrates that the MnO4 oxidizes both
the monomeric molecules of SDS as well as the aggregated
units (micelles). Surfactant monomers rapidly join and leave
micelles, and the aggregation number represents only an aver-
age over time. Micelles are not fixed entities, but have a tran-
sient character.33 Therefore, according to the multiple equilib-
rium models, the distribution of the surfactant (D1) between
various states of aggregation is controlled by a series of dy-
namic association–dissociation equilibria:
8
C. A. Bunton, ‘‘Kinetics and Catalysis in Microheterogene-
ous System,’’ ed by M. Gratzel and K. Kalyanasundaram, Marcel
Dekker, New York (1991).
9 L. Mukherjee, N. Mitra, P. K. Bhattacharya, and S. P.
Moulik, Langmuir, 11, 2866 (1995).
10 J. K. Thomas, Chem. Rev., 80, 285 (1980).
11 E. H. Cordes and R. B. Dunlap, Acc. Chem. Res., 2, 329
(1969).
12 E. H. Cordes, Pure Appl. Chem., 50, 617 (1978).
13 M. N. Khan, J. Colloid Interface Sci., 170, 598 (1995).
14 Kabir-ud-Din, J. K. J. Salem, S. Kumar, M. Z. A. Rafiquee,
and Z. Khan, J. Colloid Interface. Sci., 213, 20 (1999).
15 E. Perez-Benito and E. Rodenas, Langmuir, 7, 232 (1991).
16 K. Sen, P. S. Tribedi, and K. K. Sen Gupta, J. Surf. Sci.
Technol., 16, 220 (2000).
D1 þ D1 ꢀ D2;
D2 þ D1 ꢀ D3;
Dnꢁ1 þ D1 ꢀ Dn:
ð16Þ
ð17Þ
ð18Þ
17 Kabir-ud-Din, J. K. J. Salem, S. Kumar, M. Z. A. Rafiquee,
and Z. Khan, J. Colloid Interface Sci., 215, 9 (1999).
18 Kabir-ud-Din, J. K. J. Salem, S. Kumar, and Z. Khan,
Colloids Surf., 168, 241 (2000).
The small aggregates of the surfactant (dimers, trimers,
tetramers, etc.) exist below the cmc.33 These small submicellar
aggregates are responsible for the oxidation of SDS by
MnO4ꢁ. The equilibrium between the micelles and sub-micel-
lar aggregates is fast. In the presence MnO4ꢁ, the equilibrium
shifts towards the right-hand side because monomeric SDS is
consumed, and is oxidized to the product. On the other hand,
in the micellar pseudo phase, the aggregated SDS molecules
(Dn) are oxidized by the MnO4ꢁ. It has been proved that the
exact reaction site cannot be proposed because the micellar
pseudo phase is regarded as a microenvironment having vary-
ing degrees of water activity.34 The oxidation of SDS was
found to increase with increases in the [Hþ] value (Table 1).
HClO4 is a strong acid, which dissociates completely. There-
fore, due to the electrostatic interactions between the anionic
head group of SDS micelle (–OSO3ꢁ) and Hþ, the local con-
centration of Hþ increases in the Stern layer (water-rich region
as activity of water at the surface of ions micelles is not differ-
ent from water activities in the aqueous pseudo phase). As a
19 Kabir-ud-Din, M. Akram, and Z. Khan, Colloids Surf., A,
178, 167 (2000).
20 Kabir-ud-Din, K. Hartani, S. Kumar, and Z. Khan, Tenside
Surfactants Deterg., 38, 238 (2001).
21 K. B. Wiberg and R. Stewart, J. Am. Chem. Soc., 77, 1786
(1955).
22 J. F. Perez-Benito and C. Arias, Int. J. Chem. Kinet., 23,
717 (1991).
23 J. F. Perez-Benito, C. Arias, and E. Amat, J. Colloid
Interface Sci., 177, 288 (1996).
24 V. Pimienta, D. Lavabre, G. Levy, and J. C. Micheau,
J. Phys. Chem., 98, 13294 (1994), and the references cited therein.
25 F. Freeman and J. C. Kappos, J. Am. Chem. Soc., 107, 6628
(1985).
26 D. H. Macartney and N. Sutin, Inorg. Chem., 24, 3403
(1985).
27 Z. Khan, Raju, M. Akram, and Kabir-ud-Din, Int. J. Chem.
Kinet., 36, 359 (2004).
28 J. F. Perez-Benito and C. Arias, J. Colloid Interface Sci.,
149, 92 (1992).
29 F. Mata-Perez and J. F. Perez-Benito, Can. J. Chem., 63,
988 (1985).
30 R. T. Powell, T. Oskin, and N. Ganapathisubramanian,
J. Phys. Chem., 93, 2718 (1989).
31 D. R. Rosseinsky and M. J. Nicol, Trans. Faraday Soc., 61,
2718 (1965).
ꢁ
result,ꢁ–OSO3 converted to –OSO3H. On the other hand,
MnO4 acts as HMnO4 at the interfacial junction at the region
of the Stern and Gouy-Chapman layers; as the reaction pro-
ceeds the –OSO3H form an intermediate with the HMnO4
(Scheme 1). The probable reaction site may be the Stern and
Gouy-Chapman layers’ junctural region.
32 J. I. Morrow and S. Perlman, Inorg. Chem., 12, 2453 (1973).
33 D. O. Shah, ‘‘Micelles, Microemulsions, and Monolayers,’’
ed by D. O. Shah, Marcel Dekker, New York (1998).
34 S. Tascioglu, Tetrahedron, 52, 11113 (1996).
References
1 Z. Khan, Raju, and Kabir-ud-Din, Colloids Surf., A, 225,
75 (2003).