Indian Journal of Chemistry Vol. 48A, February 2009, pp. 189-193
Kinetics of permanganate oxidation of synthetic macromolecule poly(vinyl alcohol) Maqsood Ahmad Malik, Mohammad Ilyas & Zaheer Khan*,† a
Department of Chemistry, Jamia Millia Islamia (Central University), Jamia Nagar, New Delhi, 110 025, India Email:
[email protected] Received 22 July 2008; re-revised and accepted 27 December 2008
Oxidation of poly(vinyl alcohol) (PVA) by permanganate has been studied spectrophotometrically at 525 and 420nm. Under pseudo-first order conditions ([PVA] >> [MnO4-]), the reaction rate increases with [PVA] and the kinetics reveals complex order dependence on [PVA]. The second-order kinetics with respect to [H2SO4] at low concentration shifts to firstorder at higher concentration. Water soluble colloidal MnO2 has been identified as an intermediate in the reduction of MnO4by PVA. The hydrogen ions decrease the stability of colloidal MnO2. Poly(vinyl ketone) is found to be the final oxidation product of PVA. Inorganic electrolytes like NaF, Na4P2O7 and MnCl2 (a product of the reaction), have inhibitory and composite effects (inhibition and catalysis) on the reaction rate. Arrhenius and Eyring equations have been used to evaluate the activation parameters. The observed results are discussed in terms of Michaelis-Menten kinetic model. A mechanism has been proposed on the basis of experimental findings. Keywords: Kinetics, Reaction mechanisms, Oxidations, Poly(vinyl alcohol), Permanganate oxidations IPC Code: Int Cl.8 C07B33/00
The science of polymer (macromolecule) behavior in solution is a rapidly developing field with regard to their wide applications in the preparation of colloidal and nano-size metal particles1-5. Stabilizers (polymers, surfactants, complexing agents) have a significant effect on orientation and interaction of metal particles1-7. It has been established that the most important role of the stabilizing polymer is to protect the nanoparticles from coagulation. Poly(vinyl alcohol) is one of the favoured macromolecule for carrying out synthesis of polymer-nanoparticle composite8. Preparation of colloidal- and nano-metal particles depends upon the reduction of metal salt by suitable reducing agent as well as the presence of stabilizers9-12. PVA and its oxidation product (polyvinyl ketone) have reducing and complexing properties, respectively, towards various transition metal ions13-15. It has been established5,6,16 that secondary –OH groups in the macromolecule skeleton are responsible for the oxidation of PVA by permanganate and chromic acid in aqueous alkaline and acidic media. It has been reported on several occasions that the possible intermediate(s) species are Mn(VI), Mn(V), Mn(IV), _________ † Present address: Department of Chemistry, Faculty of Science, King Abdul Aziz University, P.O. Box 80203, Jeddah, 21589, Saudi Arabia
and Mn(III) if permanganate serves as an oxidizing agent in alkaline, neutral and /or acidic medium14,17,18. Mn(II) (reaction product) acts as an autocatalyst. Permanganate is used as oxidizing agent for removing organic molecules and heavy metals from the nuclear wastes19. Due to this, permanganate oxidations provide chemical kinetics with challenging mechanism. As a part of our investigations on the preparation, characterization and catalytic behavior of water soluble colloidal MnO29,10,20-23, we report herein the kinetics of the formation of water soluble colloidal MnO2 as an intermediate during the permanganate oxidation of poly(vinyl alcohol) in presence of H2SO4 medium and examine the inferences drawn from the inorganic electrolytes (NaF, Na4P2O7 and MnCl2) trapping experiments. Materials and Methods CO2-free, deionized and doubly distilled water (first distillation from alkaline permanganate) was used as solvent for the preparation of stock solutions of all the reagents (PVA, potassium permanganate, sodium fluoride, sodium pyrophosphate and manganese(II) chloride; all from Merck, India, 99%). Permanganate and PVA solutions were prepared by literature methods13. Permanganate solutions were standardized by titration against a standard sodium
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oxalate solution and stored in a dark glass bottle. Sulphuric acid was used to maintain [H+] of the reaction mixture and solutions were standardized by titration with sodium hydroxide to the phenolphthalein end point. Reaction mixture containing permanganate and other reagents (whenever necessary) was thermally equilibrated at desired temperature (25 ± 0.1°C), and to this was added a solution of PVA, pre-equilibrated at the same temperature. The reactions were started in glass-stoppered two-necked flask fitted with a double walled condenser to arrest the evaporation. The zero time was taken when half of the PVA solution had been added. The progress of PVA oxidation was followed spectrophotometrically by monitoring the absorbance due to unreacted permanganate at different time intervals at 525 nm using Spectronic21D spectrophotometer with cell of path length 5 cm. During the entire study, the reactions were carried out under pseudo-first order