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Colloids and Surfaces A: Physicochem. Eng. Aspects 302 (2007) 102–106

A kinetic study of water-soluble colloidal MnO2 formed bythe reduction of permanganate by thiourea

Salman Ahmad Khan, Parveen Kumar, Kishwar Saleem, Zaheer Khan ∗Department of Chemistry, Jamia Millia Islamia, Jamia Nagar, New Delhi 110025, India

Received 28 April 2006; received in revised form 20 January 2007; accepted 1 February 2007Available online 6 February 2007

bstract

A conventional spectrophotometric technique was used to study the oxidation of thiourea by permanganate in a dilute perchloric acid solution.reliminary observations indicate that the reduction of permanganate by thiourea proceeds through the fast formation of colloidal MnO2 as an

ntermediate. Various experiments were performed to confirm the nature of intermediate(s) formed during the reduction of permanganate by

hiourea. The kinetic and spectroscopic data are consistent with the formation of water-soluble colloidal MnO2. The stoichiometry was found toe 1:2 (MnO2:thiourea). The stability of water-soluble colloidal MnO2 depends upon the [H+]. The effect of total [MnO4

−], [thiourea] and [H+] onhe reaction rate was determined. On the basis of various observations, a mechanism is proposed for the oxidation of thiourea by permanganate. 2007 Elsevier B.V. All rights reserved.

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eywords: Thiourea; Permanganate; Oxidation; Kinetics; Colloidal MnO2

. Introduction

Oxidation by permanganate ion is applied extensively inrganic synthesis [1]. During the oxidation, it is evident thatermanganate is reduced to various oxidation states in acidic,lkaline and neutral media. Soluble forms of colloidal man-anese dioxide are formed as products in many permanganateeactions carried out in near-neutrality aqueous solutions, anduite often they play an important role in autocatalytic [2] andscillating [3] reactions. They are also involved as detectablentermediates in some paramagnetic reactions carried out incidic solutions. Furthermore, the mechanism by which the mul-ivalent oxidant oxidizes a substrate depends not only on theubstrate but also on the medium used for the study.

Thiourea (NH2CSNH2), one of the simplest of the thio com-ounds, has many industrial applications [4,5]. Thiourea andts derivatives are known corrosion inhibitors [6]. Thiourea isoxic [7] and a cancer support agent [8]. These and other envi-

onmental concerns have promoted studies on the destructionf thiourea [9]. Thiourea can be oxidized by a wide variety ofxidizing agents [10–17]. The reaction pathways and the final

∗ Corresponding author.E-mail address: drkhanchem@yahoo.co.in (Z. Khan).

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927-7757/$ – see front matter © 2007 Elsevier B.V. All rights reserved.oi:10.1016/j.colsurfa.2007.02.006

roducts of the oxidation reaction depend on the pH and theondition of the reaction mixtures.

Reactions involving sulfur-containing compounds are gener-lly known to be quite complex. This complexity is highlightedy free-radical mechanisms [18], autoxidations [19], and theormation of sulfur–sulfur bonds leading to various polymericulfur species [20].

A survey of literature shows that the oxidation of thioureay permanganate has not been investigated so far. We reportn this paper a kinetic study of the oxidation of thiourea byermanganate, which proceeds through the formation of stableater-soluble colloidal MnO2 as an intermediate.

. Experimental

.1. Materials

Double-distilled, and CO2-free water was used throughouthe studies. Potassium permanganate (oxidant), thiourea (reduc-

ant), sodium fluoride, and manganese(II) chloride (all E. Merck,ndia, 99%) were used without further purification. Perman-anate solutions were standardized by titration against oxalate21]. Perchloric acid (HClO4, Thomas Baker, 70% reagent) wassed for [H+].

hysicochem. Eng. Aspects 302 (2007) 102–106 103

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Fig. 1. Absorption spectra of reaction mixture containing [MnO4−] = 2.0 ×

10−4 mol dm−3; [thiourea] = 2.0 × 10−3 mol dm−3 after completion of reactionaf[

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The stability of colloidal MnO2 depends strongly onthe pH and nature of the reductant [33–36], and thewater-soluble colloidal MnO2 can exist in aqueous-neutral

