Submicellar catalytic effect of cetyltrimethylammonium bromide in the oxidation of ethylenediaminetetraacetic acid by MnO4−

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Colloids and Surfaces B: Biointerfaces 64 (2008) 42–48

Submicellar catalytic effect of cetyltrimethylammonium bromide in the oxidation of ethylenediaminetetraacetic acid by MnO4− Maqsood Ahmad Malik, Zaheer Khan ∗ Department of Chemistry, Jamia Millia Islamia (Central University), Jamia Nagar, New Delhi 110025, India Received 26 November 2007; received in revised form 1 January 2008; accepted 5 January 2008 Available online 16 January 2008

Abstract The effects of cetyltrimethylammonium bromide (CTAB), sodiumdodecyl sulphate (SDS) and Triton X-100 (TX-100) on the oxidative degradation of ethylenediaminetetraacetic acid (EDTA) by MnO4 − have been studied spectrophotometrically at 525 and 420 nm, respectively. It was found that cationic surfactant catalyse the reaction rate while anionic and non-ionic have no effect. The premicellar environment of CTAB strongly catalyses the reaction rate which may be due to the favorable electrostatic binding of both reactants (MnO4 − and EDTA) with the positive head groups of the CTAB aggregates. The influence of different parameters such as [MnO4 − ], [EDTA], [H+ ] and [surfactants] were also considered. The reaction follows the first- and fractional-order kinetics with respect to [MnO4 − ] and [EDTA]. The proposed mechanism and the derived rate law are consistent with the observed kinetics. © 2008 Elsevier B.V. All rights reserved. Keywords: Premicellar catalysis; CTAB; SDS; EDTA; MnO4 −

1. Introduction Surfactants are amphipathic molecules having both hydrophobic and hydrophilic properties. Surfactants properties have attracted growing attention for use in biochemistry, biological and chemical research applications [1]. Researches on surfactant behavior are completely multidisciplinary in nature. The investigation of electron transfer processes in organized molecular assemblies (e.g., micelles) has added a new dimension to biochemical research [2–5]. Micellar catalysis has received considerable attention in view of analogies drawn between micellar and enzyme catalysis [6–8]. The salient 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 [9–11]. The chemical literature contains abundant reports aimed towards understanding the mechanism of manganese(VII) oxidation of S-, O- and N-containing inorganic and organic ∗ Corresponding author. Present address: Department of Chemistry, Faculty of Science, King Abdul Aziz University, P.O. Box 80203, Jeddah 21589, Saudi Arabia. E-mail address: [email protected] (Z. Khan).

0927-7765/$ – see front matter © 2008 Elsevier B.V. All rights reserved. doi:10.1016/j.colsurfb.2008.01.003

reductants in acidic neutral and alkaline media [12–16]. Kinetic and mechanistic studies of the oxidation of biomolecules, such as hydroxy acids (lactic, oxalic, malic) [16–18], ascorbic acid [19], amino acids (cysteine, glycine, methionine) [20], fructose [21] and paracetamol [22]) by water-soluble colloidal MnO2 have carried out to understand the role of MnO2 sols. Ethylenediaminetetraacetic acid (EDTA) (multifunctional ␣amino acid) forms complexes with a large number of cations, including those of ions of the main-group metals. The complex formed by calcium with EDTA is used to treat lead poisoning. Therefore, its susceptibility to biodegradation during wastewater treatment and in the aquatic environment is an important criterion for assessing its environmental impact and toxicity. Oxidation of EDTA by permanganate under different experimental conditions (acidic and alkaline media) has been the subject of several investigators [23,24]. Surprisingly, despite a large body of information being available on the kinetic and mechanistic aspects of micellar catalysis, studies of their effects upon redox reactions of EDTA have not attracted due attention. For this reason we have performed kinetic studies of the EDTA–MnO4 − reaction in presence of three surfactants (cationic, anionic and non-anionic). These studies are useful when discussing the effects of micelles on electron transfer reactions. Therefore, in this paper we wish to report the results of the

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oxidative degradation of EDTA by permanganate in presence of surfactants.

