Effect of surfactants on the oxidation of oxalic acid by soluble colloidal MnO2

Colloids and Surfaces A: Physicochem. Eng. Aspects 234 (2004) 159–164

Effect of surfactants on the oxidation of oxalic acid by soluble colloidal MnO2 Kabir-ud-Din a,∗ , Waseefa Fatma a , Zaheer Khan b b

a Department of Chemistry, Aligarh Muslim University, Aligarh, 202002, U.P., India Department of Chemistry, Jamia Millia Islamia, Jamia Nagar, New Delhi 110025, India

Received 28 May 2003; accepted 12 December 2003

Abstract Kinetics of the oxidation of oxalic acid by water-soluble colloidal MnO2 has been studied spectrophotometrically in the absence and presence of non-ionic octylphenoxypolyethoxyethanol (Triton X-100, TX-100) and anionic sodium dodecyl sulfate (SDS) surfactant micelles. Anionic SDS micelles were found ineffective whereas those of TX-100 increased the rate that reached a plateau at higher [TX-100]. The reaction followed first- and fractional-order kinetics with respect to [MnO2 ] and [H+ ], respectively, both in the absence and presence of TX-100. The activation parameters (Ea , H# , and S# ) were evaluated and the effect of added manganese(II) sulphate is also reported in both aqueous and aqueous–micellar media. On the basis of experimental findings, a probable mechanism is proposed. The reaction proceeds through the adsorption of oxalic acid and hydrogen ions on the surface of the colloidal MnO2 . Freundlich isotherm is used to explain the adsorption of oxalic acid on colloidal MnO2 . © 2004 Elsevier B.V. All rights reserved. Keywords: Surfactants; Manganese dioxide; Oxalic acid; Kinetics; Sodium dodecyl sulfate; Triton X-100

1. Introduction Electron-transfer processes in micellar systems can be considered as models to get insight into electron transport occurring in biological phenomena [1]. A number of interfacial electron-transfer processes have been investigated in vesicles, poly-electrolytes, and micellar surfaces including photoredox reactions [2]. While kinetic studies in micellar media of both inorganic [3–7] and organic reactions [1,2,8–15] have been in the scope of interest of many researchers for a long time, works on colloidal aggregates as a reactant are scarce [16–18] despite of sufficient literature being available on adsorption of surfactants on sols [19–21]. In a number of reports on the oxidation of organic compounds by potassium permanganate, particular focus has been made on the involvement of manganese dioxide as an active autocatalyst [22–26] and on its role in oscillating reactions [27–30]. It has been used (in the form of aqueous suspensions) as catalyst or oxidizing agent in both

∗

Corresponding author. Tel.: +91-571-2703515. E-mail address: [email protected] (Kabir-ud-Din).

0927-7757/$ – see front matter © 2004 Elsevier B.V. All rights reserved. doi:10.1016/j.colsurfa.2003.12.015

organic [31] and inorganic [32] reactions. Perez-Benito et al. [16,17,33] have recently established the stability of water-soluble form of colloidal MnO2 prepared from the reduction of aqueous permanganate by sodium thiosulphate under neutral conditions. However, micellar effects on the redox reactions of water-soluble colloidal MnO2 has not been studied so far except the report of Tuncay et al. [18]. Therefore, the kinetics of the redox reaction between colloidal MnO2 and oxalic acid was studied to gain further insight on the role of both electrostatic and hydrophobic interactions on the metal ion oxidation of organic reductants.

2. Experimental 2.1. Materials Aqueous solutions of potassium permanganate (E. Merck, 98.5%), sodium thiosulfate (E. Merck, 99%), and manganese(II) sulfate-1-hydrate (E. Merck, 99%) were prepared by dissolving the requisite amounts in doubly distilled and deionized water. The surfactants used were sodium dodecyl sulfate (SDS, Sigma, USA, 99%) and Triton X-100 (Fluka,

