CHAPTER – I

Introduction and

Review of Literature

1

OXIDATION OF SUGARS

Waters1a in a review of oxidation of organic compounds by

chromium and manganese derivatives explained by a close model

based on Oppenauer oxidation reversible enzymatic oxidations of

polyhydroxy derivatives like carbohydrates.

Littler1b in an important communication discussed the

oxidations of hydroxy compounds, olefins and other organic systems

by two electron oxidants like Cr(VI), Mn(VII), I(VII) and Pb(IV).

Oxidation of organic compounds by V(V) and halogens are also

subjected to analysis by Zimmerman treatment in the light of selection

rules for electrocyclic reactions. The conclusions are very interesting

because Cr(VI) easily oxidized alcohols but I(VII) is ineffective for

oxidation of alcohols. These facts and I(VII) and Pb(IV) being

extremely good for glycol oxidation were easily explained. Olefin

oxidation leading to cis hydroxylation by Mn(VII) was also

rationalized. This demonstrated that the selection rules are very

important in determining mechanisms of many organic reactions.

Bakore and others2 studied the oxidation of D-glucose in acid

medium. The reaction was first order each in D-glucose and oxidant.

The acid dependence is two. The results show that the chromic acid

oxidation of D-glucose is similar to that of secondary alcohols and

pinacol. The products are a mixture of formaldehyde, formic acid and

carbon dioxide.

Singh et al3 investigated the kinetics of oxidation of D-glucose,

D-galactose and D-fructose by alkaline ferricyanide. The rate of

2

oxidation is zero order on [Fe(CN)6]-3 and directly proportional to

[OH–] and to [sugar]. An ene-diol as an intermediate has been

suggested.

Singh and others4 studied the oxidation of D-xylose and

L-arabinose by Cu(II) in alkaline medium. The reaction is zero order

with respect to Cu(II) and first order each in [sugar] and [OH–].

Catalysis by Cu2O, formed as a heterogeneous product has been

observed.

Singh et al5 studied the oxidation of a few sugars by Ag(I) in

alkaline medium. It is postulated that silver ammonium hydroxide

complex is the oxidizing species and the products formed are the

ammonium salts of carboxylic acids.

Singh and Saxena6,7 have shown that oxidation of D-xylose

L-arabinose, D-glucose, D-fructose, D-mannose, D-galactose and

L-sorbose by Cu(II) is first order in [sugar] and [OH–]. The rate is

independent of [oxidant]. The enolisation rates have been found to be

the rates of oxidation in these studies.

Pottenger and Johnson8 investigated the oxidation of glucose

and cellulose in 1.0M perchloric acid by Ce(IV). A mechanism for the

oxidation of glucose has been postulated involving the formation of a

chelate complex with Ce(IV) and breakdown of this complex in the rate

determining step to form free radical. Subsequently the free radical

reacts with another mole of Ce(IV) to form products.

Mushran and others9 have studied the kinetics of oxidation of

xylose, arabionse and galactose by chloramine-T (CAT) in highly

3

alkaline medium. The reaction involves a brief initial induction period

after which the observed order with respect to [CAT] and [aldose] are

one each. The order in [alkali] is two. The authors suggested a rate

determining step between oxidant, alkali and substrate anion.

Kumar and Mehrotra10 have suggested a rate limiting

bimolecular reaction between vanadium (V) and the reducing sugars,

leading to the formation of a free radical. The free radical is rapidly

oxidized to products by V(V).

Bhatnagar and Fadnis11 investigated the oxidation of D-ribose

by V(V) in acid medium. The reaction is first order each in [sugar] and

[oxidant]. The dependence on [H+], at lower concentrations is one and

at higher concentrations is two. The same authors12 reported that in

the oxidation of D-xylose by Mn(III) pyrophosphate, the reaction is

first order with respect to Mn(III). The fractional dependence on

[xylose] indicates the rapid formation of reversible cyclic complex

between the sugar and Mn(III). The cyclic complex breaks down in

slow step to products.

Krupenskii and others13,14 in the series of investigations have

reported the oxidation of various reducing sugars by transition metal

ions. The products have been analysed both qualitatively and

quantitatively.