conditions using an excess of [PVA] over [MnO4¯]. The pseudo-first-order rate constants (kobs, s−1) were determined from the plots of log (absorbance) versus time. Other details of the kinetic measurements were the same as described elsewhere22, 23. Results and Discussion Reaction-time curves shows the changes in the log(absorbance) of MnO4−-PVA system with definite time intervals in H2SO4 media. The linearity in these plots implies that the reaction is first-order with respect to [MnO4−]. Attempts were made to determine the stoichiometry of the redox reaction by spectrophotometric titrations. Absorbance of solution containing different initial [MnO4−] and a constant [PVA] were measured after 1 h at 25°C. Results obtained by this method indicate that 2.5 moles of PVA was consumed per mole of permanganate. OH + 2MnO4 + 5 CH2 CH + 6H -
_
2+
2Mn
5
O CH2_ C + 8H2O
…(1) The colour of the reaction mixture changed from purple to colourless at the end of the reaction. The oxidation product of PVA was confirmed as follows: a solution of PVA was added slowly to the acidic solution of permanganate at 25°C. After completion of the reaction, a saturated solution of 2,4-dinitro-
phenylhydrazine in 2M HCl was added to the reaction mixture and was left overnight in a refrigerator. The yellow precipitate was filtered, washed and dried. The ketone was identified by infra red spectrum of the yellow precipitate, which showed carbonyl stretching at 1725 cm-1. It is well known that one –OH group of secondary alcohols oxidized in to one C=O (carbonyl group) with the same number of carbon atoms. Therefore, n number of –OH group of PVA are responsible for the oxidation which leads to the formation of n number of keto-derivative with the same number of carbon atoms. It was also observed that the ketone does not undergo further oxidative degradation under our kinetic experimental conditions because a spot test for carboxylic acid was negative. To confirm the intervention of free radical as an intermediate, a solution of permanganate was added to a mixture of PVA and saturated solution of mercuric(II) chloride. The formation of precipitate was observed which suggests that free radicals responsible for the reduction of Hg(II) ions were produced in this system. In the first set of experiments, the oxidation of PVA was studied as a function of [MnO4−] at constant PVA (0.20g dm−3), H2SO4 (1.47 mol dm−3) and temperature (25°C). MnO4− concentration varied in the range of 1.0 × 10−4 to 3.0 × 10−4 mol dm−3. The kobs values were found to be 3.6 ± 0.2 ×10−4 s−1 at [MnO4−] values of 1.0, 1.4, 2.0, 2.4 and 3.0 ×10−4 mol dm−3. The reaction follows first-order kinetics with respect to [MnO4−]. The rate law is, therefore, represented by Eq. (2):
−
d [MnO −4 ] = kobs [MnO 4− ]T dt
…(2)
In the second set of experiments, the effect of [PVA] was studied at constant MnO4− (2.0 ×10−4 mol dm−3) and temperature (25 ºC) at three different [H2SO4] (0.71, 1.10 and 1.47 mol dm−3). The values of rate constant increased with increase in [PVA] (Table 1). The double logarithmic plots between kobs and [PVA] gave two straight lines with different positive slopes. The study on effect of [PVA] indicates that order with respect to [PVA] decreases with increasing [H2SO4], suggesting involvement of H+ prior to the electron-transfer step (the rate determining step). On the other hand, the plots of 1/ kobs versus 1/[PVA] were linear with a positive intercept on the y-axis at constant [H2SO4] (Fig. 1). Such plots are indicative of Michaelis-Menten
MALIK et al.: KINETICS OF PERMANGANATE OXIDATION OF POLY(VINYL ALCOHOL)
behaviour (kinetic proof for complex formation between PVA and MnO−4). The effect of [H2SO4] on the oxidation rate was also studied by carrying out the reaction at solutions of different [H2SO4]. At PVA = 0.20 g dm−3, MnO4− = 2.0 × 10−4 mol dm−3 and temperature = 25ºC, the reaction rate increased with increase in [H2SO4] (kobs × 104 = 0.0, 0.4, 1.1, 1.8, 2.4, 3.6 and 4.9 s−1 at [H2SO4] = 0.0, 0.32, 0.73, 0.82, 1.10, 1.47 and 1.83 mol dm−3, respectively). Due to high acidity employed in this study, no attempt was made to keep the ionic strength constant. The change in the rate with acidity may be explained by the Zucker-Hammett hypothesis, where log kobs is related to either –Ho, the Hammett acidity function, or to log[acid]. The order of the reaction was found to be 2.1 and 1.0 at lower (0.45 to 1.1 mol dm−3) and higher (1.0 to 1.47 mol dm−3) [H2SO4], respectively, indicating the role of H+ before the rate determining step. The effect of [Na2SO4] was Table 1 — Pseudo first-order rate constants (kobs) for the oxidation of PVA by permanganate (2.0 × 10−4 mol dm−3) in the presence of H2SO4 at 25ºC [PVA] (g dm−3)
104 kobs (s−1) at [H2SO4] (mol dm−3) = 0.71 1.1 1.47 0.0 0.0 0.0 0.0 0.8 1.7 0.4 1.3 2.9 0.6 1.8 3.3 1.0 2.4 3.6 1.1 2.7 4.5 1.2 2.9 5.9 1.2 3.4 7.8
0.0 0.05 0.10 0.15 0.20 0.30 0.40 0.50
1.2
4
10 /kobs(s)
0.9
0.6
0.3
0.0 0
4
8
12 -1
16
20
3
1/[PVA] (g dm )
Fig. 1 — Double reciprocal plots between kobs and [PVA]. [Reaction cond.: [MnO−4] = 2.0 × 10−4 mol dm−3; [H2SO4] = 1.1 (●) and 1.47 mol dm−3 (▲), Temp. = 25°C].