Table 1Values of first-order rate constant for the oxidation of thiourea (=2.0 ×10−3 mol dm−3) by MnO4

− in HClO4 (=1.8 × 10−4 mol dm−3) at 40 ◦C

[MnO4−] (×104 mol dm−3) kobs (×104 s−1)

1.4 2.7

S.A. Khan et al. / Colloids and Surfaces A: P

.2. Kinetic procedure

Kinetic experiments were carried out in a temperature-ontrolled (±0.1 ◦C) water bath. The reaction was initiated bydding the required quantity of thiourea to a thermally equi-ibrated mixture of permanganate, perchloric acid, and othereagents. The zero time was taken when half of thiourea solu-ion had been added. The course of the reaction was followed by

easuring the absorbance of the remaining colloidal MnO2 at25 nm on a Bausch & Lomb Spectronic-20 spectrophotometer.he oxidation was carried out under pseudo-first-order condi-

ions with at least a ten fold excess of reductant over oxidantoncentration. The pseudo-first-order rate constants were calcu-ated from the slopes of plots of log(absorbance) versus time.ther details of the kinetic procedure are the same as described

lsewhere [22–25].

.3. Product analysis

At the completion of the reaction between thiourea andotassium permanganate, 200 cm3 of ethanol and 60 cm3 of con-entrated HCl were added to the reaction mixture. It resultedn the formation of white crystals of formamidine disul-de. The compound was identified by the reported method26].

. Results and discussion

.1. General consideration

During the reduction of MnO4− by the thiourea at room

emperature (=30 ◦C), the reaction mixture becomes pinkλmax = 525 nm) to brown immediately. In order to confirmhe nature of brown colour, the spectrum of the reaction mix-ure containing thiourea (=2.0 × 10−3 mol dm−3) and MnO4

−=2.0 × 10−4 mol dm−3) was recorded at the end of the reac-ion (Fig. 1). It is well known that if the brown color is due tohe formation of water-soluble colloidal MnO2 as an interme-iate, the spectrum will mainly due to the scattering of lightRayleigh’s law [27,28]: A = C/λ). The plot of log A versus log λ

as linear with slope (=−5.5) (Inset-Fig. 1). These results aren good agreement with the observations of Freeman et al. [28]nd Perez-Benito et al. [29]. Thus, we may safely conclude thatolloidal MnO2 was formed as an intermediate during the reduc-ion of MnO4

− by thiourea, which can be represented by theollowing equation:

ermanganate + thioureafast−→colloidal MnO2 (1)

.2. Stoichiometry

The stoichiometry of the reaction was determined usingpecterophotometric titrations. Several reaction mixtures

ith [oxidant] > [reductant] at constant [H+] (=1.8 × 10−4

ol dm−3) were prepared and kept for 30 min at 40 ◦C. Theesidual absorbance of colloidal MnO2 was monitored at 425 nmhere there was no absorbance from the Mn(II) species. The

1122

t 40 ◦C. Inset: plot of log(absorbance) vs. log λ for the product (MnO2) obtainedrom the MnO4

− oxidation of thiourea. [MnO4−] = 2.0 × 10−4 mol dm−3;

thiourea] = 2.0 × 10−3 mol dm−3; temperature = 40 ◦C. Slope = −5.5.

toichiometry for the oxidation of thiourea by MnO2 is 2:1thiourea:MnO2).

.3. Effect of [MnO4−]

The effect of [MnO4−] on the reaction rate was stud-

ed at constant [thiourea] (=2.0 × 10−3 mol dm−3), [HClO4]=1.8 × 10−4 mol dm−3) and temperature (=40 ◦C). It wasbserved that kobs decreased as the initial [MnO4

−] increasedTable 1). It is well established that pseudo-order rate constantre independent on the initial reactant concentration in defect.he abnormal behavior is due to possible coagulation (floccula-

ion) of the colloidal MnO2 particles (intermediates). This typef behavior (dependent of pseudo-first-order rate constant on thenitial concentration of the reactant in defect) has been observedn many permanganate reactions, and especially for those reac-ions, which has an autocatalytic reaction path [30–32].