Pati and Mahapatro15 have investigated the kinetics and

mechanism of oxidation of D-arabinose, D-galactose, D-xylose,

D-mannose and D-glucose by phenyl iodosoacetate in aqueous acetic

acid -perchloric acid medium. The reaction is found to be total second

4

order, first order each in substrate and oxidant. The order with respect

to [H+] is inverse unity. A mechanism involving a transition state

complex between sugar and oxidant has been postulated which breaks

down in a slow step.

Gupta and others16 investigated oxidation of few aldoses by

Cr(VI) in acid medium. The reaction has been observed to be first

order on both [Cr(VI)] and [aldose]. The order with respect to [H+] is

complex,the major product of oxidation is reported to be aldonic acid

when the [aldose] is in excess.

Pati and Panda17 have studied the oxidation of some aldo and

keto sugars by V(V) in different binary solvent compositions. The

reaction is of second order, first order each in oxidant and substrate.

Keto sugars exhibited direct unit dependence and aldo sugars

exhibited direct fractional dependence on acidity. A mechanism routed

through a chelate complex between sugar and oxidant has been

suggested. A correlation has been made between the rate of oxidation

of sugars and the free aldehyde sugar concentration and it is suggested

that the free aldehyde sugar is participating in the reaction.

The same authors18 extended the oxidation of sugars to Ru(III)

catalysed by bromate in the presence of sulphuric acid and mercuric

acetate. They reported first order with respect to oxidant and zero

order with respect to substrate. The order on catalyst is unity. The

dependence on acidity is inverse: 1.6, 1.4 and 1.7 for D(+)- glucose,

D(+)-xylose and L(+) –arabinose respectively indicating the reactive

species to be the bromate ion. A mechanism involving slow formation

of outer complex between the oxidant and Ru(III) has been suggested

5

which disproportionates in a fast step to give Ru(V). The Ru(V) thus

formed, oxidizes the substrate in a fast step to give Ru(III) and

products.

In subsequent investigations the same authors carried out the

oxidation of sugars by peroxy disulphate catalysed by Ag(I) and

reported first order with respect to oxidant and zero order with

respect to substrate. The dependence on acidity is minimal and free

radical scavengers inhibited the rate of oxidation. A radical

mechanism has been suggested.

Singh and others19 studied the oxidation of maltose and

cellobiose by Nessler’s reagent (HgI ) in alkaline medium. In each case

the reaction proceeds after a slight induction period. The reaction rate

is first order with respect to reducing [sugar] and completely

independent of initial [Hg(II)]. It follows first order kinetics at lower

concentration of OH–, which becomes zero order at higher [OH–]. With

increase in iodide ion concentration the decrease in rate is noticed. A

general mechanism which involves intermediate enediol has been

proposed with HgI as the reacting species.

Fadnis and Kulshreshtha20 reported the oxidation of D-sorbitol

with vanadium (V) in sulphuric acid medium.

Varadarajan and others21 studied the D-glucose oxidation with

pyridinium fluorochromate at constant ionic strength and at different

acidities. The reaction is acid catalysed and at low temperature short

induction period is observed. They reported arabinose and formic acid

as the main products and the rate of oxidation of arabinose is

insignificant.

6

Kistayya and others22 studied the oxidation of aldoses by

N-bromosuccinimide (NBS) in aqueous acetic acid with ruthenium as a

catalyst. The reaction is first order in [NBS] both in the presence and

absence of catalyst. In the absence of catalyst it is first order in

substrate which becomes fractional order in the presence of catalyst.

The reactivity order is D-arabinose>D-xylose>D-galactose>D-mannose

>D-glucose.

Dhar23 studied the oxidation of D-glucose and its various C-1

and C-2 substituted derivatives by pyridinium chloro chromate (PCC)

and found the reactivity order as 2-deoxy-D-glucose >D-glucose>

1-O-methyl--D-glucopyranoside>2-amino-2-deoxy-D-glucose hydro-

chloride which is clearly explained by inductive, steric and shielding

effects.

A comparative study of and -anomers of some

monosaccharides revealed that -anomer is oxidized faster than

-anomer.