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also investigated under the same kinetic experimental conditions in presence of H2SO4. The kobs values were found to be 3.6, 3.5, 3.7 and 3.6 × 104 s−1 at 105 [Na2SO4] of 2.0, 3.0, 4.0 and 5.0 mol dm−3. These findings indicate that the reaction rate remains the same in the presence of SO42−. In order to evaluate the various activation parameters (activation energy, enthalpy and entropy of activation), a series of experiments were performed at four different temperatures, viz., 25, 35, 45 and 55ºC at different [PVA] with constant [MnO4−] and [H2SO4]. The values of rate constant (k) were calculated from the intercepts of the plots of 1/kobs versus 1/[PVA] and found to be 3.5, 5.8, 12.8 and 27.1 × 104 s−1 at temperature = 25, 35, 45 and 55ºC, respectively. Activation parameters (Ea = 55 kJ mol−1, ∆H# = 53 kJ mol−1 and ∆S# = -127 J K−1 mol−1) were calculated by Arrhenius (log k versus 1/T) and Eyring (logk/T versus 1/T) plots. Interestingly, the magnitude of enthalpy of activation (∆H#) is significantly same (53 and 56 kJ mol−1) for the PVA oxidation by acidic (present work) and alkaline MnO4−, respectively13, showing that the same transition state is formed in both media. The large negative value of ∆S# indicates the existence of compact activated state stabilized by strong hydrogen bonding and large solvation in the electron transfer step in acidic media. Fluoride and pyrophosphate ions are efficient trapping agents for Mn(III). In order to confirm the formation of Mn(III) as an intermediate during the reduction of permanganate by PVA, different sets of experiments were also performed in presence of P2O74− and F− ions. At constant [PVA], [MnO4−], [H2SO4] and temperature, the kobs decreased continuously with increase in [NaF and Na4P2O7] (kobs × 104 = 3.6, 2.3, 2.4, 2.2, 1.2, 1.1 and 3.6, 0.9, 1.0, 1.0, 1.1, 1.1s−1 at [NaF] and [Na4P2O7] = 0.0, 1.0, 2.0, 3.0, 4.0, 5.0 × 10−3 mol dm−3, respectively). The decreasing kobs may be due to complex formation between Mn(III) and F− and P2O74− ions. As a result, the free [Mn(III)] decreases at the reaction site. Thus, we may safely conclude that the Mn(III) is formed as an intermediate during the course of reaction. However, when we attempted to monitor the formation of Mn(III) at 470 nm, we failed to detect any build-up of Mn(III). The kinetic experiments were also performed in presence of externally added Mn(II) at constant concentrations of other reagents. Initially, there was a very sharp decrease in kobs, followed by gradual increase with increasing [Mn(II)]
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(kobs × 104 = 3.6, 2.3, 4.1, 5.2, 6.8 and 9.1 s−1 at [Mn(II)] = 0.0, 0.2, 0.6, 1.0, 1.4 and 2.0 × 10−3 mol dm−3, respectively). It was observed that the oxidation rate decreases with [Mn(II)] while at higher [Mn(II)], the results are opposite (kobs increases with increasing [Mn(II)]). In presence of [Mn(II)], there is competition between Mn(II) and PVA to react with MnO−4 and its intermediate state, i.e., Mn(IV). Thus, we may conclude that Mn(II) acts as an autocatalyst. Due to the complicated features of the reaction, the exact dependence of kobs on [Mn(II)] could not be estimated in presence of PVA. As far as the formation of Mn(IV) is concerned, the oxidation of PVA was also studied by monitoring the absorbance at 420 nm, where the contribution from MnO4− is negligible. Under our experimental conditions, (H2SO4 > 0.71 mol dm−3), the formation of Mn(IV) was not observed. These observations are not unexpected; in presence of high acid concentration, the Mn(IV) is unstable and undergoes acid hydrolysis24. On the basis of above results and discussion the mechanism shown in Scheme 1 is proposed for the oxidation of PVA by permanganate. In Scheme 1, Eq. 3 represents protonation of MnO4. The next reaction (Eq. 4) shows formation of a complex between PVA and permanganate. In analogy with previous studies10,22, we assume that the complex undergoes a one-step, two electron oxidationreduction mechanism to Mn(V) and oxidation product of PVA (corresponding ketone). Equation 5 is the
