Spectral Study of Ruthinium (III) Catalyzed Oxidation of Maltose by Potassium Permagnate in Acidic Medium

Introduction

Permanganates ions oxidize a greater variety of substrates and find extensive applications in organic syntheses,^1–7 especially after the advent of phase transfer catalysis,^3,4,6 which permits the use of solvents such as methylene chloride and benzene. Kinetic studies constitute an important source of mechanistic information on the reaction, as demonstrated by results referring to unsaturated acids in both aqueous^1,3,7,14-17 and non-aqueous media^.8. During oxidation by permanganate, it is evident that the Mn(VII) in permanganate is reduced to various oxidation states in acidic, alkaline and neutral media. Furthermore, the mechanism by which this multivalent oxidant oxidizes a substrate depends not only on the substrate but also on the medium⁹ used for the study. In strongly alkaline media, the stable reduction product^10,11 Potassium permanganate has been used to oxidize various compounds in both acidic and alkaline medium ^[1-6]. Little attention has been paid however, to the reactivity of KMnO₄ in the presence of catalyst ^[7,8] and nearly no investigation has so far been reported on the catalytic role of ruthenium(III) chloride with potassium permanganate  as an oxidant in acidic medium. This fact prompted us to undertake the present investigation namely, “Mechanistic study of ruthenium(III) catalyzed oxidation of maltose by acidic potassium permanganate”.

Experimental

Devices

The change in absorbance with respect to time was done by  Shimadzu UV 1800-Double beam-UV–Vis spectrophotometer and IR studies were performed by Shimadzu8400 S- FT IR,

Materials

An aqueous solution of Maltose (BDH), potassium permanganate (CDH, AR), KNO₃ and were prepared by dissolving the weighed samples in triple distilled water. H₂SO₄ acid (SD.fine) was used as a source of H⁺ ion. The Ru(III) solution was prepared by dissolving a known weight of RuCl₃ (SDFine-Chem) in HCl (0.20 mol dm^–3). Hg was added to the Ru(III) solution, to reduce any Ru(IV) formed during the preparation of the Ru(III) stock solution, which was set aside for 24 h. The Ru(III) concentration was then assayed by EDTA titration.^15.All other reagents were of analytical grade and their solutions were prepared by dissolving the requisite amounts of the samples in doubly distilled water. H₂SO₄ and KNO₃ were used to provide the required acidity and to maintain the ionic strengthrespectively. Reaction vessels were painted black so as to prevent photochemical decomposition, if any.

Kinetics

All kinetic measurements were performed under pseudo-first order conditions with excess of [maltose] over [MnO₄ ^– ] at a constant ionic strength of 0.30 mol dm^–3. The reaction was initiated by mixing previously thermostated solutions of MnO₄ ^– and maltose, which also contained the necessary quantities of Ru(III), H₂SO₄ and KNO₃, to maintain the required acidity and ionic strength respectively. The temperature was uniformly maintained at 30 ± 0.1°C. The course of reaction was followed by monitoring the decrease in the absorbance of MnO₄^– in a 1 cm quartz cell of Shimadzu UV 1800-Double beam-UV–Vis spectrophotometer , at its absorption maximum of 525 nm as a function of time. The application of Beer’s law to permanganate at 525 nm is verified, giving Ɛ=2250±25dm³ mol^–1 cm^–1 (literature Ɛ = 2200 dm³ mol^–1 cm^–1). The first-order rate constants, were evaluated by the plots. The course of reaction was followed by recordinMnO₄^– at its absorption maximum, 525 nm, as a function of time⁹. It was verified that there was no interference from other reagents at this wavelength. The absorption spectra of the reactants and typical traces of the spectral changes are shown in fig A.

Scheme 1  Click here to View scheme

 

Figure A: Spectroscopic changes occurring in the Ru(III) catalysed oxidation of maltose by permanganate with[S] = 1.0 X  10–3, [MnO4] = 1×5 ´ 10–4 [Ru(III)] = 1×10–6, [H+] = 1×10–3, and  I = 0.30 mol dm–3 at 30°C, scanning time interval = 120 s. Click here to View figure

 

Results and Discussion

The stoichiometry of the reaction was ascertained by equilibrating the reaction mixture containing an excess of potassium permanganate. Estimation of unconsumed MnO₄^– revealed that one mole of substrate consumed four  moles of MnO₄^–

S  = maltose  P  =  Malturic acid

Ru(III) catalyzed oxidation of maltose was studied at several initial concentrations  (Table-1). The plots of unconsumed manganate versus time on varying concentrations showed first order w.r.t. [MnO₄^–] which is further confirmed by plotting a graph between (-dc/dt) vs [MnO₄^–] (Fig.1.1). The first order rate constant shows nearly the same value at different concentrations of Maltose, showing zero order dependence on  substrate concentrations. The reactions were observed to be first order with respect to Ru^III. This has also been confirmed by “least- square method” (Fig.3). A log (-dc/dt) versus log [H⁺] plot (Fig.4) was found to have  a slope value of +0.24 at 30⁰ C, indicating a positive fractional order  effect of [H⁺] on the reaction rate. Experimental data showed that the reaction was unaffected by chloride ion and ionic strength on the rate.