Singh and others24 investigated kinetics and mechanism of

oxidation of lactose and maltose by [Cu(Bipyridyl)2]2+ in alkaline

medium by spectrophotometric method. Kinetic data revealed that the

rates of oxidation are independent of [Cu(II)] and first order in [OH–]

and [sugar]. The studies of dielectric constant demonstrate that there

is a small increase in average zero order rate constant with the

decreasing dielectric constant of the medium and it is independent of

ionic strength of medium. A general mechanism involving the

intermediate enediol anion has been proposed.

7

Sharma and others25 reported the bromide catalysed oxidation

of dextrose by Ce(IV) in aqueous sulphuric acid solution. The reaction

shows first order dependence in dextrose and Ce(IV). Positive catalytic

effect on the reaction rate is shown by bromide ion. Reaction rate

decreases with increase in concentration of hydrogen ion and [HSO4–]

or [SO42-].

Sah26 studied bromide catalysed oxidation of fructose by Ce(IV)

in aqueous sulphuric acid solution. The reaction is first order in

fructose and Ce(IV).

Banerjee and others27 determined polyols by vanadium (V) in

perchloric acid. The oxidation of fructose on Pt/C catalyst has been

reported28. The major products are 2-keto-D-gluconic acid and

D-threo-hexo-2,5-diulose.

Singh and others29 studied the kinetics of Pd(II) catalysed

oxidation of D-arabinose, D-xylose and D-galactose by

N-bromosuccinimide (NBS) in acidic solution. The kinetic data shows

first order kinetics in each pentose and hexose at low concentrations

but at high concentrations change to zero order. First order in NBS

and Pd(II) and inverse fractional order in [H+] and [Cl-]. The

corresponding acids were identified as the main products.

Kinetics and mechanism of the oxidation of reducing sugars

by osmium tetroxide in alkaline medium was studied by Singh and

others 30. It shows pseudo unimolecular kinetics with respect to OsO4.

At lower concentration of sugar and OH– reaction shows first order

and changes to zero order at higher [sugar] and [OH–].

8

Gupta and others31 studied spectrophotometrically the kinetics

and mechanism of oxidation of some aldoses, amino sugars and

methylated sugars by tris (pyridine-2-carboxylato) manganese (III) in

sodium picolinate-picolinic acid buffer medium. The reaction is first

order in [manganese (III)] and [sugar].

Aldonic acid and Cr3+ are the final products in the oxidation of

D-glucose, D-allose, D-mannose, D-galactose, 6-deoxy-D-galactose and

2,6-dideoxy-D-ribo-hexose by [Cr(VI)]32.

Upadhyay and others33 investigated the kinetics and mechanism

of ruthenium (III) catalysed oxidation of arabinose, xylose, galactose,

glucose, fructose, lactose and maltose by chloramine-T. It shows first

order dependence on [substrate], [chloramine-T] and [OH–].

Dolezal and others investigated the oxidative degradation of

D-glucose using peroxo disulphate34. A number of products are

isolated like 3-hydroxy-2-pyranose, furan-2-carboxylic acid and furan

2-aldehyde as some of the products in addition to formic and acetic

acids. 4-oxo pentanoic acid also isolated as the minor product.

Gupta and others35 investigated the kinetics of oxidation of

some aldoses and amino sugars by potassium bromate in hydrochloric

acid medium. The reactions appear to proceed through the

intermediate formation of bromate esters followed by the

decomposition of esters to give products.

The kinetics of electron transfer reactions of aquo-thallium (III)

perchlorate with D-xylose, D-arabinose and D-ribose have been

studied under the pseudo first order conditions by Fadnis and others36.

9

The fractional order with respect to aldopentose concentrations

suggest the Michaelis-Menten type of kinetics.

The absorption kinetics of D-glucose at the alumina solution

interface was studied colorimetrically by Bajpai and others37. Various

kinetics and thermodynamic parameters of the absorption processes

have been evaluated.

The kinetics and mechanism of oxidation of D-glucose and

D-mannose with bromamine-T in alkaline medium was investigated by

Rangappa and others38. The rate of reaction is influenced by a change

in ionic strength of the medium and the dielectric effect is found to be

positive. The product analysis indicates that the sugars are oxidized to

a mixture of aldonic acid consisting of arabinonic, ribonic, erythronic

and glyceric acids.