.4. Effect of [HClO4]

.6 1.8

.8 1.8

.0 0.9

.2 0.9

104 S.A. Khan et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 302 (2007) 102–106

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aitFsrf4[i[thiourea] is linear passing through the origin indicating first-order dependence on [thiourea]. Comparison of Figs. 4 and 5clearly suggest that [H+] is altered during the course of reaction.

ig. 2. Effect of [HClO4] on kobs for the MnO4− oxidation of thiourea.

MnO4−] = 2.0 × 10−4 mol dm−3; temperature = 40 ◦C; [thiourea] = 3.0 (A),

.0 × 10−3 mol dm−3 (B).

edium [37,38]. The rate constant, obtained as a functionf [HClO4] at constant [MnO4

−] (=2.0 × 10−4 mol dm−3),thiourea] (=2.0 × 10−3 mol dm−3) and temperature (=40 ◦C),as found to increase with increasing [HClO4] (Fig. 2).

.5. Effect of temperature

In order to calculate the values of activation parame-ers, the reaction was studied at three different tempera-ures, viz., 40, 50, and 60 ◦C at [thiourea] (=2.0 × 10−3

ol dm−3), [MnO4−] (=2.0 × 10−4 mol dm−3) and [HClO4]

=1.8 × 10−4 mol dm−3). The value of rate constants was foundo increase with temperature. The plot of log kobs versus 1/T wasinear (Fig. 3). The value of Ea (activation energy) was calcu-ated from the slope of Fig. 3, and is recorded in Table 2 alongith other parameters.

.6. Effect of [thiourea]

The effect of [thiourea] on the reaction rate was stud-ed at constant [MnO4

−] (=2.0 × 10−4 mol dm−3), [HClO4]=1.8 × 10−4 mol dm−3) and two different temperatures i.e., 40

able 2alues of first-order rate constants at different temperatures and activa-

ion parameters (Ea, �H#, �S#) for the oxidation of thiourea (=2.0 × 10−3

ol dm−3) by MnO4− (=2.0 × 10−4 mol dm−3), [HClO4] (=1.8 × 10−4

ol dm−3)

emperature (◦C) kobs (×104 s−1)

0 0.90 2.60 6.9

ctivation parametersEa (kJ mol−1) 80�H# (kJ mol−1) 78�S# (J K−1 mol−1) −66

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ig. 3. Arrhenius plot for the oxidation of [thiourea] = 2.0 × 10 mol dmy MnO4

− = 2.0 × 10−4 mol dm−3 in the presence of [HClO4] = 1.8 × 10−4

ol dm−3.

nd 50 ◦C. The values of rate constant increased with increasen [thiourea] (range: 1.0–3.0 × 10−3 mol dm−3) and decrease inhe range of 4.0–6.0 × 10−3 mol dm−3 (Fig. 4). Inspection ofig. 4, clearly suggest that effect of [H+] and [thiourea] operatesimultaneously during the course of reaction. In order to see theole of [thiourea] separately, some kinetic runs were also per-ormed in presence of sodium acetate–acetic acid buffer at pH.0. It was observed that reaction rate increase with increasingthiourea] at constant buffer. These results are depicted graph-cally in Fig. 5 as kobs–[thiourea] profiles. Plot of kobs versus

ig. 4. Effect of [thiourea] on kobs for the MnO4− oxidation of thiourea.

MnO4−] = 2.0 × 10−4 mol dm−3; [HClO4] = 1.86 × 10−4 mol dm−3; tempera-

ure = 50 (A), temperature = 40 ◦C (B).

S.A. Khan et al. / Colloids and Surfaces A: Physicochem. Eng. Aspects 302 (2007) 102–106 105

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ig. 5. Effect of [thiourea] on kobs for the MnO4 oxidation of thiourea inresence of sodiumacetate–acetic acid buffer. [MnO4

−] = 2.0 × 10−4 mol dm−3;H = 4.0.