Table 1: Effect of variation of reactants on the reaction rates for maltose at 30⁰C

[Substrate] x 102 M	[KMnO4] x 103 M	Ru(III) x 106 M	[HClO4]  x 103 M	(-dc/dt) x 107 mol l-1 s-1
0.33 0.40 0.50 0.66 1.00 2.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00	1.00 1.00 1.00 1.00 1.00 1.00              0.28 0.33 0.40 0.50 0.66 1.00   1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00	96 96 96 96 96 96 96 96 96 96 96 96 48 96 144 192 240 288   96 96 96 96 96	1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 1.00 0.33 0.40 0.50 0.66 1.00 2.00	8.47 8.65 8.45 8.55 8.59 8.34 2.23 2.8 3.21 4.15 5.35 8.59 4.8 8.59 11 16.2 20.33 24.21 5.83 6.05 6.84 7.34 8.59 9.50

[KNO₃] =  0.30 M

Figure 1.1: Plot between [KMnO4] x 10-3 M vs. (-dc/dt) x 10-7 M L-1 S-1 for oxidation of Maltose at 30°C. Click here to View figure

 

Figure 1.2: Plot between Absorbance (remaining KMnO4) and time at 30°C Click here to View figure

 

Figure 1.3: Plot between remaining Log Absorbance-(Remaining KMnO4) and time at 300C Click here to View figure

 

Figure 2: Plot between [Rh(III)] x 10-5 M  vs. (-dc/dt) x 10-7  M L-1 S-1 for oxidation of Maltose at 30°C. Click here to View figure

 

Figure 3: Plot between log [Rh(III)]  vs. log (-dc/dt) for oxidation of Maltose at 300C. Click here to View figure

 

Figure 4: Plot between 4 + log [H+] vs.  8 + log (-dc/dt) for oxidation of Maltose at 300C.   Click here to View figure

 

The rate measurements were taken at 30⁰ – 45⁰C and a plot of log k vs I/T, which was linear (fig.5). In the light of the discussion regarding the reactive species of potassium permagnate in acidic medium and that of ruthenium trichloride in acidic medium it has been concluded beyond doubt that MnO₄^–  and [RuCl₂(H₂O)₄]⁺ are the actual reactive species of potassium permagnate and ruthenium trichloride respectively in acidic medium.Now on the basis of above results following mechanism can be given for the oxidation of aforesaid substrates:

Figure 5: Plot between 4 + log k vs. 1/T x 104 K-1 for oxidation of  Maltose at 300C.                      Click here to View figure

 

Considering the above slow steps and applying steady state treatment with a reasonable approximation, the rate law for Maltose may be  written in terms of consumption of  [MnO₄^–] as equation:-

where K₁   = (k₁/k_-1)             and     [Ru(III)]_T  = [C₁]  +  [C₂]

Spectral Evidence for the Formation of Complexes During Course of Reaction

In the oxidation of maltose with permanganate ion in acidic medium, order with respect to sugars concentration was first throughout its 10-fold variation. This shows that OH group of sugar involves in the oxidation before the rate determining step although it will combine with the most reactive complex to form the reaction products along with the regeneration of Ru catalyst .In order to probe the possible formation of complex MnO₄ ^– and Ru (III) spectra for the solution of MnO₄^–, H+ and for the solution of MnO₄^– , H+ and Ru(III)  have been collected Figures.

From the spectra, it is clear that the addition of Ru (III) solution there is a decrease of absorbance from 1.567 to 1.465 (λmax 525 nm). This decrease in absorbance can be considered as an indication of quick formation of the complex between reactive species of Ru (III) in acidic medium. When maltose and lactose solutions were added to the solutions of MnO₄ – , H⁺ and Ru(III)  a decrease in absorbance from 1.465   to 1.376 (λmax 525 nm)for maltose was noted (Fig B- Fig.D). This shows that the complex formed subsequently combines with maltose to give a highly solvated activated complex.

