Anionic-micelles assisted oxidation of tartaric acid by permanganate: A kinetic and mechanistic approach

Article abstract of DOI:10.1016/j.molliq.2016.12.081

In this study, we demonstrate the catalytic effects of sodium dodecylbenzenesulphonate (SDBS) on the oxidation of tartaric acid by permanganate. UV–visible spectroscopic data revealed that the reaction proceeds through the formation of a transient species between the reactants. The values of ??ε?=?12.8?×?10⁵?mol^??1?dm³?cm^??1 and K_c?=?6.5?×?10^??4?mol^??1?dm³ were calculated from the kinetic data. Reaction-time profiles show that the reaction has an induction period (slow step) followed by an auto-acceleration path (fast step). SDBS is found to be stable in presence of MnO₄^?. The reaction follows first-order kinetics with respect to tartaric acid under pseudo-first-order conditions in absence and presence of SDBS. The observed results are discussed in terms of pseudo-phase model of the micelles proposed by Menger and Portony. The values of activation energy?=?99?kJ?mol^??1 and 59?kJ?mol^??1 were calculated by using Arrhenius equation, respectively, in absence and presence of SDBS micelles. The probable mechanisms have also been proposed and discussed. Presence of silver ions also catalyzed of reaction rates.

Full text of DOI:10.1016/j.molliq.2016.12.081

Anionic-micelles assisted oxidation of tartaric acid by permanganate: A
kinetic and mechanistic approach
Zoya Zaheer, Elham Shafik Aazam
PII:
DOI:
Reference:
S0167-7322(16)33348-7
doi:10.1016/j.molliq.2016.12.081
MOLLIQ 6764
To appear in:
Journal of Molecular Liquids
Received date:
Revised date:
Accepted date:
26 October 2016
16 December 2016
22 December 2016
Please cite this article as: Zoya Zaheer, Elham Shafik Aazam, Anionic-micelles assisted
oxidation of tartaric acid by permanganate: A kinetic and mechanistic approach, Journal
of Molecular Liquids (2016), doi:10.1016/j.molliq.2016.12.081
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Anionic-micelles assisted oxidation of tartaric acid by permanganate: A kinetic
and mechanistic approach
Zoya Zaheer *, Elham Shafik Aazam
Department of Chemistry, Faculty of Science, King Abdulaziz University, P.O. Box
80203, Jeddah , 21589, Saudi Arabia
Abstract In this study, we demonstrate the catalytic effects of sodium
dodecylbenzenesulphonate
(SDBS) on the oxidation of tartaric acid by
permanganate. UV-visible spectroscopic data revealed that the reaction proceeds
through the formation of a transient species between the reactants. The values of ∆ε
= 12.8 × 105 mol-1 dm3 cm-1 and Kc = 6.5 × 10-4 mol-1 dm3 were calculated from the
kinetic data. Reaction-time profiles show that the reaction has an induction period
(slow step) followed by an auto-acceleration path (fast step). SDBS is found to be
-
stable in presence of MnO4 . The reaction follows first-order kinetics with respect to
tartaric acid under pseudo-first-order conditions in absence and presence of SDBS.
The observed results are discussed in terms of pseudo-phase model of the micelles
proposed by Menger and Portony. The values of activation energy = 99 kJ mol-1
and 59 kJ mol-1 were calculated by using Arrhenius equation , respectively, in absence
and presence of SDBS micelles. The probable mechanisms have also been proposed
and discussed. Presence of silver ions also catalyzed of reaction rates.
Key words: Micellar effects; Kinetics; Tartaric acid; Permanganate
*Correspondence to: E-mail : zoya.zaheer@gmail.com (Z. Zaheer).