Table 1 Critical micelle concentration (cmc) values of CTAB and SDS surfactants in the absence and presence of reactantsa

2. Experimental section

Solution

2.1. Materials Potassium permanganate (Merck, India, 99%) and disodium salt of ethylenediaminetetraacetic acid (Merck, India, 99%) were used as received. Surfactants used were cetyltrimethylammonium bromide, sodiumdodecyl sulphate and Triton X-100, purchased from Merck (India). CO2 -free, deionized and distilled water was used for the preparation of stock solutions of all the reagents. Permanganate and EDTA solutions were stored in a dark glass and polythene bottles, respectively, because EDTA solution gradually leaches metal ions from glass containers, resulting in a change in the effective [EDTA]. An ELICO LI-120 digital pH meter fitted with a CH-41 combination electrode was used for pH measurements. 2.2. Rate measurements The reactions were started in glass-stoppered two-necked flask fitted with double walled condenser to check evaporation. A mixture containing required amount of permanganate, CTAB, water and other reagents (whenever necessary) was thermally equilibrated at desired temperature (25 ± 0.1 ◦ C) and to this was added a measured amount of EDTA solution, preequilibrated at the same temperature. The reaction volume was always 50 cm3 . Over the entire range of this study, reactions were carried out under pseudo-first-order conditions using an excess of [EDTA] over [MnO4 − ]. The rate of disappearance of permanganate ion was monitored at 525 nm using Spectronic21D Spectrophotometer with cell of path length 1 cm. A value of 2320 dm3 mol−1 cm−1 was calculated for the molar absorption coefficient of MnO4 − . The pseudo-first-order rate constants (kobs1 , s−1 ) were determined from the initial part of the plots of log (absorbance) versus time with a fixed-time method (vide infra). Other details of the kinetic measurements were the same as described elsewhere [19,21]. The pH of the reaction mixture was also measured at the end of each kinetic run and observed that pH drift during the reaction is very small (within 0.05 unit). 3. Critical micelle concentration (CMC) determination To determine the CMC, the conductivity measurements of the surfactants (CTAB and SDS) solutions were made with conductivity bridge (305, Systronic, India) using conductivity cell (CM 82T; cell constant = 1.02 cm−1 ), the cmc values of these surfactants were determined from plots of the specific conductivity versus [surfactant] in the absence and presence of MnO4 − and EDTA. The break point of nearly two striaght-line portions in the plot are taken as an indication of micelle formation and this corresponds to the CMC of surfactant [25]. The experiments were carried out at 25 ◦ C under varying conditions, that is, water + EDTA, water + MnO4 − and water + MnO− 4 + EDTA. The results are given in Table 1.

Water EDTA MnO4 −

104 cmc (mol dm−3 ) CTAB

SDS

10.1 (9.8) 8.8 8.7

80.0 (80.1) b 80.0 80.0

a [MnO − ] = 2.0 × 10−4 mol dm−3 , EDTA = 2.0 × 10−3 mol dm−3 , tempera4 ture = 25 ◦ C. b The literature values are quoted in parentheses at 25 ◦ C (Ref. [25]).

4. Results and discussion 4.1. Effects of [reactants] in presence of [surfactant] It is well established that surfactant can alter reaction mechanism, molecularities and orders by virtue of their medium effect; and that they can be utilized as mechanistic probes for reaction mechanism [26]. Preliminary observations showed that the solution of CTAB became turbid in presence of HClO4 . Therefore, H2 SO4 was used to maintain the acidic strength constant. The most interesting features of the present observations are the very fast decrease in the absorbance of the reaction mixture containing MnO4 − + EDTA + CTAB at 525 nm in the presence of H2 SO4 . Therefore, the choice of the best conditions for the kinetic experiments is a crucial problem that we address first. In order to examine the effect of variables, experiments were tried at [MnO4 − ] (=1.0 × 10−4 to 3.0 × 10−4 mol dm−3 ) and [EDTA] (=1.0 × 10−3 to 5.0 × 10−3 mol dm−3 ) in presence of [CTAB] (=0.5 × 10−4 to 8.0 × 10−4 mol dm−3 ), [TX100] (=8.0 × 10−4 mol dm−3 ) and [SDS] (=2.0 × 10−3 to 8.0 × 10−3 mol dm−3 ). It should be emphasized here that reactions were studied without adding HClO4 and H2 SO4 . Fig. 1 shows examples of some of the kinetic curves of the oxidation of EDTA by MnO4 − from which the values of rate constants were calculated. From the inspection of the plots of Fig. 1, it is clear that oxidation proceed in two stages. The time up to the linearity can be considered to be the non-catalytic path (first stage). After this point, the deviation in linear plot could be called the autocatalytic path (second stage). The time at which the deviation commenced was found to be decrease with the [EDTA]. These observations are in good agreement with our previous results [16,19]. The pseudo-first-order rate constants (kobs1 ) were obtained from the measurements of slopes of the initial tangents to the plots of log(absorbance) versus time. It was observed that the rate constants decreased as the initial [MnO4 − ] increased at fixed [EDTA] (=2.0 × 10−3 mol dm−3 ), [CTAB] (=2.0 × 10−4 mol dm−3 ) and temperature (=25 ◦ C) (Table 2). The basic trend in chemical kinetics is that the pseudo-first-order rate constants are independent of the initial concentration of the reactant in defect. The same type of defect has been obtained in many MnO4 − reactions and especially in those having an autocatalytic character [14,15]. At constant [MnO4 − ] (=2.0 × 10−4 mol dm−3 ), [CTAB] (=2.0 × 10−4 mol dm−3 ) and temperature (=25 ◦ C), the reac-