The oxidation behavior and relative reactivity of methyl-

α-D-glucopyranoside and methyl-β-D-glucopyranoside towards

permanganate and acid chromate in perchloric acid medium was

studied by Parthasarathi39. The reactions are first order with respect

to [glucopyranoside] and [oxidant]. The reaction rate increases with

the increase in [H+].

Singh and others40 reported kinetics and mechanism of Ru(III)

and Hg(II) co-catalyzed oxidation of D-galactose and D-ribose by

N-bromoacetamide(NBA) in perchloric acid. The kinetic data indicates

first order in NBA at lower concentrations which changes to zero order

at higher concentrations, first order in [sugar], [Ru(III)], inverse

10

fractional order in [H+] and [acetamide]. Formic acid, D-lyxonic acid

and formic acid, L-erythronic acid are main products obtained in the

oxidation of D-galactose and D-ribose respectively.

Das and others41 studied the kinetics and mechanism of Ce(IV)

oxidation of D-mannitol and D-glucose in aqueous H2SO4 acid media in

the presence and absence of Ir(III). It is found that Ir(III) catalyses the

reaction even at very low concentration. It shows a first order

dependence in Ce(IV), [substrate] and [catalyst]. The reaction is acid

catalysed and inhibited by [HSO4-].

Fadnis and Sonali have studied42 the kinetics and mechanism of

electron transfer reactions of iron(III) perchlorate with D-glucose, D-

galactose and D-mannoses in presence of complexing 2,2|-bipyridyl

under the pseudo-first order conditions. The results suggests the

formation of 1:2 complex between Fe(III) and 2,2|-bipyridyl.

The kinetics of periodate oxidation of carbohydrates and

polymeric substrates have been reported by the method of isothermal

calorimetry43. The kinetic rates are dependent on the molecular weight

of the monomer and on its concentration. The order of reactivity is

trehalose > maltose > cellobiose.

Zaheer Khan et al44 reported D-fructose oxidation by

vanadium(V) in H2SO4 medium. This reaction shows an induction

period followed by auto acceleration, and the reaction is followed by

spectrophotometry by observing changes in absorbance at 350nm. It

shows first and fractional order in [V(V)] and [D-fructose] respectively,

however dependence on [H+] is complex.

11

Kinetics and mechanism of Ir(III) catalysed and Hg(II) co-

catalysed oxidation of reducing sugars by N-bromoacetamide in acidic

medium has been studied45, as a continuing study of their previous

investigation of Ru(III) and Hg(II) co-catalysed oxidation of sugars.

Isomerisation of a few monosaccharides has been

experimentally studied46 through electrophoresis. The carbohydrates

xylose and glucose isomerised under the lyocell conditions of

experimentation.

The kinetics of Pt(IV) catalysed chloramine-T oxidation of

glucose, galactose and fructose in alkaline medium was investigated47

by Upadhyay and Neelu Kambo. The reactions are first order in

oxidant, while the order of reaction in substrate and OH– decreases

from unity at higher concentration of substrate and [OH–] respectively.

A review work on oxidation of monosaccharides with

N-metallo-N-halo aryl sulfonamides was reported by Rangappa48.

Mechanisms of oxidation of monosaccharides such as erythrose and

threose series of pentoses and hexoses, 6-deoxyhexoses uranicacids

and amino sugars are studied with mild oxidizing agents such as Cl+,

Br+ or I+ in detail. The product profile was confirmed by HPLC and

GLC-MS Data.

Gowda et al49 studied the kinetics and mechanism of oxidation

of D-fructose and D-glucose by sodium salts of N-(chloro) mono /

di-substituted benzene sulfonamides in aqueous alkaline medium. The

kinetic orders are first order each on oxidant, substrate and alkali. The

authors proposed that the higher activation energy is responsible for

the lower reactivity of glucose compared to fructose. It is further

12

reported that the substituents in benzene ring of the aryl sulfonamides

affect the kinetic rate.

Kabir and others50 investigated the kinetics of oxidation of

L(+)Arabinose by Ce(IV) have been followed by monitoring the

disappearance of absorbance of Ce(IV) at 385nm in absence and

presence of surfactants. Whereas anionic micelles of sodium dodecyl

sulphate (SDS) have no effect on the oxidation rate, a two-fold increase

has been observed in the presence of cationic micelles of cetyl

trimethyl ammonium bromide (CTAB). The reaction obeys first order

kinetics with respect to [L(+)arabinose] in both the media. The

observed catalytic role has been analysed in terms of Menger-Portnoy

model.