rate-determining step in the redox process. Mn(V) is highly unstable in an acidic medium with respect to disproportionation25 and immediately gets converted into Mn(IV) (Eqs 6 and 7) (this species is commonly involved in the oxidation of organic reductants by MnO4-). In presence of a large amount of PVA, the intermediate(s), Mn(IV) and/or Mn(III), immediately gets converted into the stable products (Eqs 8 and 9). According to Scheme 1, the overall reaction rate is given by Eq. (10). −
kK c K a [H + ]2 [PVA][MnO −4 ] d [MnO 4− ] = dt 1 + K a [H + ] + K a K c [H + ][PVA]
…(10)
kK c K a [H + ][PVA] 1 + K a [H + ] + K a K c [H + ][PVA]
…(11)
− kobs =
The equilibrium in step 3 lies well to the right and hence, the inequality 1> Ka [H+] will evidently exist and Eq. (11) is reduced to Eq. (12), which explains the experimental results, i.e., complex- and fractionalorder dependence on [H+] and [PVA], respectively. − kobs =
kK c K a [H + ]2 [PVA] 1 + K a K c [H + ][PVA]
…(12)
The inequality 1 > Ka Kc [H+] [PVA] is evident at low [H2SO4], and the rate law, Eq. (12), then reduces to Eq. (13), which explains the second-order dependence of the reaction on [H2SO4] at lower concentrations. kobs=kKcKa[H+]2[PVA]
…(13)
The above inequality will be valid in the reverse direction at higher [H2SO4], that is, Ka Kc [H+] [PVA] > 1 will hold, and Eq. (12) is reduced to Eq. (14), which explains the first-order dependence at higher concentrations of [H2SO4] at constant [PVA]. kobs=k[H+]
Scheme 1
…(14)
Thus, Eqs (13) and (14) correspond to two extreme conditions between which the reaction order should vary; between two to unity and one to zero with respect to [H2SO4] and [PVA], respectively. The rate law, Eq. (12), which is consistent between the extreme conditions of Eqs (13) and (14) is verified in the form of Eq. (15).
MALIK et al.: KINETICS OF PERMANGANATE OXIDATION OF POLY(VINYL ALCOHOL)
1 1 1 = + + 2 kobs kK c K a [H ] [PVA] k[H + ]
…(15)
The plots of 1/kobs versus 1/[PVA] at each acidity shows the expected linear relationship with positive slopes and positive intercepts on the y-axis (Fig. 1). The values of k and Kc have been calculated using the values of Ka (= 2.99× 10−3 mol−1dm3 ) , intercept and slope of Fig. 1 and found to be 3.5 × 10−4 s−1 and 9.7 × 103 mol−1dm3, respectively, at constant [H2SO4] (= 1.1 mol dm−3). As predicted, the intercepts and slopes of the plots were found to be dependent on [H2SO4]. Inspection of Fig. 1 clearly suggests that these plots deviate from linearity at higher [PVA] at each [H2SO4]. In solutions of high acidities, however, the oxidation of PVA is not as simple as indicated by Eqs (4) and (5). The oxidation rate and order of the reaction markedly increase at high [PVA] as well as acidity of the medium. References 1 Napper D H, Polymeric Stabilization of Colloidal Dispersions (Academic Press, London) 1983. 2 Platzer O, Amblard J, Marignier J L & Belloni J, J Phys Chem, 96 (1992) 2334. 3 Toshima N, in Fine Particles Sciences and Technology-From Micro-to New Particles, edited by E Pellizzetti (Kulwer, Dordrecht) 1996, p. 371.
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