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Fig. 1. Plots of log(absorbance) versus time at 525 nm for the oxidation of EDTA by MnO4 − in presence of CTAB. Reaction conditions: [MnO4 − ] (=2.0 × 10−4 mol dm−3 ); [EDTA] (=2.0 × 10−3 mol dm−3 ); temperature = 25 ◦ C; [CTAB] = (A) 1.0, (B) 2.0, (C) 3.0, (D) 4.0 and (E) 5.0 × 10−4 mol dm−3 .

tion rate increased with an increase in [EDTA] (=1.0 × 10−3 to 5.0 × 10−3 mol dm−3 ). The plot of log kobs1 versus log [EDTA] is linear with slope ca. 0.58, indicating fractional-order dependence of kobs1 on [EDTA]. On the other hand, the plot of 1/kobs1 versus 1/[EDTA] is linear with a positive intercept and positive slope (Fig. 2). Such plot is indicative by Michacelis–Menten behavior (kinetic proof for complex formation between MnO4 − and EDTA). The same experiments were also performed in presence of [TX-100] (=8.0 × 10−4 mol dm−3 ) and [SDS] (=8.0 × 10−3 mol dm−3 ). No appreciable change was observed in rate constants in presence of these surfactants. In order to calculate the activation parameters, a series of kinetic runs performed at different temperatures at constant [MnO4 − ] (=2.0 × 10−4 mol dm−3 ) and [EDTA] (=2.0 × 10−3 mol dm−3 ) in absence and presence of [CTAB] (=2.0 × 10−4 mol dm−3 ). The activation energy (Ea ) and other parameters H# and S# ) of this system were evaluated from Arrhenius (log kobs versus 1/T; Fig. 3) and Eyring (log kobs /T versus 1/T) plots. The results are summarized in Table 3. A comparison between the Ea values in aqueous and micellar media

Fig. 2. Plot of 1/kobs1 versus 1/[EDTA]. Reaction conditions: [MnO4 − ] (=2.0 × 10−4 mol dm−3 ); [CTAB] (=2.0 × 10−4 mol dm−3 ); temperature (=25 ◦ C).

Fig. 3. Arrhenius plots for the oxidation of EDTA by MnO− 4 in presence (A) and absence (B) of CTAB. Reaction conditions: [MnO4 − ] (=2.0 × 10−4 mol dm−3 ); [EDTA] (=2.0 × 10−3 mol dm−3 ); [CTAB] (=2.0 × 10−4 mol dm−3 ).

Table 2 Dependence of pseudo-first-order rate constants on [MnO4 − ] and [EDTA] for the reduction of MnO4 − by EDTA at 25 ◦ C in presence of CTAB (=2.0 × 10−4 mol dm−3 ) 104 [MnO4 − ] (mol dm−3 ) 0.6 1.0 1.4 2.0 2.4 2.0

a

103 [EDTA] (mol dm−3 )

pH

2.0

5.26

1.0 2.0 3.0 4.0 5.0

5.50 5.26 5.17 5.10 5.07

The kobs1 values in the absence of CTAB are given in parentheses.