In an earlier study from these laboratories51 the kinetics of

oxidation of maltose by dichloro isocyanuric acid (DCICA) at pHs 2, 4.5,

6.85 and 12. This reaction has zero order dependence on [oxidant] and

the relative orders on [Ru(III)] and on [sugar] are one each.

Further it was also reported by the same authors52 about the

study of the kinetics and mechanism of oxidation of D-ribose, D-

glucose and D-fructose by dichloroisocyanuric acid (DCICA) in aqueous

acetic acid – perchloric acid mixtures catalysed by Ru(III). The

corresponding lactones are the products in each case along with

formaldehyde in case of D-fructose under the conditions of [sugar] >

[DCICA].

In another communication the same authors reported53 the

kinetics of oxidation of D-glucose by 1,3,5-trichloro – 1,3,5-triazine –

2,4,6-trione under catalytic conditions using transition metals like

13

Ru(III), Os(VIII) and Mn(II) in aqueous acetic acid – perchloric acid

mixtures.

Rajput and others54 investigated the kinetics of oxidation of

aldoses by N-bromosuccinimide.

Goel and others55 investigated the kinetics of oxidation of some

aldoses like glucose, galactose, xylose and ribose by hexacyano

ferrate(III) ions in aqueous alkaline buffered medium. The kinetic

results indicate the zero order kinetics in hexacyanoferrate(III) and

first order in aldoses and OH–.

Kinetic data for the chromium(VI) – D-glucose redox system in

presence of complexing agents are reported for the first time by

Zaheer Khan and others56. The reaction is first order each in [Cr(VI)]

and [D-glucose]. The kinetics reveals complex order dependence with

[HClO4].

Rajput and Aarti studied57 the kinetics of oxidation of dextrose

under unctalaysed and catalysed conditions by N-bromo succinimide

in basic medium. The reaction has been found to be first order with

respect to each oxidant, substrate and OH– in both uncatalysed and

catalysed reactions. A first order dependence on catalyst is also

observed.

Sheila and others58 investigated the kinetics of Rh(III) catalysed

oxidation of dextrose and maltose by sodium periodate in acidic

medium. The reaction is carried out in the presence of mercuric

acetate as a scavenger for iodide ion. The rate shows first order

kinetics with respect to oxidant, sodium periodate and Rh(III) for both

dextrose and maltose.

14

Thus even though a lot of work was reported in literature on the

oxidation of sugars using transition metal catalysts very few reports

are noticed on the kinetic and mechanistic studies of oxidation of

carbohydrates using N-halo derivatives like dichloro dimethyl

hydantoin (DCDMH) under catalysed conditions.

15

References:

1. (a) W.A.Waters, Chem. Soc. Quart. Reviews, 12 (1958) 277 (Oxidations by Compounds of Chromium and Manganese). (b) J.S.Littler, Tetrahedron., 27(1) (1971) 81.