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Switzerland, 99%). All the other reagents (oxalic acid, perchloric acid, and gum arabic) were commercially available analytical reagents and used as received. 2.2. Preparation of the colloidal MnO2 sol Water-soluble colloidal MnO2 was prepared by the method of Perez-Benito et al. [16,17,33]. The respective concentrations of potassium permanganate and sodium thiosulphate were 5.0 × 10−4 and 1.875 × 10−4 mol dm−3 [34]. 2.3. Kinetic spectrophotometric measurements Kinetic experiments were performed at 30±0.1 ◦ C (except while studying the effect of temperature variation) in a paraffin oil-bath. The reactions were carried out in a three-necked reaction vessel blackened from outside to check the effect of light on the rates. The reaction was initiated by adding temperature pre-equilibrated solution of oxalic acid (2.8 × 10−4 mol dm−3 ) to thermally equilibrated mixture of other reagents including colloidal MnO2 . The kinetics was monitored spectrophotometrically by measuring the absorbance of remaining colloidal MnO2 at known time intervals at 425 nm (λmax of colloidal MnO2 [34]). The reactions were usually followed up to not less then 80% completion. The concentration of oxalic acid was always in excess over MnO2 and pseudo-first-order rate constants (kobs or kψ in s−1 ) were evaluated by plotting log(A425 ) versus time.

Fig. 1. Absorption spectra of mixtures containing fixed amount of KMnO4 (=4.0 × 10−4 mol dm−3 ) and varying amount of Na2 S2 O3 (0.0 mol dm−3 , (䊏), 2.0 × 10−4 mol dm−3 , (䉲)) at 30 ◦ C.

2.0 × 10−5 mol dm−3 , the kobs remained constant with increasing [MnO2 ] (Table 1). According to the basic tenets of chemical kinetics the pseudo-first-order rate constants should be independent of the initial concentration of MnO2 , the initial decrease of kobs may then be due to possible flocculation of the colloidal MnO2 particles. In order to obtain constant values of rate constants, all the experiments were repeated in presence of gum arabic (a well known stabilizer of

3. Results and discussion 3.1. Characterization of colloidal MnO2 The UV-Vis spectrum of aqueous KMnO4 solution possesses absorption band located at λmax = 530 nm (Fig. 1 shown by the symbol 䊏). With the addition of Na2 S2 O3 solution, the band gradually disappeared with the appearance of a new single broad band of high intensity at 420 nm. The concentration of Na2 S2 O3 needed to stabilize the new spectrum (Fig. 1) is 2×10−4 mol dm−3 . As the new stabilized spectrum is that of colloidal MnO2 [34], the wavelength of 425 nm was chosen to follow the kinetic runs. 3.2. Reaction in the absence of surfactants The kinetic results of the reduction of colloidal MnO2 by oxalic acid obtained in aqueous medium under different experimental conditions can be summarized as follows: (i) The order with respect to colloidal MnO2 was determined by finding kobs at different initial concentrations of MnO2 (1.6 × 10−5 to 2.8 × 10−5 mol dm−3 ) with other parameters remaining fixed ([oxalic acid] = 2.8× 10−4 mol dm−3 , [HClO4 ] = 0.4 × 10−4 mol dm−3 , 30 ◦ C). With an initial decrease up to [MnO2 ] ≤

Table 1 Effect of varying [colloidal MnO2 ] and [MnSO4 ] on the pseudo-first-order rate constants (kobs or kψ ) for the oxidation of oxalic acid (=2.8 × 10−4 mol dm−3 ) by colloidal MnO2 in the absence and presence of TX-100 at 30 ◦ C and [HClO4 ] = 0.4 × 10−4 mol dm−3 105 [colliodal MnO2 ] (mol dm−3 )

105 [MnSO4 ] (mol dm−3 )

103 kobs or 103 kψ (s−1 ) Aqueousa

TX-100b

1.6 1.8 2.0 2.2 2.4 2.6 2.8 2.8

0.0

4.6 3.8 2.7 2.7 2.7 2.7 2.7 2.7 3.1 3.8 4.2 –

6.5 5.0 3.4 3.5 3.4 3.4 3.4 3.4 3.8 4.6 5.4 5.7

0.0 0.2 0.4 0.6 0.8

(1.7) (1.6) (1.6) (1.6) (1.6) (1.7) (1.6)

a Values obtained in presence of gum arabic (=0.0511%) are given in parentheses. b [TX-100] = 0.8 × 10−3 mol dm−3 .