104 kobs1 (s−1 )

104 kcal1 (s−1 )

(kobs1 − kcal1 )/kobs1

1.4 2.2 2.7 3.0 3.4

0.00 0.00 −0.08 −0.07 0.00

(1.7)a

3.5 2.9 (1.6) 2.6 (1.4) 2.2 (1.2) 1.3 (0.9) 1.4 2.2 2.5 2.8 3.4

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Table 3 Values of pseudo-first-order rate constants (kobs1 ) and activation parameters for the reduction of MnO4 − by EDTA in presence of CTAB (=2.0 × 10−4 mol dm−3 ) Temperature (◦ C)

103 [EDTA] (mol dm−3 )

104 [MnO− 4 ] (mol dm−3 )

104 kobs1 (s−1 )

25 35 45

2.0

2.0

2.2 (1.8)a 4.5 (3.8) 9.3 (8.3)

Activation parameters (kJ mol−1 )

Ea H# (kJ mol−1 ) S# (J K−1 mol−1 ) a

Aqueous

Micellar

60 57 −517

56 54 −503

The kobs1 values in the absence of CTAB are given in parentheses.

indicates that the CTAB surfactant act as catalyst and provide a new reaction path with a lower value of Ea . To study the effect of pH, a series of kinetic runs were also performed in presence of H2 SO4 (range: 1.0 × 10−4 to 5.0 × 10−4 mol dm−3 ) at constant concentrations of other reagents (EDTA, MnO4 − and CTAB). The rate constants, obtained as a function [H2 SO4 ] was found to increase with increasing amounts of [H2 SO4 ] (kobs1 × 104 = 1.5, 1.8, 2.1, 7.7 and 9.2 s−1 at [H2 SO4 ] = 1.0, 2.0, 3.0, 4.0 and 5.0 × 10−4 mol dm−3 , respectively). We may thus safely conclude that the oxidation of EDTA depends upon the acidity of the medium. 4.2. Rate constants—[surfactants] profile for oxidation of EDTA by MnO4 − In order to see the role of surfactants on the oxidation of EDTA by MnO4 − , the effects of [CTAB], [SDS] and [TX-100] were studied at 2.0 × 10−3 mol dm−3 EDTA and 2.0 × 10−4 mol dm−3 MnO4 − . Preliminary observations showed that the oxidation of EDTA by permanganate is very fast in presence of CTAB (>6.0 × 10−4 mol dm−3 ). Therefore, the kinetic studies were limited in the [CTAB] range of 0.5 × 10−4 to 5.0 × 10−4 mol dm−3 .The plot of kobs1 against [CTAB] shows gradual increase of rates of nearly 10-fold with the increase in [CTAB] (Fig. 4) which clearly demonstrate the CTAB catalytic effect not only above but even below CMC, i.e., micellar as well as submicellar catalysis are observed. These observations are most interesting instead of kobs1 , attained constant values (for unimolecular reactions) or pass through a maximum (for bimolecular reactions) with [CTAB], [7,27]. The observed catalytic effect may, therefore, be due to (i) presence of premicelles and/or (ii) preponement of micellization by reactants [28] (as is also confirmed by CMC decreases at reaction conditions, Table 1). The anionic SDS surfactant neither catalysed nor inhibited the oxidation reaction (Fig. 4). It is not surprising because there is an electrostatic repulsion between the reactants (MnO4 − and EDTA) (vide infra) and the negative head groups of SDS micelles. As a result both the reactants are located in the bulk aqueous medium and the rate remains unaffected. On the other hand, reaction mixture containing MnO4 − , EDTA and non-ionic TX-100 surfactant became turbid after some time due to the instability of polyoxyethylene chain of TX-100.

4.3. Analysis of kobs1 —[CTAB] data It is well established that an aqueous surfactant solution has three components: surfactant monomers in the aqueous solution, micellar aggregates, and monomers absorbed as a film at the interface. The surfactant is in dynamic equilibrium among all these components [29]. Surfactant monomers rapidly join and leave micelles (as micelles have a transient character), and the aggregation number represents only an average over time. The premicellar catalytic effect can be brought in the fact that small aggregates of CTAB (dimers, trimers, tetramers, etc.) exit below the CMC, these small submicellar aggregates can interact physically with the reactants forming active entities. On the basis of the above description, the possible cause of rate enhancement may be discussed. The presence of negative charge on the EDTA and MnO4 − must be considered. It is certainly possible that the negative charge on MnO4 − forms an ion pair (Q+ MnO4 − ) with the positive quaternary ammonium (Q+ ; –N+ (CH3 )3 ) head group of CTAB molecules which

Fig. 4. Effect of [surfactants] (CTAB () and SDS (䊉)) on the rate constants of the oxidation of EDTA by MnO4 − . Reaction conditions: [MnO4 − ] (=2.0 × 10−4 mol dm−3 ); [EDTA] (=2.0 × 10−3 mol dm−3 ); temperature (=25 ◦ C).