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1.Introduction
Surfactants are usually organic compounds that are amphiphilic, which exhibits a
double affinity as a polar- a polar duality and contain both hydrophobic and
hydrophilic groups at their tails and heads [1-3]. They will diffuse in water and adsorb
at interfaces between air and water. The water-insoluble hydrophobic group may
extend out of the bulk water phase, while the water-soluble head group remains in the
water phase. The behavior of surfactants in solution were the subject of considerable
interest because of its wide applications, in different industries, medical sciences, life
sciences, analytical chemistry, pollution control and even in the fields of physics and
engineering. They behaves as a electrolyte at lower concentration and formed
aggregates due to unusual property of self-assembly at certain concentration. It has
been established that the rate enhancement of a chemical reaction in solution by the
addition of a surfactant at a concentration higher than its critical micelle concentration
may be due to higher concentration of the reactants in the aggregates of surfactant [4-
6]. The utilization of surfactants as reaction media affects reaction rates, yield of
products, mechanism, and stereochemistry; they may also be useful in chemistry for
kinetic studies [7]. The hydrophobic, electrostatic, and hydrogen bonding forces plays
an important role in the solubilization of water insoluble organic substance into the
micelles.
Menger and Portony [8], Bunton et al.[9, 10], Vera and Rodenas [11], Moya et al.
[12], Tuncay et al.[13] , and others developed various models to explain the kinetic
and mechanistic aspects of various micellar catalyzed and inhibited reactions. Bunton
in his pioneering review discussed the mechanism of redox reactions on the basis of
micellar kinetics [14]. Khan reported the effects of anionic micelles on the intra
molecular general-base catalyzed hydrazinolysis and hydrolysis of phenyl salicylate.

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They also suggested that the micelle behaves as a porous cluster having water
cavities with rough surface [15]. The silent properties of the surfactants that affect
electron transfer reactions are localization and compartmentalization effect, pre-
orientational, polarity and counter ion effect and the effect of charged interfaces [10,
16]. Tartaric acid is a dihydroxyl derivative of succinic acid, which has been used as a
antioxidant, leavening-, and chelating-agent in the field of food preparation,
pharmaceuticals, wine making industries [17]. They also formed effervescent salts
and metal complexes in presence of citric acid, and various metal ions, respectively,
which improved the taste of oral medications, and prevent the metal ions from the
corrosion. The stability and redox chemistry of tartaric acid have significant impacts
in the food making processes. The kinetics and mechanism of amino acids, carboxylic
acids (citric, succinic, tartatic and oxalic acids) and drugs containing these functional
-
groups with different oxidants (MnO4 , V(V), and Cr(VI)) in presence of cationic
cetyltrimethylammonium bromide (CTAB), cetylpyridinium chloride (CPC) anionic
sodium dodecylsulphate (SDS), and nonionic TrotonX-100 surfactants have been the
subject of various investigators [18-26], but the details of systematic kinetics of
permanganate ion (strong oxidizing agent in acidic, neutral, and alkaline media, its
reduction products depends on the experimental conditions) oxidation of tartaric acid
are not yet known in presence of anionic SDBS surfactant. In the present
investigation, we examine the effects of sodium dodecylbenzene sulfate on the
-
oxidation kinetics of tartaric acid-MnO4 redox system for the first time. We
determine the various parameters associated with an aqueous and micellar chemistry.
2.Experimental and methods
2.1.Materials
Tartaric acid (99%, C4H6O6; molar mass 150.087; reductant; BDH England),
potassium permanganate (99%, KMnO4; oxidant; molar mass 158.034 ; BDH

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England), Dodecylbenzenesulphonic acid sodium salt (99 % ; C18H29NaO3S; molar
mass 348.48; BDH England), NaClO4 ( 99 %; molar mass 122.44 ; BDH England)
and inorganic salts were used as received. The water used as the solvent was
previously subjected to deionization, followed by double distillation (first time from
alkaline KMnO4). The purity of SDBS was ascertained by the reported method [27].
2.2.Kinetic procedure
In a typical experiment, all reactants solution (except tartaric acid) were taken in a
three necked reaction flask fitted with a double walled spiral condenser ( to arrest
evaporation) and immersed in an thermostatic water bath maintained at the desired
temperature. The equilibrated tartaric acid solution was added in the reaction flask ,
-
and progress of the reaction (decay of MnO4 absorbance ) was followed by
measuring the absorbance at definite time intervals at 530 nm on a Spectro UV-vis
double beam PC scanning spectrophotometer, Model UVD-2960, Labomed. It is to be
noted that the absorbance remains constant for ca. 40 min and than decreases with
time very fast, which might be due to the two states kinetics of tartaric acid oxidation,
i.e., the time up to linearity can be considered to be the first stage and after this point,
the deviation in linearity could be called the second stage. For the first step, the
pseudo-first-order rate constants with and without SDBS were calculated from the
slopes of the tangents to the plots of log(absorbance) versus time with a fixed-time
method for all kinetic experiments [28, 29]. Duplicate runs gave results that were
reproducible to within ± 4 %. The pH of the reaction mixtures were measured just
after mixing of reactants and completion of the reaction. No interference due to other
species at 410 nm for Mn(IV) was verified.
2.3.Critical micelle concentration (CMC) determination

