Kinetics and Mechanism of Uncatalysed Oxidation of D-ribose by Cerium (IV) in Aqueous Acidic Medium

 

Manoj K. Ghosh* and Surendra K. Rajput

Department of Chemistry, Govt. Nagarjuna P.G. College of Science, Raipur, 492010 (C.G.), India

*Corresponding Author E-mail: mkghosh01@yahoo.co.in

 

 

ABSTRACT:

Kinetics and Mechanism of uncatalysed oxidation of D-ribose by cerium (IV) in aqueous acidic medium have been  investigated titrimetric method in the temperature range 308-333K. The reaction has been found to be first order with respect to D-ribose in an uncatalysed reactions. The effect of [HSO4-] has also been observed. The increase in ionic strength of the medium decreases the rate of uncatalysed reaction. A 1:2 stoichiometry is observed in the oxidation reaction. On the basis of the experimental results, a reasonable mechanism has been proposed. Rate equation derived from this mechanism can explain all the experimental results. From the effect of temperature on the reaction rate, the Arrhenius equation and various activation parameters have been computed.

 

KEY WORD: Kinetics,  uncatalysed, d-ribose, cerium (IV) , Arrhenius equation.

 


 

INTRODUCTION:

The products of oxidation of aldopenoses depend on the nature of the oxidant used and other experimental conditions. The oxidative degradation of organic and inorganic compounds with cerium(IV) are potentially interesting .Due to the specific coordination properties of cerium(IV), unique reactions of this ion can be expected. The kinetics of the oxidation of various monosaccharides by cerium(IV) has been reported on several occasions.

 

Chemistry of cerium is a very broad area which has received considerable attention through the years, resulting in substantial advance both in the synthetic1-2and mechanistic categories. Cerium(IV) is a well known oxidant in acidic media3-5 having reduction potential6-7of the couple CeIV/CeIII (1.70V) and is stable only in high acid concentration. In sulphuric acid and sulphate media, several sulphate complexes6,9,10 of  cerium (IV) form, but their role has not received much attention so far. Kinetic provides the most important indirect evidence in the mechanism. The metal ion oxidants have been widely employed in the synthetic chemistry11-12 including carbohydrates chemistry13-16. These are stable, in expensive and can readily be stored and handled.

 

The kinetic method of analysis have been widely developed and accepted in chemical analysis of  different samples17-18 and kinetic study of Ru (III) catalysed oxidation of sucrose by sodium periodate  in acidic medium have been studied by various scientists21-22. Literature survey indicate that, there is no reports on the ribose oxidation of aqueous  sulphuric acid medium by  cerium(IV).Hence we have investigate  the uncatalysed  oxidation of D-ribose  by cerium(IV) in order to understand  the behavior of active species of oxidant in sulphuric acid media and to propose a suitable mechanism. 

 

MATERIALS AND METHODS:

Chemicals

Chemicals of absolute pure quality were used without further purification. Stock solution of D-ribose and potassium bisulphate is prepared in double distilled water. Ce (IV) stock solutions were prepared by dissolving ceric ammonium sulphate in aqueous sulphuric acid. Sodium thiosulphate solution was standardized with standard iodine solution. All chemicals were purchased from E. Merck.

 

Experimental

Kinetic runs were performed in stopper glass vessels in a controlled temperature ± 0.1°C on water bath.  Requisite volume (90cm2) of all reagents, including substrate, were introduced into a reaction vessel and thermally equilibrated to 35 0C. A measured volume (10cm3) of cerium solution, also at 35 0C was rapidly poured into the reaction vessel. The kinetics of the reaction was studied under the conditions [Substrate]>>[Oxidant] in the case of Ce(IV).The kinetics were followed by estimating aliquots of the reaction mixture for cerium(IV) iodometrically, using starch  indicator18. Doubly recrystallised sugars(E. Merck) were used for the kinetic studies. Cerium solution was always made up and stored in black coated flask to prevent photochemical reaction. The solution was then standarised with sodium thiosulphate solution using starch indicator. Aqueous solutions of ribose were prepared fresh each day. All other reagents were of analytical grade. Conductivity water used throughout the study.

 

RESULTS AND DISCUSSION:

Kinetics of redox reactions between D-ribose and  cerium(IV) in sulphuric acid medium was followed under pseudo first order reaction at constant temperature. The oxidation of D-ribose by cerium(IV) in absence of catalysed  is normal. The order of reaction with respect to [Oxidant] was determined in (Table-1). Result shows that the rate constant is directly proportional to the concentration of  cerium(IV). Linear line is obtained in (Fig-1),,when we plot of  k1 v/s Ce(IV) concentration. This indicates the first order kinetics with respect to oxidant.

 

 

Table-1: Effect of variation of [Cerium(IV)], on the reaction rate

102[D-ribose]=5.00 mol dm-3; 102[H2SO4]=3.00 mol dm-3, 103[KHSO4]= 5.00 mol dm-3; Temp.= 308K

Run No.