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Scheme 1. Proposed mechanism for the oxidative degradation of EDTA by MnO4 − .

brings the reactants together through electrostatic interactions. Coordination of a Q+ cation by the permanganate anion would decrease electron density of the MnO4 − which, in turn, increases the oxidizing power of the permanganate. This is a reasonable explanation for the submicellar catalysis in the present work. Table 1 suggesting that the MnO4 − and EDTA are interacts with the CTAB surfactant, and submicellar aggregates are formed [30,31]. Both the reactants will be preferentially located at the positively charged CTAB aggregated molecules and, therefore, the kinetic CMC of CTAB is lower than in water. 5. Mechanism In aqueous solution, various EDTA species like H6 Y2+ , H5 Y+ , H4 Y, H3 Y− , H2 Y2− , HY3− and Y4− (Y = EDTA) exist in equilibrium and nature of these species depends upon the pH of the solution [32]. The first two species (H6 Y2+ and H5 Y+ ) are relatively strong acids and normally are not of importance in evaluation of dissociation constants. Therefore, EDTA has only four values of dissociation constants, i.e., Ka3 = 1.02 × 10−2 ; Ka4 = 2.14 × 10−3 ; Ka5 = 6.92 × 10−7 ;

Ka6 = 5.5 × 10−11 . Under the experimental conditions used in this work ([HClO4 ] = 0.0 mol dm−3 ), H4 Y species exists in significant concentration and this species is reactive towards complexation with permanganate. The most satisfactory mechanism to fit the experimental data is represented by Scheme 1. In Scheme 1, the reactive species of EDTA (H4 Y) and permanganate (MnO4 − ) readily form complex (complex 1; Eq. (2)). By analogy with previous results [12,13,15,17] we assume that complex1 decomposes in a rate-determining one-step oneelectron oxidation–reduction mechanism to give free radical and Mn(VI) (Eq. (3)). After the slow step, the radical reacts with a molecule of MnO4 − to yield the imine cation and Mn(VI) (Eq. (5)). In the next step, imine cation gives HCHO and ethylenediaminetriacetic acid (oxidation products of EDTA) (Eq. (6)) after hydrolysis. It is well known that in the MnO− 4 reduction, various species of manganese (Mn(VI), Mn(V), Mn(IV), and Mn(III)) are formed as an intermediate(s). The presence of Mn(VI) and Mn(V) is ruled out by the fact that they are highly unstable in an aqueous acidic neutral media [12]. The intermediate(s), Mn(VI) and Mn(V) may decompose to Mn(IV) species. This oxidation state is commonly involved in the MnO4 − oxidation of organic compounds [12,13]. A rate law consistent with Scheme 1 may

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be expressed as Eq. (7). kobs1 = 1 kobs1

=

k1 Kes1 [EDTA] 1 + Kes1 [EDTA]

(7)

1 1 + k1 Kes1 [EDTA] k1

(8)

Thus, a plot of 1/kobs1 versus 1/[EDTA] should give a straight line (Fig. 2). The values of k1 and Kes1 were calculated from the intercept and slope of the Fig. 2, and are 5.0 × 10−4 s−1 and 403 mol−1 dm3 , respectively. Therefore, it was conformed that the redox reaction of MnO4 − and EDTA occurs in two kinetically distinguishable steps. The first step is a fast formation of complex between MnO4 − and EDTA (Eq. (2)). The second step is a slower electron transfer from EDTA to the manganese(VII) with in the complex (Eq. (3)). Using values of k1 , [EDTA] and Kes1 , the kcal1 , can be generated for various kinetic runs (Table 2). The close agreement between kobs1 and kcal1 , provides the supportive evidence for the proposed mechanism (Scheme 1) and conforms the validity of the ratelaw (Eq. (7)). The proposed mechanism is further supported by analysis of the products. Formaldehyde has been detected as 2,4-dinitrophenylhydrozone derivative. Ethylenediaminetriacetic acid was detected by reported methods [33,34]. Carbon dioxide was also identified by Ba(OH)2 [35]. Formation of radicals during the redox process was confirmed by the addition of saturated solution of HgCl2 in the reaction mixtures (white precipitate was observed). It has been established that Mn(IV) species is responsible for the auto catalysis observed in many permagnate reactions [10–12,36]. Therefore, in order to confirm the formation of Mn(IV) as an intermediate, some kinetic runs were also performed at 420 nm, where the contribution from MnO4 − is negligible. All attempts to observe Mn(IV) at 420 nm were unsuccessful. Therefore, EDTA did not form Mn(IV) complex stable enough to be detected under conditions used in this present kinetic measurements. In the light of above observations and discussion, the proposed mechanism is given in Scheme 2 for the autoacceleration pathway.