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-
In order to determine the CMC of SDBS solution with and without MnO4 and tartaric
acid , the specific conductance of different solutions (SDBS + H2O, SDBS + tartaric
-
acid, and SDBS + MnO4 ) were measured on a Microprocessor conductivity meter of
o
WTW, Model LF/2000C (Germany) at 25 C. The specific conductance was recorded
-
as a function of [SDBS] at a fixed [tartaric acid] and [MnO4 ]. The CMC values were
determined from plots of the specific conductivity versus [SDBS]. The break point of
nearly two straight-line portions in the plot are taken as an indication of micelle
formation and this corresponds to the CMC of surfactant. The results are summarized
in Table 1.
2.4.Products analysis
-
In a typical experiment, a solution of [MnO4 ] =1.0 ×10-4 mol dm-3 was added to a
reaction mixture containing [tartaric acid] = 10.0 ×10-4 mol dm-3 with or with out ,
SDBS = 10.0 ×10-4 mol dm3. As the reaction proceeded, the evolved carbon dioxide
was flushed out by passing a continuous current of pure N2 gas (free from O2 and
CO2) and was absorbed in standard Ba(OH)2 solution and then [CO2] estimated
quantitatively by titration with standard HCl solution [30]. On the other hand, after
completion of the reaction, the resultant mixture was collected to a round-bottom
flask and distilled. The distillate was collected in a saturated solution of 2, 4-
dinitrophenylhydrazine in 2N HCl. We did not observed the formation of yellow
precipitate . Thus ruled out the possibility to the formation of any type of oxidation
product of tartaric acid containing aldehyde moiety.
3.Results and discussion
3.1.Aqueous chemistry
-
In a typical experiment, unusual kinetic behavior was observed when MnO4 (1.0 ×
10-4 mol dm-3) was allowed to react with tartaric acid ( 10.0 × 10-4 mol dm-3) .

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Absorbance of the reaction mixture increases initially with reaction time at fixed
[tartaric acid] (Fig. S1; supporting information ) at 530 nm and than decreases as the
-
reaction proceeds, which might be due to the initial complexation between MnO4 and
tartaric acid (Eq. 2) [31].
KC
MnO4- + tartaric acid
complex
(1)
The value of Kc was calculated by using the Eq. (2).
-
[MnO ]
[tartaric acid]
[tartaric acid]
1
4
=
+
(2)
Kc
(absorbance)
l

l
-
where, ∆(absorbance) = difference in absorbance between the complex and MnO4 at
-
525 nm when both the un-complex and complexed forms of MnO4 absorb. ∆ε =
difference in absorption coefficients and l is the path length (1.0 cm). According to
-
Eq. (2), the plot of [MnO4 ]T [tartaric acid] / ∆(absorbance) against [tartaric acid]
should be linear and this was found to be so (Fig. S2; supporting information) with
positive intercept on the Y-axis. The values of ∆ε = 12.8 × 105 mol-1 dm3 cm-1 and
Kc = 6.5 × 10-4 mol-1 dm3 were calculated from the intercept and slope of Fig.S2.
-
In order to determine the reaction kinetics with respect to [MnO4 ], and [tartaric
acid], a series of experiment were performed under different experimental conditions.
The results are summarized in Tables 1 to 3. As can be seen in Fig. 1 (typical
-
example), absorbance of MnO4 remain unchanged for some time ca. 40 min and than
decreases with reaction-time , indicating the formation of intermediate having higher
reaction rates in comparison to tartaric acid. Finally, the pink -purple color became
transparent colorless after ca. 90 min. Reaction-time curves (typical example), plots
-
of absorbance versus time clearly indicate that the oxidation of tartaric acid by MnO4
has an induction period (absorbance remains constant for ca. 40 min) followed by an
-
auto-acceleration (fast disappearance of MnO4 color), which might be due to the