.7. Identification of the intermediates and formation ofolloidal MnO2

It has been established that the oxidation of organic substratesy permanganate is characterized by two principal processes:

n(VII) + organic reductant(s) → intermediate,

ntermediate → Mn(II)

here intermediate = one (or several) reaction intermediate(s).In the case of MnO4

− as an oxidizing agent, the possiblentermediates are Mn(VI), Mn(V), Mn(IV), and Mn(III). Theresence of Mn(VI) and Mn(V) is ruled out by the fact that theyre highly unstable in an acidic medium.

Therefore, in order to gain further insight into the mech-nistic aspects and the role of hydrogen ions, a series ofinetic experiments were carried out at constant [MnO4

−]=2.0 × 10−4 mol dm−3), [thiourea] (=2.0 × 10−3 mol dm−3)nd different temperatures (=40 and 50 ◦C). The rate of for-ation and decomposition of colloidal MnO2 increased with

he increase in [HClO4]. On the other hand, at higher [HClO4]≥1.0 mol dm−3), the formation and decomposition of colloidal

nO2 was not observed. Thus, in presence of HClO4, the inter-ediate colloidal MnO2 is unstable and undergo acid hydrolysis.

n order to see the role of reduction product of permanganate.e. Mn(II), the oxidation kinetics of thiourea was carried out inresence of externally added Mn(II). Fig. 6 illustrates the depen-ence of the initial reaction rate on [Mn(II)], indicating sigmoid

ependence of kobs on [Mn(II)]. Our results seem to indicate thatn(II) is a true active autocatalyst (the Mn(II) formed as reac-

ion product contributes to accelerate the oxidation of thioureay MnO4

−). r

ig. 6. Effect of [F−], [Mn2+] on kobs for the oxidation of [thiourea] = 2.0 ×0−3 mol dm−3 by [MnO4

−] = 2.0 × 10−4 mol dm−3 in the presence ofHClO4] = 1.8 × 10−4 mol dm−3.

It is well known that Mn(III) species is reasonably stable onlyn acidic solution (pH < 3.0) and disproportionate in solutionsf higher pH [39]. Therefore, to substantiate the formation ofn(III) as an intermediate during the present reaction, different

mounts of sodium fluoride were added to the reaction mixture.he values of kobs increase with increasing [F−] [40,41]. This

ndicates that Mn(III) is also formed as an intermediate in theeduction of permanganate by thiourea.

.8. Mechanism

On the basis of experimental results the following mechanisman be proposed:

(2)

(3)

(4)

omplex 2k−→radical + Mn(III) (5)

n(III) + thioureafast−→Mn(II) + radical (6)

(7)

In the above scheme, equations (3) and (4), respectively, rep-esent the adsorption of thiourea and hydrogen ion on the surface

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06 S.A. Khan et al. / Colloids and Surfaces A:

f colloidal MnO2 through the electrostatic interaction. Eq. (5)s a one-step, one electron oxidation-reduction, rate-determiningtep. This reaction results in the formation of other intermediate,.e. Mn(III). The rate law consistent with the above scheme maye expressed by the following equation:

obs = k Ka Kad K′ad[H+]2[thiourea]T

(1 + Ka[H+])(8)

ccording to Eq. (8), the plots of kobs versus [thiourea] should beinear passing through the origin with positive slope at constantH+]. Such plots have been realized in the present study (Fig. 5).q. (8) also predicts that the plot of kobs versus [HClO4] at con-tant [thiourea] should be non-linear (Fig. 2) of upward concaveature, indicating complex kinetic behavior with respect to [H+].

. Conclusion

The important finding of this study is the fast reductionf permanganate by thiourea leading to the formation of col-oidal MnO2 as a stable intermediate. We are unaware of anyrecedence in the redox chemistry of permanganate-thioureaystem. The linearity in the plot of log absorbance versus log λ

learly indicates that the reaction proceeds through the forma-ion of water-soluble colloidal MnO2. We have also observedhat Mn(III) also formed as an other autocatalytic intermediate.inally, we can state that, when a excess of thiourea over per-anganate is used, further oxidation of the intermediate occurs

o yield Mn(II) as the final product. In presence of externallydded Mn(II), the reaction system become more complicatedue to the oxidation of Mn(II) by permanganate.

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