2. G.V.Bakore and M.P.Tondon, Z.Phys.Chem., (Leipzig), 222 (1963) 320.

3. M.P.Singh and N.Nath, J.Phy.Chem., 69 (1965) 2038.

4. S.V.Singh and M.P.Singh, Z.Phy.Chem. (Frankfurt am Main), 50 (1966) 11.

5. M.P.Singh, H.S.Singh, S.C.Tiwari, K.C.Gupta, A.K.Singh, V.P.Singh and R.K.Singh, Indian J.Chem., 13 (1975) 819.

6. S.V.Singh, M.P.Singh and O.C.Saxena, Indian J.Chem., 8 (1970) 529.

7. S.V.Singh, M.P.Singh and O.C.Saxena, J.Am. Chem.Soc., 92 (1970) 537.

8. C.R.Pottenger and D.C.Johnson, J.Polym.Sci., Part A-1, 8 (1970) 301.

9. S.P.Mushran, M.Sanehi and R.N.Mehrotra, Proc. Nat. Acad. Sci., 43A (1973) 105.

10. A.Kumar and R.N.Mehrotra, J.Org.Chem., 40 (1975) 1248.

11. A.G.Fadnis and R.K.Bhatnagar, J.Indian Chem. Soc., LIII (1976) 999.

12. A.G.Fadnis and R.K.Bhatnagar, Monatsh.Chem., 109 (1978) 329.

13. A.Y.A.Krupenskii and I.V.Dolgaya, Khim.Drev (Russ) 2 (1978) 81.

14. A.Y.A.Krupenskii and I.V.Dolgaya, Khim.Drev (Russ) 2 (1978) 84.

15. S.C.Pati and R.C.Mahapatro, Proc. Indian Acad. Sci. 88A (1979) 203.

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17. S.C.Pati and M.Panda, Int. J.Chem. Kinet., XI (1979) 73.

18. M.Panda – Ph.D. Thesis, Berahampur University (India) (1980).

19. M.P.Singh, R.K.Singh, A.K.Singh and Amita Srivastava, Indian J. Chem., 19A (1980) 547.

20. A.G.Fadnis and S.K.Kulshreshtha, J. Indian Chem. Soc., LVIII (1981) 763.

21. R.Vardarajan and R.K.Dhar, Indian J.Chem., 25A (1986) 474.

16

22. T.Kistayya, M.Surekha Reddy and Kandlikar Sushama, Indian J.Chem., 25A (1986) 905.

23. Raj K Dhar, Indian J.Chem., 31A (1992) 97. 24. A.K.Singh, Anita Parmar, Aparajita Tiwari, Amar Singh and Ranjana

Gupta, Proc. Indian Natn. Sci. Acad., 56A (1990) 71. 25. J.Sharma and M.P.Sah, J.Indian Chem. Soc., 71 (1994) 613. 26. Maheswar Prasad Sah, J.Indian Chem. Soc., 72 (1995) 173. 27. Amalendru Banerjee, Dinabandhu Mandal, Anuva Putatunda and

Gopal Chandra Banerjee, J.Indian Chem. Soc., 74 (1997) 667. 28. Heinen, Annenieke.W, Peters, Joop A and Van Bekkum Herman,

Carbohydr. Res., 304 (1997) 155. 29. Ashok Kumar Singh, Deepti Chopra, Shahla Rahmani, Bharat Singh,

Carbohydr. Res. 314 (1998) 157.

30. S.Hari Singh, Gupta Arti, K.Anil Singh and Singh Bihari, Transition Met. Chem., 23(3) (1998) 277.

31. Kalyan Kali Sen Gupta, Bilkis Ara Begum, Carbohydr. Res. 315 (1999) 70.

32. Signorella, Sandra; Daier, Veronica; Garcia, Silvia; Cargnello, Roxana; Gonzalez, Juan Carlos; Carbohydr. Res. 315 (1999) 14.

33. Neelu Kambo and Santosh. K.Upadhyay, Trans. Met. Chem., 25 (2000) 461.

34. M.Dolezal, O.Novotny, J.Velisek, Prague Czech Rep.Czech J. Food Sci, 18 (2000) 92.

35. K.K.Sen Gupta, N.Debnath, B.Nandini, B.Amalendu and B.Surendranath J.Indian Chem. Soc., 77 (2000) 152.

36. Anand G.Fadnis and Sapana Arzare, J.Indian Chem. Soc., 77 (2000) 235.

37. A.K.Bajpai and A.Choubey, J.Indian Chem. Soc., 77 (2000) 238.

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39. Parthasarathi Tribedi, J.Indian Chem. Soc., 78(2001) 287.

40. Ashok Kumar Singh, Vineeta Singh, Ajaya Kumar Singh, Neena Gupta and Bharat Singh, Carbohydr.Res., 337 (2002) 345.

41. Aparna Roy and Asim K Das, Indian J.Chem., 41(A) (2002) 2468.

17

42. Anand G.Fadnis and Sonali Rawat, J.Indian Chem. Soc., 80 (2003) 759.

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45. Ashok Kumar Singh, Shahla Rahmani, Bharat Singh, Ramesh Kumar and Manju Singh, J. Phy. Org. Chem., 17(3) (2004) 249.