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oxidation rates of intermediate(s) (oxidation product of tartaric acid) along with the
parent reaction. Table 2 shows that all kinetic runs with an excess of tartaric acid
-
exhibit pseudo-first-order kinetics with respect to [MnO4 ], indicating that the total
-
[MnO4 ] must be used in the rate equation (Eq. 3).
-
-
- d[MnO4 ] / dt = kobs [MnO4 ]T
(3)
The effect of varying [tartaric acid] (from 1.0 × 10-4 mol dm-3 to 50.0 × 10-4 mol dm-
-
3) on the reaction rate was studied at constant [MnO4 ] =1.0 × 10-4 mol dm-3 at room
temperature. The kobs values are summarized in Table 2. Double log-log plot between
kobs and [tartaric acid] is a straight line (with slope = 0.98), indicating the first-order
with respect to [tartaric acid]. These results are in good agreement to the observations
of Hassan and his co-workers regarding the first-order kinetics with respect to tartaric
acid [21]. Some kinetic experiments were also performed to determine the role of [H+]
(HClO4 was used as a source of hydrogen ions concentrations; from 1.0 × 10-3 mol
-
dm-3 to 10.0 × 10-3 mol dm-3) on the reaction rate at fixed [MnO4 ] =1.0 × 10-4 mol
dm-3 and [tartaric acid] = 20.0 × 10-4 mol dm-3. The values of rate constant increases
with increasing [H+] and the reaction followed first-order kinetics. It is known that
tartaric acid is a dibasic weak acid (pK1 = 3.03, and pK2 = 4.34), exists in the various
species in aqueous solution and the ratio of these species depends upon the pH of the
working solution. Thus, under our experimental conditions (pH = 2.8), unionized
form of tartaric acid to be the reactive and major existing species. Hassan [21] and
Hassan et al. [25] reported that the ionic-strength has significant effects on the redox
-
kinetics of MnO4 with alanine, and tartaric acid, respectively. Therefore, to examine
the influence of ionic strength on the reaction rates, a series of kinetic experiments
were carried out in NaClO4 solutions of 1.0 × 10-2 - 4.0 × 10-2 mol dm-3 at constant

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-
[MnO4 ] =1.0 × 10-4 mol dm-3 and [tartaric acid] = 20.0 × 10-4 mol dm-3. The kobs
values were found to constant with increasing ionic strength (Table 2), indicating the
reaction might occur between two dipolar neutral molecules [21]. Based on the
observed facts, a straightforward reaction mechanism is put forward which involve a
complex formation between reactive species of tartaric acid and MnO4 − (Scheme 1).
−
Scheme 1. Mechanism to the oxidation of tartaric acid by MnO4
In Scheme 1, Eq (4) is the ionization of tartaric acid. Fig. S1 shows that the reaction
−
proceed through a complex formation between the reactants. By analogy with MnO4
ion oxidation of other organic reductants [32, 33], we assume that reaction proceeds
-
through the formation of an ester like species between MnO4 and tartaric acid (Eq
−
(5)). The complex has smaller molar absorptivity as compared to that of MnO4 ion
at 525 nm. In the rate determining step (Eq. (6)), the ester- complex undergoes
oxidative decomposition leading to the formation of Mn(V), CO2 and an
intermediate. In the next step, Mn(V) reacts with other molecule of tartaric acid and
give Mn(III) (Eq. (7)). Finally, Mn(III) further reacts with tartaric acid and
immediately gets converted in to the stable species of Mn, i.e., Mn(II) and a radical
(Eq. 8). The formation of free radical during the redox process was identified as
follows. Addition of a saturated solution of HgCl2 and/or acrylonitrile to a reaction
mixture containing tartaric acid and permanganate, formation of white precipitate was
observed after some times, indicating the free radical intervention in the oxidation
reaction.
On the other hand, unstable oxidation product of tartaric acid (Scheme 1;
-
intermediate) further react with MnO4 , leads to the formation of CO2 as a final

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oxidation product of tartaric acid. Mn(V) species is highly unstable in aqueous acidic
medium, and immediately converted into the MnO2 (generally formed as an
-
intermediate in the MnO4 redox reactions with organic reductants. Its stability
strongly depends on the pH and nature of reducing agents. In order to establish the
-
formation of MnO2, spectra of reaction mixture containing MnO4 and tartaric acid
were recorded at different time intervals. Inspection of Fig. S3 (supporting
-
information) shows that the absorbance decreases with time and MnO4 pink color
disappeared completely at the end of the reaction. We did not observed the formation
of stable and/or unstable yellow color . Thus ruled out the formation of colloidal
MnO2 as an stable intermediate during the course of reaction (Scheme 2).
Scheme 2. Mechanism to the oxidation of intermediate , oxidation product of
tartaric acid
Schemes 1 and 2 clearly suggests that there is a competition between, Mn(V), MnO2,
-
Mn(III), intermediate(s) and radical to react with tartaric acid and/or MnO4 . There is
an involvement of a series of cyclic reactions (Eqs. 7, 8 , and Scheme 2). Finally, we
-
can state that, when an excess of MnO4 over tartaric acid is used, further oxidation of
the intermediate(s) occurs to yield carbon dioxide as a final product. Under these
conditions the following stoichiometric equation (Eq. 9) can be proposed, which is a
-
sum of the oxidation of tartaric acid and its possible oxidation product(s) by MnO4 .
HO-CH-COOH
HO-CH-COOH
2MnO4- +
+ 6H+
2Mn2+ + 4CO2 + 6H2O
(9)
In order to determine the activation parameters, the reaction was also carried out at
different temperatures in the range of 25 to 35 οC . The Arrhenius plot of log(kobs)