103x[Ce(IV)] mol dm-3

K1x104(sec-1)

1

2

3

4

5

6

7

0.50

1.00

2.00

3.00

4.00

5.00

6.00

1.04

1.01

0.99

0.92

0.90

0.84

0.81

 

Figure-1

 

The order of reaction was determined at different concentration of substrate [D-ribose] and at fixed concentrations of other reactants in Table-2. Plot of k1 v/s D-ribose concentration is found to be a straight line (Fig.2a), which indicates that the rate of the reaction is directly proportional to the substrate concentration. The plot of log k1 v/s log [D-ribose] are linear. This indicates that the order with respect to [substrate] is one. (Fig.2 b)

 

Table-2:Effect of variation of [D-ribose], on the reaction rate

103[Ce(IV)]=3.00 mol dm-3; 102[H2SO4]=3.00 mol dm-3,

103[KHSO4]= 5.00 mol dm-3; Temp.= 308K

 

Run No.

102x[D-ribose] mol dm-3

K1x104(sec-1)

1

2

3

4

5

6

 

0.50

1.00

2.00

3.00

4.00

5.00

3.50

6.03

11.05

16.40

22.10

26.00

 

 

Figure-2(a)

 

Figure-2(b)

 

The effect of [H+]   ion concentration, on the rate of reaction, kinetic runs were carried out keeping the concentrations of all other reactants constant and varying the [H+] with sulphuric acid. The reaction has been carried out at various initial concentration of sulphuric acid tabulated in (Table-3). It has been observed that rate of reaction decreases with  increase of sulphuric acid concentration (Fig.3). The plot of kv/s 1/[H+] and log k1 v/s log [H+] was linear [Fig.3(a) and 3(b)].The result indicates that the order with respect to  [H+] is inverse first.

 

Table-3:Effect of variation of [H+], on the reaction rate

103[Ce(IV)]=3.00 mol dm-3;102[D-ribose]=5.00mol dm-3,

103 [KHSO4]=5.00 mol dm-3; Temp.= 308K

Run No.

102x[H2SO4] mol dm-3

K1x104(sec-1)

1

2

3

4

5

6

0.50

1.00

2.00

3.00

4.00

5.00

3.50

2.00

1.20

0.90

0.85

0.80

 

Figure- 3(a)

 

Figure-3(b)

 

The reactions were studied at different concentration of salt [KHSO4], while the other reactants are constant. The observations are given in (Table-4). The graphical plot of k1 v/s  [KHSO4]  is found to be a linear in nature (Fig.-4), which indicates that the rate of the reaction is inversely proportional to HSO4- ion concentration. Thus the addition of salts viz. KHSO4 did not have much effect on the rate of reaction.

 

 

Table-4:Effect of variation of [KHSO4], on the reaction rate

103[Ce(IV)]=3.00 mol dm-3; 102[D-ribose]=5.00 mol dm-3,

102[H2SO4]= 3.00 mol dm-3; Temp.= 308K

Run No.

103 x [KHSO4] mol dm-3

K1x104(sec-1)

1

2

3

4

5

6

7

0.50

1.00

2.00

3.00

4.00

5.00

6.00

3.50

3.00

2.75

2.30

1.90

1.50

1.30

Figure- 4(a)

 

Figure- 4(b)

 

To observe the effect of temperature on the reaction rate, the reaction was studied at different temperatures, while keeping all other reactants are constant in (Table-5). The kinetic data shows that the velocity of reaction increases with rise in temperature, showing the validity of the Arrhenius equation (Fig.5). So an attempt has been made to correlate the various activation parameters on the reaction mechanism.

 

Table-5:Effect of variation of Temperature, on the reaction rate

103[Ce(IV)]=3.00 mol dm-3;102[D-ribose]=5.00 mol dm-3, 102[H2SO4]=3.00 mol dm-3

Temperature in K

1/T x 10-3

K1x104(sec-1)

308

318

323

328

333

3.24

3.14

3.09

3.04

3.00

3.5

6.00

7.40

8.80

10.20

Kinetic and Activation parameters  for uncatalysed reaction

Parameter

D-ribose

Ea* (kJ mol-1)

ΔH* (kJ mol-1)

ΔS* (J mol-1)

ΔG* (kJ mol-1)

log A

36.50

33.81

-38.80

46.34

10.77

 

Figure-5.

 

Activation parameters

The result shows the value of activation energy (Ea) was found to be 36.5 kjmol-1.The value of enthalpy of activation (ΔH*) at 323 k is 33.81 kJmol-1, entropy (ΔS*) at 323 k is -38.8Jmol-1,frequency factor (A) at 323 k is10.77 and free energy(ΔG*) 46.34 kJmol-1.

 

In order to seen that the high positive value of change in free energy (ΔG*) indicates reaction is highly solvated transition state, while negative value of change in entropy (ΔS*) suggested, the formation of an activated complex with reduction in the degree of freedom of molecules. 