46. Immanuel Adorjan, John Sjoberg, Thomas Rosenau, Andreas Hofinger and Paul Kosma, Carbohydr. Res., 339 (2004) 1899.

47. Neelu Kambo and Santosh K. Upadhyay, India J. Chem., 43A(2004) 1210

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18

OXIDATION OF KETONES

Lapworth1 studied the bromination of acetone in aqueous acidic

solutions. The reaction is dependent on [acetone], [H+] and

independent of [bromine]. In this reaction, the mechanistic process

involves an acid-catalysed enolisation.

Dawson and others2,3 investigated the aqueous iodination of

acetone catalysed by various acids and bases. The rate of reaction is

independent of [halogen]. Same results are obtained for alkali

catalysed chlorination and bromination of acetone. In case of base

catalysis enolisation itself was a composite process consisting of terms

due to rate controlling enolisation and due to rate controlling

halogenation.

Kinetics of oxidation of ketones by selenium dioxide was

investigated by Mel’nikov and Rokitskaya4. They established that the

rate of oxidation decreases gradually with increase of molecular

species. Aliphatic ketones are more easily oxidized than aromatic

ketones. All cyclic ketones are more rapidly oxidized compared to fatty

ketones.

The same authors5 extended the work to pyruvic and levulic

acids. The results indicate the correctness of mechanism postulated

earlier through enolisation.

Kinetics of oxidation of acetone by selenious acid was

invetigated by Duke6. The coordination of acetone through the oxygen

to selenium metal of selenium dioxide the results in redistribution of

charge with in acetone molecule causing enhance reactivity.

19

The halogenation of acetone in acid medium is investigated

by Bell et al7-9. This reaction is independent of the nature and

concentration of the halogen and dependent on initial concentration of

substrate and acid. The rate determining step involves an acid

catalysed enolisation and later kinetically very fast step(s) involving

the reaction(s) of the molecular enol with the halogen and proving to

the measured pseudo zero order kinetics on halogen.

In many cases where the rate of halogenation might dependent

on the [halogen], it was interpreted that by decreasing the [halogen]o

to a very low value (i.e. below some threshold concentration, ca 10-5M

or less) and high [H+], the rate of its attack on the enol could become

lower than the rate of enol formation that’s why attack by the halogen

becomes rate controlling and that the rate was proportional to the

concentration of halogen.

Eg: The rate of chlorination of acetone by molecular chlorine

representing to the enol-Cl2 reaction as follows.

S + H+ Enol + H+

Enol + Cl2 Products

Rate = = o

o

ClkHkSClHkk][][][]][[

221

212

Os(VIII), Ru(III)10,12,19,20,21,27 have extensively used for the

oxidation of ketones both in acid and alkaline medium by various

oxidants. The conclusions are first order in catalyst, zero order in

oxidant, first or fractional order in substrate and first or fractional

order in H+. Generally a mechanistic pathway of enolisation resulting

in enol formation in an equilibrium step followed by a complexation

k1

k2 k –1

20

between enol and catalyst is postulated. This complex breaks down in

a fast step without involving oxidant when there is zero order kinetics

on oxidant. Alternatively the complex breaks down in the presence of

oxidant resulting in first or fractional order in oxidant.

Hine et al11 and Guthrie23 studied the kinetics of transformation

of acetone-d6 to acetone-d5 and then to acetone-d4 in aqueous

solutions in the presence of H+, OH– and general buffers.

Mushran et al13,16,17,18,35 , Rao21 reported the oxidation of several

aliphatic, aryl aliphatic, cyclic ketones and acetyl pyridines by CAT in

aqueous and aqueous ethanol medium under alkaline conditions. In all

cases, the reaction shows first order on each of [substrate], [CAT] and

[OH–]. The corresponding 1,2-diketones are the products.

The different isotope effects14,15 on the various steps of

enolisation of ketone enolisation and ketonization are isolated by

kinetic measurements of the iodination and bromination of acetone

and acetone – d6, in the presence of H2SO4 or HClO4 using normal and

heavy water and under conditions such that enolization and halogen

addition to the enol occurred at comparable rates. For the ketonization

the H2O – D2O isotope effect (3.7) corresponds to the ratio of rate

constants for addition of H+ and D+. The same value is observed for the

hydrolysis of ethyl ether of acetone-enol. The different rate constants

of ketonization and enolization in different H2O – D2O mixtures are

also explained.