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versus 1/T was used to the determination of activation energy. The enthalpy and
entropy of activation (∆H# and ∆S# 1) were also calculated by using Eyring equation.
Table 1 record the values of activation parameters along with the values of CMC.
k = (kBT / H) exp ( S#/R) exp(- H#/RT)
(10)
3.2.Micellar chemistry
It has been established that sodium dodecylsulphate, and Triton X-100 were unstable
-
with MnO4 in presence of acidic media [34, 35]. Therefore, a series of spectra of
-
MnO4 were recorded with SDBS as a function of time. Fig. 2 shows that the spectra
-
of MnO4 remain unchanged for ca. 100 min under our experimental conditions,
-
suggesting that SDBS is stable. Initially, oxidation kinetics of tartaric acid-MnO4
reaction was studied at several initial concentrations of all the reactants in the
presence of SDBS (Table 2). It is to be mentioned that pH and ionic strength of the
working solutions could not be maintained constant for the micelles mediated
reactions. Tondre and his co-workers [36, 37] have advised to avoid the use of even
buffer solutions to maintain pH of micellar solutions (pH shift by about minus two
units is expected at the surface of SDBS micelles with respect to the bulk pH and
control of pH is not as straightforward in micellar solutions as in ordinary solvents
[38]) Therefore, no mineral acid and electrolyte was used to maintain the pH , and
ionic strength constant in presence of SDBS. In order to establish the pre-, and post-
-
micellar effect of anionic micelles of SDBS on the MnO4 -tartaric acid redox reaction
, the reaction rates were determine as a function of [SDBS] (from 0.0 × 10−4 to 30.0
−
×10−4 moldm−3) at constant [MnO4 ] = 1.0 ×10−4 mol dm−3,and [tartaric acid] = 20.0
×10−4 mol dm−3 at room temperature. The observed results are depicted graphically as
reaction-time profiles and kobs–[SDBS] in Figs. 3 and 4, respectively. It has been
found that kobs increases (nearly 19-fold) steeply by increasing [SDBS] from 0.0 to
20.0 ×10-4 mol dm-3 (Fig.4). Surprisingly, reaction was completed, 100 %
−
disappearance of MnO4 pink color with in 50 min at lower [SDBS] , where as
reaction became slow at higher [SDBS]. The decrease in reaction rates absorbance at
higher [SDBS] might be due to the dilution effect. Similar behavior has been observed
by Hassan [21] in the oxidation of tartaric acid in the presence of non-ionic Tritin X-
100 surfactant. The pre-micellar SDBS kinetic effects as well as the micellar SDBS

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effects on oxidation rate can be rationalized by considering the distribution and/or
solubilization of the reactants and/or intermediate(s) among the different pseudo
−
phases present in the reaction medium. The order of reaction with respect to [MnO4 ]
was found pseudo-first-order kinetics which shows that rate does not depend upon the
−
initial [MnO4 ] (Table 2). The above studies confirm that the order of reaction with
−
respect to [ MnO4 ] and [tartaric acid] is the same as that in aqueous medium.
Organic additives and inorganic electrolytes might produce significant alterations in
the formation of surfactant aggregates. Generally , the CMC value of a particular
surfactants decrease in presence of inorganic inert salts and organic additives at a
specific temperature, which might be due to the interaction, incorporation, and
solubilization of additives into the micelles. This causes also changes in the water
structure [39]. Several techniques have been developed to the determine the CMC of
surfactants. In the present studies, the CMC value in water estimated from the specific
conductance plots versus [SDBS] (Fig.5) is 3. 7 ×10-4 mol dm-3 for pure SDBS,
which is in good agreement with the literature values (Table 1). Fig. 5 also shows that
-
the MnO4 and tartaric acid have the ability to reduce the CMC.
3.3.Probable role of SDBS
It has been established that pre- , and post- micellar effects of surfactants on the
reaction rates might be due to the formation of some physical active entities and
incorporation or exclusion of the reactants from aqueous micellar phase, respectively
[40]. The hydrophobic, electrostatic and hydrogen bonding are the main factors
involved in the rate enhancement of a bimolecular reaction. As a typical example, Fig
4 shows that the reaction rate increases with [SDBS] (≥ 2.0 ×10-4 mol dm-3).
Accordingly to the multiple equilibrium model, the distribution of surfactant (D1)
between various states of aggregation is controlled by a series of dynamic association-
dissociation equilibria [41].