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Table 2 Effect of varying the [oxalic acid], [HClO4 ], and temperature on the pseudo-first-order rate constants (kobs or kψ ) for the oxidation of oxalic acid by colloidal MnO2 (=2.8 × 10−5 mol dm−3 ) Temperature (◦ C)

104 [oxalic acid] (mol dm−3 )

104 [HClO4 ] (mol dm−3 )

2.8 3.2 3.6 4.0 4.4 4.8 5.2 2.8

0.4

30

0.0 0.2 0.4 1.0 1.1 1.2 1.6 2.0 2.4 2.8 0.4

30

2.8

Ea (kJ mol−1 ) H# (kJ mol−1 ) −S# (J K−1 mol−1 ) a

103 kobs or 103 kψ (s−1 ) Aqueous

20 25 30 35 31 33 184

2.7 3.4 4.0 4.8 5.8 5.8 6.1 1.6 2.3 2.7 3.0 3.1 3.3 3.8 4.4 5.0 5.4 1.9 2.5 2.7 3.8

TX-100a

2.6 3.1 3.4 4.6 28 30 193

[TX-100] = 0.8 × 10−3 mol dm−3 .

colloidal MnO2 [35]). This time the kobs values were found to be independent of [MnO2 ] (Table 1, shown in parentheses), which indicate that the reaction follows the first-order kinetics with respect to [MnO2 ]. (ii) The order with respect to oxalic acid was deduced from the values of kobs obtained at several [ox-

alic acid] (2.8 × 10−4 to 5.2 × 10−4 mol dm−3 ) with fixed [MnO2 ] (=2.8 × 10−5 mol dm−3 ), [HClO4 ] (=0.4 × 10−4 mol dm−3 ), and temperature (30 ◦ C). The results, in Table 2 and Fig. 2, indicate a sort of saturation approach with increasing [oxalic acid]. A double-logarithmic fit yielded a slope of 1.2 with

Fig. 2. Plot showing dependence of the pseudo-first-order rate constants on the oxalic acid concentration for its oxidation by colloidal MnO2 (=2.8 × 10−5 mol dm−3 ) in the presence of HClO4 (=0.4 × 10−4 mol dm−3 ) at 30 ◦ C.

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r = 0.9848. Thus, the order with respect to [oxalic acid] lies between 1 and 2. (iii) The reaction rate was found to increase with increasing [HClO4 ]. These results are summarized in Table 2. The plot of log kobs versus log[H+ ] is linear with slope 0.31, indicating the order to be fractional with respect to [H+ ]. Both pH-dependent and -independent pathways are involved. In view of the above points, the empirical rate law is written as −

d[MnO2 ] dt = {k + k

[H+ ]0.31 }[MnO2 ][oxalic acid]1.2

(1)

(iv) Effect of temperature was also seen within the range 20–35 ◦ C at constant [MnO2 ] (=2.8×10−5 mol dm−3 ), [HClO4 ] (=0.4 × 10−4 mol dm−3 ), and [oxalic acid] (=2.8 × 10−4 mol dm−3 ). The kobs values are given in Table 2. Activation parameters (Table 2) were calculated using Arrhenius and Eyring equations by linear least-squares method. (v) It is well known that Mn(II) ion acts as an active autocatalyst and, if formed as a reaction product, contributes to accelerate the destruction of colloidal MnO2 [17,34]. To substantiate the role of Mn(II) during the oxidation of oxalic acid by colloidal MnO2 , varying amounts of Mn(II) (as MnSO4 ) were added initially to the reaction mixture. It was found that the oxidation rate of oxalic acid was increased by adding Mn(II) ions (Table 1). The observation suggests that Mn(II) has an autocatalytic role. Taking cognizance of the [H+ ] range (0.2 ×10−4 to 2.8 × 10−4 mol dm−3 ) used in the kinetic experiments (Table 2) and that of the dissociation constants of oxalic acid [36], the following mechanism is proposed for the noncatalytic reaction part: (a) pH-independent pathway (2)