 

Reaction Mechanism

The kinetics of the forgoing reactions were studied and showed that substrates and oxidant interact in an equilibrium step to form an intermediate complex which is assumed to disproportionate forming a free radical and reduced Ce(IV). It is believed that involvement of both C1and C2 hydroxyls in a complex formation. On the basis of above statement and observed first order dependence on [oxidant] as well as  [substrate] a probable mechanism (Scheme-1) is proposed for the oxidation D-ribose such complex formation between the oxidant and substrate was observed in earlier studies.

 

 


Scheme- 1

 

Rate Law

The proposed mechanism involves the formation of complex in a reversible manner which reach with the substrate at rate determining steps to form [Ce(IV)-S] complex followed by a slow redox decomposition giving rise to aldoxide radical which oxidized by Ce(IV) rapidly.

 

The oxidation of D-ribose at different temperatures from 308K to 333K was studied. The rate of disappearance of cerium(IV) in this reactions increases sharply with increasing concentration of D-ribose. The plots of k1 against T were linear for uncatalyzed oxidation. The Arrhenius activation energy Ea for the uncatalyzed oxidation of D-ribose was 36.50 kJ mol-1.

 

The observed19stoichiometry of the reaction carried out by taking cerric ammonium sulphate in large excess as compared  to the carbohydrates in  different ratios and the observations were made for 36 hrs. at 350C.The reaction can be represented as in equation (1):

 

C5H10O5+2Ce(IV) +H2O=C4H8O4+ HCOOH+ 2Ce (III) + 2H+ -----(1)

D-ribose[S]            Aldotetrose     Formic acid

On the plot of 1/kobs against 1/[S] is made from which the constants 1/ks and k2/ksk1 are determined form the slop and intercept respectively. According to the equation mentioned in the above, when plot  between 1/kobs and 1/[S], a positive intercept would be observed which confirms the validity of the  mechanism and also the rate law.

 

CONCLUSIONS:

The proposed mechanism is well supported by the moderate values of energy of activation and thermodynamic parameters. The high positive value of the energy of activation (ΔG*) and (ΔH*) indicate that the transition state is highly solvated where as the negative value of entropy of activation (ΔS*) indicated that the activated complex is cyclic nature.

 

 

ACKNOWLEDGEMENT:

The authors are thankful to Dr. K.N. Bapat, Principal, Dr. S. Nigam, Professor and Head, for Lab facilities and also grateful to Dr. Sanjay Ghosh (Asstt. Prof.), Department of Chemistry, Govt. N.P.G. College of Science, Raipur.

 

REFERENCES:

1.       Mathew B, Narayana NV, Sreekumar and Vipin ; Microchim Acta, 144;2005: 291.

2.       Kolitsch V and Shwendtner K ; Acta.Cryst.,60C, 2004,89.

3.       Pol  PD, Katharic C P and Nandibewoor ST; Transition met. Chem., 27; 2002: 807.

4.       Thabaj K A, Chimatadar SA, and Nandibewoor T; Transition met. Chem.,31;2006:186.

5.       Chimatadar SA, Madawale S V and NandibewoorS T; Transition met. Chem., 32;2007:634.

6.       Day MC and Selbin J. Theoretical Inorg. Chemistry. Reinhold Pub.Corp., New York ,1964.

7.       Vogel AI. A text Book of Quantitative Inorganic analysis. Longmans, London, 1961; 3rd Ed.

8.       Pottenger CR, Johnson DC; J. of Polymer Science part A-1: Polymer Chem. 8(2) ; 2003: 301-318.

9.       Patil R K, Chimatadar SA and Nandibewoor T; Transition Met. Chem., 33; 2008:625.

10.     Kharzeoua S E and Serebrennikou V; Russ. J. Inorg. Chem., 12; 1967: 1601.

11.     Chinn LJ ,Selection of oxidants in synthesis, oxide. at carbon atoms; Marcel Dekker, New York.1971

12.     Augustin RL; Oxidation, I and II; Marcel Dekker, New York.1969.

13.     Evans WL; Chem. Rev., 6;1929: 281.

14.     Butterworth R F and Haneesian, S. Synthesis, 70;1971:121-124.

15.     Heyns  K and Paulsen H; Adv. Crabohydr. Chem., 17;1962:169-176.

16.     Guthine  RD; Adv. Crabohydr. Chem., 19;1962:109-116.

17.     Vogel A I.A text book of quantitative inorganic Analysis. Longmans, London,1961; 348.

18.     Laidler,K J; Chemical Kinetics, McGraw Hill, New York.1965

19.     Singh RB and Siddhartha SP ; Asian J. Exp. bio. Sci. 1(1); 2010: 204-207.

20.     Muller H, Pure Appl. Chem., 67(4); 1995: 601.

21.     Singh NK, Singh SK and Singh S, Kinetic study of Ru (III) catalysed oxidation of sucrose by sodium periodate in acidic medium, J. Chemtracs., 11(1);2009: 289-292.

22.     Singh AK, Chopra D, Rashmani S and Singh B; Carbohydrate Res., 314; 1998: 157-159.

 


 

 

 

 


 

 

Received on 09.08.2012        Modified on 20.08.2012

Accepted on 25.08.2012        © AJRC All right reserved

Asian J. Research Chem. 5(8): August, 2012; Page 1047-1052