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D2
D1 + D1
(11)
(12)
(13)
D3
Dn
D1 + D2
Dn -1 + D1
The pre-micellar catalytic effects of SDBS can be sought in the fact that small
aggregates of the SDBS exist below the CMC; these small sub micellar aggregates
(monomers, dimmers, trimers, etc ) can interact physically with the reactants forming
active entities [40]. Therefore, ionic micelles behaving as electrolytes, exert kinetic
salt effects and affects reaction rates in the aqueous region [5]. The polarity of the
Stern layer (most of the ionic micelle mediated organic reactions are believed to occur
in this region]) is generally considered to be lower than the bulk water and micellar
surfaces are water-rich and do not provide a uniform reaction medium because a
micelle is a porous cluster with a rough surface and deep-water-filled cavities. In
micelle-mediated reactions, it is not possible to precisely locate the exact site of the
reaction but, at least, localization of the reactants can be considered. Due to the
presence of water in the micellar Stern layer, the solubilization of tartaric acid and
permanganate into the micellar pseudo phase can not be completed ruled out. On the
other hand, as suggested by the Bunton that the consequent charge redistribution
(charge-charge interaction) in formation of the transition state can not be ruled out
completely in the present redox reaction [14]. A possible arrangement (although
highly schematic) could be that as shown in Scheme 3.
Scheme 3. Solubilization of reactants and intermediate into the SDBS micelles
The observed catalytic role of SDBS micelles role may be explained in terms of the
pseudo phase model of micelles proposed by by Menger-Portony [8] modified by
Bunton for the bimolecular reaction (Scheme 4) [40]. The assumptions involved in

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this model, the advantages, and the disadvantages are critically discussed by Bunton
[5, 40]. As the rates were observed to increase with [SDBS], it is likely that both the
reactants are partitioned between aqueous and the micellar pseudo-phases.
Scheme 4. Micellar pseudo-phase model for the solubilization of reactants
Bunton considered the total volume of the micelles as a separate phase and both the reactants
are associated/ incorporated into micelles and for bimolecular reaction, Eq. [14] can be
derived from Scheme 4 [5].
kw [TA]T + ( Ks km - kw) MTS [Dn]
kobs
(14)
(1 + Ks [Dn])
S
where TA = tartaric acid, kw = k’w / [( TA )w] and km = k’m MT ; kw and km are the second -
S
order rate constants and MT = mole ratio of TA associated to the micellar head group, i.e.,
S
[(TA)m] / [Dn]). The value of MT was obtained by considering the equilibrium:
KT
(TA)w + Dn
(TA)m
(15)
(16)
[(TA)m]
[(TA)m] ( [Dn] - [(TA)m] )
KT
A quadratic Eq. (17) can be obtained with the help of mass balance of tartaric acid, [TA]T =
[(TA)w] + [(TA)m].
KT [(TA)w ]2 - ( 1 + KT [Dn] + KT [TA]T ) [(TA)m] + KT [Dn] [TA]T = 0
(17)
which is solved for [(TA)m] with the help of a computer programme [6], where KT as an
2
adjustable parameter . The value of KT for which the value of Σdi turned out to be minimum
was taken as the best value of KT [42]. The values of kw = 2.2 ×10-4 s-1, KS = 50 mol-1 dm3,
km = 10.2×10-4 s-1 , KT = 70 mol−1 dm3, kw = 11.2 ×10-2 mol-1 dm3 s-1(second-order rate
constant in ansence of SDBS), k2m = 1.8 ×10-4 mol1 dm3 s-1 ( based on the relation: k2
=
m