C1

C1 − → (MnO2 )x−1 + Mn(II) + 4CO2 + 2H2 O

(3)

(b) pH-dependent pathway KII

(MnO2 )x − (H2 C2 O4 )2 + H+
(MnO2 )x C2

−(H2 C2 O4 )2 − (H+ )

(4)

kII ˙ → (MnO2 )x−1 + Mn(II) + 3CO2 + H2 O + COOH C2 −

(5) ˙ (MnO2 )x + COOH + 3H+ fast

υ = {kI + kII

[(H+ )s ]}[MnO2 ][(H2 C2 O4 )s ]2

(7)

where kI = kI KI , kII

= kII KI KII (kI and kII

are the rate constants corresponding to pH-independent and -dependent pathways, whereas (H+ )s and (H2 C2 O4 )s refer to adsorbed species on the colloid surface), and for the first-order rate constants: υ (8) kobs = = {kI + kII

[(H+ )s ]}[(H2 C2 O4 )s ]2 [MnO2 ] Adsorption of organic acids (oxalic, formic, and lactic) on the surface of colloidal particles in quasi-equilibrium reactions prior to the redox rate-determining steps is widely accepted. Therefore, adsorption isotherm can be used to explain the observed result. Writing Freundlich adsorption isotherm for the species present in the bulk solution and adsorbed on colloidal MnO2 , we get, according to equilibrium steps (2) and (4)

[(H+ )s ] = y [H+ ]n

(9)

and



[(H2 C2 O4 )s ] = y

[H2 C2 O4 ]n

(10)

where y and n are the adsorption parameters for the respective species. Eq. (7) can now be written as





υ = {kI (y

)2 + kII

y (y

)2 [H+ ]n }[MnO2 ][H2 C2 O4 ]2n

(11) Eq. (11) agrees very well with the experimental results (see Eq. (1)) with k = kI (y

)2 and k

= kII

y (y

)2 . Also, n = 0.31 and n

= 1.2/2 = 0.6. It is well known that for applicability of the Freundlich isotherm, the exponent values must lie within 0–1 (involvement of two oxalic acid molecules, Eq. (2), was considered because of this reason). 3.3. Reaction in the presence of surfactants

KI

(MnO2 )x + 2H2 C2 O4
(MnO2 )x − (H2 C2 O4 )2 kI

The total reaction rate, derived by the sum of the contributions of the pH-independent and -dependent pathways is:

− → (MnO2 )x−1 + Mn(II) + CO2 + 2H2 O

(6)

The effect of varying anionic [SDS] in the range 4 × 10−3 to 28 × 10−3 mol dm−3 was studied at constant [oxalic acid] (=2.8×10−4 mol dm−3 ), [MnO2 ] (=2.8×10−5 mol dm−3 ), [HClO4 ] (=0.4 ×10−4 mol dm−3 ), and temperature (30 ◦ C). It was observed that SDS has no affect on the value of kψ (Table 3). The result suggests that there is repulsion between anionic micellar aggregates of SDS (the micelles possess net negative charge due to –OSO3 − ) and the negatively charged colloidal MnO2 [33,34]. In order to see the rate of non-ionic TX-100 surfactant, the pseudo-first-order rate constants (kψ ) for the oxidation of oxalic acid by colloidal MnO2 were determined in presence of TX-100 (Tables 1–3). The plots between log(A425 ) versus time were linear in presence of various [TX-100], confirming that the reaction is first-order in the presence of TX-100 also. The order in [MnO2 ] remains the same as that observed in

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Table 3 Effect of varying the [surfactant] on the kψ for the oxidation of oxalic acid (=2.8×10−4 mol dm−3 ) by colloidal MnO2 (=2.8×10−5 mol dm−3 ) at 30 ◦ C and [HClO4 ] = 0.4 × 10−4 mol dm−3 103 [SDS] (mol dm−3 )

103 kψ (s−1 )