Figure B: Absorption Spectra of MnO4 -, H+  solutions. Click here to View figure

 

Figure C: Absorption Spectra of MnO4 –, H+ and Ru(III) solutions.  Click here to View figure

 

Figure D: Absorption Spectra of MnO4 –, H+ ,Ru(III) and Maltose solutions. Click here to View figure

 

Thus rate law (2) derived fully explains first order dependence of the reaction on ruthenium(III) chloride, zero order dependence of substrates viz. maltose in acidic medium. It also explains first order kinetics with respect to [MnO₄^–]

The rate law (2) also shows positive fractional order dependence on sulphuric  acid i.e. [H⁺] and negligible effect of potassium chloride is shown.There is negligible effect of ionic strength and Acetic acid on reaction.

The rate expression shown by equation (2) can also be written as:-

where K₁  =  (k₁ / k_-1)

With the help of equation (3) a graph plotted between (1/rate) and (1/[H⁺]), a straight line with positive intercept on Y axis (Fig. is obtained which helps us to calculate value of k₂ and slope of straight line helps to calculate K₁.

The rate law is in agreement with all observed kinetics. The proposed mechanism is in consistency with the activation parameters given in table-2.

Table 2: Activation parameters for acidic permanganate oxidation of Maltose

Substrate	Temperature coefficient	ΔEakJ mol-1	ΔH*J mol-1	ΔS*JK mol-1	ΔG*kJ mol-1	Log A	krmol-2 lit-2 sec-1
Maltose	2.09	37.39	37.3947	-214.04	64.879	1.809	63.64

 

Conclusions

In the present investigation it has been tried to experimentally explore various aspects of kinetics of oxidation of Maltose using Ruthenium (III) as a catalyst using potassium permagnate in acidic solution. The experimental results as shown reveal that the reaction rate doubles when the concentration of the catalyst Ruthenium (III) is doubled. The order w.r.t. potassium permagnate is also observed to first order.  There is fractional order w.r.t. H⁺ in reaction mixture. The rate law  is in conformity with all kinetics observations summarized there. The product is Malturic acid and is confirmed by TLC, Functional Group Analysis and also by FTIR Spectroscopy. The proposed mechanistic steps are supported by the negligible effect of ionic strength which indicates the involvement of an ion in the slow  and rate determining step. From the present investigation it is also concluded that MnO₄^–  and [RuCl₂(H₂O)₄]⁺ are the reactive species of KMnO₄ and Ruthenium(III) chloride respectively in acidic medium.The high positive values of energy of activation ,(∆G^⃰ ) indicates highly solvated transition state, while fairly high negative values of entropy of activation (∆S^⃰ ) suggest the formation of an activated complex with reduction in degree of freedom.

References 

1. Sh. Srivastava , R.K. Sharma ; Oxidation Communication , 2, 343-349 (2006).
2. Hassan R M, Mousa M A and Wahdan M H.; 1988., J Chem Soc, Dalton Trans, 605-609(1988.)
3. Sheila Srivastava  and Sarika Singh ; J. Indian Chem. Soc., 18 , 295 (2004).
4. Sheila Srivastava , Ashish Kumar and Parul Srivastava ;  J. Indian Chem. Soc.,  83 , 347-350 (2006).
5. Sheila Srivastava , Rajendra Kumar Sharma and Sarika Singh ; J. Indian Chem. Soc., 83 , 282-287 (2006).
6. Ashish Singh , Surya Singh , Ashok Singh , Bharat Singh ; Trans. Metal Chem. , 30(5) ,   610-615 (2005).
7. Iftikhar Ahmed , C. Mohammad Ashraf ; Int. J. Chem. Kinet. , 11 , 813-819 (2004).
8. Saik Sondu , Bangalore , Sethuram and Tageda Navaneeth Rao ; Transition Metal Chem. , 15(1) , 78-80 (1990).
9. Sheila Srivastava  ; Trans. Metal Chem. , 24(6) , 683-685 (1999).
10. Sheila Srivastava , Bharat Singh ; Trans. Metal Chem., 16 ,466 (1991).
11. N. Venkata Subramanian , V. Thigarajan ; Canadian J. Chem., 47 , 694 (1969).
12. J. C. Bailer , “ Chemistry of Coordination Compound ; Reinhold , New York (1964) , P-4.
13. C. S. Reddy , E. V. Sundaram ; J. Indian Chem. Soc., 62 , 209 (1985).
14. Sheila Srivastava , Ashish Kumar and Parul Srivastava ;  Oxidation communications, 29 , No. 3,660-666 (2006).
15. Reddy C S and Vijayakumar T 1995 Indian J. Chem. A34 615; 1970 Chem. Abstr. 72 50624
16. Sheila Srivastava , Ashish Kumar and Parul Srivastava ;  Oxidation communications, 29 , No. 3,622-627 (2006).
17. Sheila Srivastava , Ashish Kumar, Parul Srivastava, Lakshmi chaudhary and Shalini singh, J. Indian Chem. Soc., 86 , 1-5 (2009).

---

This work is licensed under a Creative Commons Attribution 4.0 International License.