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14
km Vm, where Vm is the molar volume of SDBS micelles), and kw / k2 = 622 were
m
calculated by a non linear least square technique. These values are also substituted in Eq.(14),
and kcal are calculated (Table 2), which are in good agreement with the observed values of
rate constants, provides the supporting evidence to the validity of Scheme 3 mechanism.
Table 1 shows that the Ea for the SDBS catalyzed reaction is lower in comparison to
the un-catalyzed reaction, indicating that SDBS acts as a catalyst and provide a new
reaction path, which has lower activation energy (Scheme 5). A complete account for
all the factors that influence ∆H# and ∆S# is not possible because kobs does not
represent a single elementary step as it is a complex function of association, ionization
and micellization constants. Further more though a change in temperature is known
produce changes in the size, shape, surface charge, ect. of the micelles and micellar
structural sensitivity to temperature is kinetically unimportant.
Scheme 5. Energy profile diagram for the reaction
3.4.Comparison with other reductants and effects of [Ag+] on reaction rates
In order to compare the reactivity of tartaric acid with other related organic acids,
oxalic and succinic acids, A series of kinetic experiments were performed under the
same experimental conditions to that of tartaric acid. Under our experimental
-
conditions, [MnO4 ] = 1.0 × 10−4 mol dm−3 and [organic acids] = 20.0 × 10−3
moldm−3 at room temperature, no oxidation was observed, even after prolonged
incubation, whereas the oxidation of tartaric acid is very fast in these experimental
conditions and completed with the 1h (Fig. 3) in presence of SDBS. It has been
established that the presence of [H+] is essential for the oxidation of oxalic and
-
succinic acids with MnO4 [43]. Therefore, the reactivity of tartaric , oxalic and
succinic acids is in the order tartaric » oxalic » succinic, which could be explained in

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terms of the presence of -OH groups [44]. To see insight into the effect of AgNO3 ,
some kinetic experiments were carried out (Fig. 1), which suggested that Ag+ ions has
catalytic role on the oxidation of tartaric acid. These observations are in good
agreement to the results of Verma et al. regarding the catalytic effect of Ag+ ions on
the permanganate oxidation of amino acids [45].
Conclusions
In this study, we have observed the sub-, and post- micellar catalytic effects of
-
SDBS aggregates (monomer, dimmer, trimmer etc.) on the MnO4 - tartaric acid redox
reaction in an aqueous media. We are unaware of any precedence in the redox
chemistry of this system, especially regarding the role of SDBS. The reaction-time
profiles showed that reaction has an induction period followed by an auto-acceleration
path, which might be due to the fast oxidation rates of tartaric acid oxidation
-
product(s) with MnO4 . To the best of our knowledge, this is the first evidence
showing the pre- and post-micellar effect of SDBS in the oxidation of tartaric acid.
The present work provides new insight that may allow a better understanding of the
electron transfer reactions occurring during the use of tartaric acid as a preservative in
the food chemistry.

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Table 1. Critical micelle concentration of SDBS different experimental conditions,
-
rate constant, activation parameters for the oxidation of tartaric acid by MnO4 .
Solutions
104 cmc (mol dm-3)a Temp (K0) 104kobs (s-1)
Aqueous
Micellar
18.8
30.3
41.2
59
SDBS + Water 3.7
298
303
308
2.2
6.1
8.2
-
SDBS + MnO4 2.2
SDBS + TA
Literatureb
3.1
Activation parameters
3.6 to 16.0
Ea (kJ mol-1)
99
∆H# (kJ mol-1) 96
56
∆S# (J K-1 mol-1) -103
-109
-
a[MnO4 ] = 1.0 × 10-4 mol dm-3; [tartaric acid = 10.0 × 10-4 mol dm-3; Temp. = 25 °C.
b K.G. Furton, and A. Norelus, J. Chem. Ed. , 1993, 70, 254-.256; S.S. Shah, K.
Naeem, S.W.H. Shah, G.M. Laghari, Colloids Surf. A 2000, 168, 77-85; E.W.
Karine, and G.S. A. Savall, J. Appl Electrochem. 2007, 37, 1337-1344.

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Table 2. Values of rate constants, and pH for the oxidation of tartaric acid by
- a
[MnO4 ]
-
104[MnO]4 ] 104[tartaric acid] 102 [NaClO4] pH
104kobs
( s-1 )
(mol dm-3)
(mol dm-3)
(mol dm-3)
1.0
1.5
2.0
2.5
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
1.0
20.0
20.0
20.0
20.0
10.0
30.0
40.0
50.0
20.0
20.0
20.0
20.0
20.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
0.0
1.0
2.0
3.0
3.5
4.0
2.8 (3.1) 2.0 (18.2)
2.8 (3.1) 2.1 (18.1)
2.8 (3.0) 2.0 (18.0)
2.9 (3.2) 2.1 (18.1)
2.8 (3.1) 1.6 ( 16.2(
2.7 (3.2) 4.8 (21.1)
2.9 (3.1) 7.9 (24.2)
2.8 (3.2) 8.4 (26.2)
2.8
2.8
2.9
2.8
2.8
2.2
2.1
2.0
2.2
2.0
aValues of pH and rate constants in presence of [SDBS] = 10.0 ×10-4 mol dm-3 are
given in the parenthesis at room temperature (298 K).