0.0 4.0 6.0 8.0 10.0 12.0 14.0 20.0 24.0

2.7 2.6 2.7 2.7 2.6 2.7 2.7 2.6 2.7

104 [TX-100] (mol dm−3 ) 0.0 1.0 2.0 3.0 4.0 5.0 7.0 8.0 −

103 kψ (s−1 ) 2.7 2.9 3.1 3.3 3.4 3.3 3.4 3.4 −

the absence of micelles. The activation parameters (Table 2) are also same as that observed in the absence of TX-100. These observations establish that the mechanism operative in the aqueous medium (see Section 3.2) is being followed in the TX-100 micellar medium too. The kψ −[TX-100] profile shows positive catalysis (Fig. 3), which may be explained in terms of adsorption, increasing stabilization, and association of the colloidal MnO2 and oxalic acid with increase in [TX-100] that reaches a limiting value. The reaction may then be discussed in terms of mathematical model proposed by Tuncay et al. [18] as log kψ = 0.0760 log[TX-100] − 2.210

(12)

According to Eq. (12), a plot of log kψ versus log[TX-100] should be linear with a slope and intercept on the y-axis. An alternative data treatment was also carried out by using Eq. (13) which predicts that the plot of 1/(kψ −kobs ) against 1/[TX-100] should be linear. b 1 =a+ kψ − kobs [TX-100]

(13)

Fig. 3. Plots showing dependence of the pseudo-first-order rate constants on the TX-100 (A) and SDS (B) concentrations for the oxidation of oxalic acid (=2.8×10−4 mol dm−3 ) by colloidal MnO2 (=2.8×10−5 mol dm−3 ) in the presence of HClO4 (=0.4 × 10−4 mol dm−3 ) at 30 ◦ C.

Fig. 4. Plot of log kψ vs. log[TX-100] for the oxidation of oxalic acid (=2.8 × 10−4 mol dm−3 ) by colloidal MnO2 (=2.8 × 10−5 mol dm−3 ) in the presence of HClO4 (=0.4 × 10−4 mol dm−3 ) at 30 ◦ C. The line is drawn according to Eq. (12).

That indeed is the behavior of the present system according to Eqs. (12) and (13) and can be seen from the plots of Figs. 4 and 5. Values of the parameters a and b were found to be 6.4 × 102 s and 0.4169 mol dm−3 s. 3.4. Probable role of TX-100 The effect of the ionic and non-ionic micelles on the reaction rates of bimolecular reactions is due to the association/incorporation through electrostatic/hydrophobic and hydrogen bonding interactions between the reactants within a small volume of the self-assemblies [37,38]. The micelle is a porous cluster with rough surfaces and deep water filled cavities. Therefore, it is difficult to ascertain the exact reaction site of a micellar-mediated reaction. It is well known that hydroxy groups of organic molecules play an important role in the adsorption of reductants, through hydrogen bonding, on the surface of colloidal particles [16–18]. As

Fig. 5. Plot of 1/(kψ −kobs ) against 1/[TX-100] for the oxidation of oxalic acid (=2.8×10−4 mol dm−3 ) by colloidal MnO2 (=2.8×10−5 mol dm−3 ) in the presence of HClO4 (=0.4 × 10−4 mol dm−3 ) at 30 ◦ C.

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the number of –OH groups increases, the probability of adsorption increases. In presence of TX-100 (I), there will be a competition between the oxalic acid and TX-100 to be adsorbed on the sites of the colloidal MnO2 . As the oxalic acid is hydrophilic, hydrophobic, and electrostatic interactions can be largely

neglected with TX-100 micelles but, due to containing two –COOH groups, hydrogen bonding may occur between the –COOH groups of oxalic acid and the ether−oxygen of the polyoxyethylene chains of TX-100. As a number of donor groups are present in one molecule of TX-100, more H2 C2 O4 molecules may bind to form multiple hydrogen bonds. The hydrogen bonding between the polar polyethoxyethylene chains of TX-100 and the reactants may, therefore, be responsible for the catalytic role of non-ionic TX-100. The surfactant thus helps in bringing the reactants closer, which may orient in a manner suitable for the redox reaction followed by rearrangement of TX-100 molecules. Although exact locations of both the reactants on the surface of the TX-100 micelle cannot be ascertained, mere presence of TX-100 affecting the rate of redox reaction reveals that this reaction occurs in the close proximity of the micellar surface.

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