Scheme 2. Mechanism for the autoacceleration path of EDTA oxidation.

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In presence of large amount of EDTA, the Mn(IV) immediately gets converted into stable product (Mn(II); Eqs. (9) and (10)). The reduction of Mn(IV) to Mn(III) by Mn(II) has also been observed [21] (Eq. (11)). Scheme 2 clearly indicates that the autocatalytic pathway is not a true path of MnO4 − –EDTA reaction. It may be a mixture of a series of reactions (Eqs. (9) to (14)). Therefore, the exact dependence of kobs2 on [EDTA] cannot be estimated. On the other hand, permanganate also oxidizes –CHO and –NH2 to –COOH and –NO2 , respectively. Hence, the final products (formaldehyde and ethylenediaminetriacetic acid, which are produced in reaction (6), could, in principle, react with MnO4 − . Finally, we can state that when an excess of permanganate over EDTA is used, further oxidation of the intermediate(s) occurs to yield CO2 and nitro compounds as the final products. 6. Conclusion It is difficult to ascertain the exact reaction site of a permicellar and micellar mediated reaction [37,8]. From the present data, one can obtain the preliminary picture of the reaction sites. The key fact are: (1) the reaction proceeds more fast in presence of cationic, (CTAB) surfactant than in aqueous phase; (2) the anionic (SDS) and non-ionic (TX-100) surfactants has no effect on the reaction rate; (3) both the reactants (MnO4 − and EDTA) proceeds towards the cationic head group of CTAB dimer, tetramer, etc.); (4) the presence [H+ ] is not essential for the oxidation of EDTA in presence of CTAB; and (5) the oxidative degradation of EDTA by MnO4 − is first order in [EDTA], but in the presence of CTAB the reaction is fractional-order in [EDTA]. This gives an indication to the relationship between the charge effect of surfactants and the molecularities and or order of the reaction [38]. References [1] D. Attwood, A.T. Florence, Surfactant Systems their Chemistry Pharmacy and Biological Properties, Chapman and Hall, NY, 1983. [2] D. Piszkiewicz, J. Am. Chem. Soc. 99 (1977) 7695. [3] W. Richmond, C. Tondre, E. Krzyzanowska, J. Szymanowski, J. Chem. Soc. Faraday Trans. 91 (1995) 657. [4] I.A. Vinokurov, J. Kankare, Langmuir 18 (2002) 6789. [5] J. Herszage, M.D. Afonso, Langmuir 19 (2003) 9684. [6] F.M. Menger, C.E. Portony, J. Am. Chem. Soc. 89 (1967) 4698. [7] C.A. Bunton, Catal. Rev. Sci. Eng. 20 (1979) 1. [8] M.N. Khan, J. Colloid Interface Sci. 170 (1995) 598. [9] J.H. Fendler, Membrane Mimetic Chemistry, Characterization and Applications of Micelles, Microemulsions, Monolayers, Bilayers, Vescicles Hostguest Systems and Polyions, Wiley, NY, 1982. [10] C.A. Bunton, J. Mol. Liq. 72 (1997) 231. [11] Kabir-ud-Din, K. Hartani, Z. Khan, Colloid Surf. A: Physicochem. Eng. Asp. 193 (2001) 1. [12] L.I. Simandi, M. Jaky, J. Am. Chem. Soc. 98 (1976) 1995. [13] F. Freeman, J.C. Kappos, J. Am. Chem. Soc. 107 (1985) 6628. [14] J.F. Perez- Benito, C. Arias, J. Colloid Interface Sci. 149 (1992) 92. [15] J.F. Perez- Benito, C. Arias, E. Amat, J. Colloid Interface Sci. 177 (1996) 288. [16] Z. Khan, Raju, M. Akram, Kabir-ud-Did, Int. J. Chem. Kinet. 36 (2004) 359. [17] Kabir-ud-Din, W. Fatma, Z. Khan, Colloid Surf. A: Physicochem. Eng. Asp. 234 (2004) 159.

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