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21
-
Table 3. Effect of [SDBS] on the oxidation rates of tartaric acid by [MnO4 ]
104[SDBS]a
(mol dm-3)
pH
104 kobs 104kcal
kobs- kcal / kobs
s-1
s-1
0.0
2.5
5.0
3.2
3.2
3.1
3.2
3.1
3.2
3.1
3.2
3.1
3.2
2.0
4.2
8.0
16.1
18.2
16.0
14.0
10.1
10.0
10.1
7.5
10.0
12.5
15.0
20.0
25.0
30.0
18.1
16.1
15.2
10.0
+0.005
-0.003
-0.085
+0.009
-
a[tartaric acid] = 20.0 ×10-4 mol dm-3, [MnO4 ] = 1.0 ×10-4 mol dm-3 at room
temperature (298 K).

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22
Scheme 1
Ka
-
HO-CH-COOH
HO-CH-COO- + H+
HO-CH-COOH
(4)
(5)
MnO4
HO-CH-COOH
(TA)
-H+ ++H
KC
HMnO4 + HO-CH-COOH
HO-CH-COOH
O=Mn-O-CH-COOH + H2O
HO-CH-COOH
(ester-complex)
k, rds
ester-complex
HMnO3 + CO2 + CH-CH-COOH
(6)
OH
(intermediate)
Mn(III) + intermediate
fast
fast
Mn(V) + TA
Mn(III) + TA
(7)
(8)
Mn(II) + HO-CH-COOH
.
O-CH-COOH
(radical)
Ka Kc k1[tartaric acid]T [H+]
(H+ + Ka)
kobs
=

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23
Scheme 2
fast
-
CH-CH-COOH + MnO4
HOOC-CH-COOH + Mn(V)
OH
OH
(tartronic acid)
fast
fast
fast
-
tartronic acid + MnO4
3CO2 + Mn(V)
2HMnO3
MnO2 + H2MnO4
Mn(II) + intermediate
Tartaric acid + MnO2

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24
Scheme 3
-
MnO4
H+
HOCHCOOH
HOCHCOOH
HMnO4
H+
-
-
-
-
-
HOCHCOOH
HOCHCOOH
O=MnOCHCOOH
-
HOCHCOOH
(Ester complex)
-
Na+
-
-
H+
Na+
-
-
H+
Na+
H+
-
-
Na+
Na+
-
Na+
-
-
Stern layer
+
HOCHCOOH
HOCHCOOH
H

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25
Scheme 4
(MnO4-)w
+ Dn
+ (Tartaric acid)w
kw
Ks
KT + Dn
Products
-
(MnO4 )m + (Tartaric acid)m
km

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26
Scheme 5

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27
Figure captions
Fig.1. Reaction -time curves showing the effects of [tartatric acid] and [Ag+] on the
-
oxidation of tartaric acid by [MnO4 ] = 1.0 ×10-4 mol dm-3.
-
Fig.2. UV-visible spectra of reaction mixture containing MnO4 + SDBS at different
-
time intervals. Reaction conditions: [MnO4 ] = 1.0 ×10-4 mol dm-3, [SDBS ] = 10.0
×10-4 mol dm-3.
Fig.3. Reaction -time curves showing the effects of [SDBS] on the redox reaction of
-
[tartaric acid] = 20.0 ×10-4 mol dm-3 by [MnO4 ] = 1.0 ×10-4 mol dm-3.
-
Fig.4. Effects of [SDBS] on the reaction rates of MnO4 + tartaric acid redox reaction.
-
Reaction conditions: [MnO4 ] = 1.0 ×10-4 mol dm-3, [tartaric acid] = 20.0 ×10-4 mol
dm-3.
Fig. 5. Plots of specific conductance versus [SDBS].

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28
Fig.1.
0.18
0.16
0.14
0.12
0.10
0.08
0.06
0.04
0.02
0.00
103[TA] (mol dm-3)
+
10, Ag = 0.0
+
20, Ag = 0.0
+
30, Ag = 0.0
+
20, Ag = 10
+
20, Ag = 20
0
20
40
60
80
100
120
Time (min)

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Fig.2.
Time (min)
0.20
0.15
0.10
0.05
0.00
only MnO4-
10
20
30
40
50
60
90
100
120
150
180
300
400
500
600
700
800
900
Wavelength(nm)