Singh and Anand24 studied the kinetics of the oxidation of some

cyclic ketones by selenium dioxide in 50% : 50% (v/v) acetic acid

21

water mixtures. A first order dependence in each selenium dioxide and

ketone has been observed. The rate of the oxidation reaction has been

found to be accelerated by the increase in the mineral acid

concentrations. The effects of solvent, solvent isotope, temperature on

the rate of oxidation have been studied in detail to determine

bimolecular oxidation. The values of energy of activation, the entropy

of activation and frequency factor for the reaction have been

evaluated.

Radhakrishnamurti and Mahapatro25 and Patnaik30 reported the

reactions of several enolisable ketones with N-iodosuccinimide and

N-bromoacetamide respectively. The reaction shows zero order on

[halogenating agent], first order each on [ketone] and [H+]. The results

were interpreted in terms of rate determining enol-formation

mechanism. The authors assumed that further rapid steps leading to

the product formation are kinetically immeasurable.

Dubois et al26 investigated kinetics of bromination and/or

iodination of cycloalkanones and aryl substituted acetophenones in

water, using very low concentrations of halogen. In addition to the

rate of enolization the rate limiting step involved the halogen addition

to the enol.

Sundaram et al29 studied kinetics of oxidation of aliphatic and

cyclic ketones in aqueous acetic acid by N-bromo saccharin in the

temperature range 10-500C. The reaction is found to be first order

with respect to [ketone] and [H+] and independent of [oxidant]. The

reaction is ion-dipole type and it does not initiate polymerization of

added acryl nitrile. Thermodynamic parameters like E, Hand S

22

are evaluated. Linearity of Exner’s plot and constancy in G

values suggest operation of a similar type of mechanism in all the

ketones. A plausible mechanism rationalizing the observed kinetic data

is proposed. The rate expression derived from the mechanism is of the

form. Rate=k[ketone]1[H+]1[NBSac]0.

Radhakrishnamurti and Rath31, Vasudevan and Venkata

Subramanian32 studied the kinetics of the reactions of several ketones

in the absence of added chloride ions in acidic solutions and the

reaction is found to be pseudo zero order. The same reaction shows

pseudo first order in the presence of added chloride36 and a linear

dependence each on [S] and [H+]. The dependence in the presence of

Cl– leads to a limiting value finally. These results are interpreted in

terms of probable mechanisms involving:

1. In the presence of Cl– a rate determining interaction of SH+ with the

most effective molecular chlorine species produced by the

hydrolysis of TCICA prior to the fast steps of product formation.

2. In the absence of Cl– a rate determining enol formation from the

conjugate acid of the ketones [SH+] is proposed.

Valchehha and Pradhan33 investigated the kineics of oxidation of

ethyl acetoacetate in aqueous acetic acid medium by selenium dioxide.

The reaction is found to be first order each in [substrate] and

[oxidant]. H+ catalysis is reported on the process. Primary salt effect is

negligible. Increase in % of acetic acid in reaction medium increases

the reaction rate. A suitable mechanism has been suggested.

23

Guthrie and Cossar34 investigated kinetics of chlorination of

simple aldehydes and of acetone. Cyclohexanone oxidation by

aquacerium(IV) ions in acid perchlorate is reported by Mehrotra28. It

shows first order dependence in [Ce(IV)], less than unity in

[cyclohexanone] and zero order dependence in [H+]. With increase in

ionic strength rate enhancement is observed. It is found that the

oxidation of cyclohexanone is faster than its rate of enolisation which

rules out the oxidation of ketone in the enol form. It is supported by

kinetic isotope effect, kH/kD = 2.0.

Kinetics of oxidation of cyclic ketones by 2,6-dichloro

indophenol is studied spectrophotometrically in presence of alkali by

Singh et al37. Results showing first order on [cyclic ketone] and second

order on [2,6 – dichloro indophenol] were reported. Rate of reaction

increases with [alkali]. It is observed that the rate does not affected by

neutral electrolytes.

Palo38 reported kinetic study of a few enolisable ketones with

trichloromelamine.

Dlingeleski and others39 investigated the kinetic study

on the acetate and phosphate ion-catalysed ketonization of

3-hydroxycyclohex-2-enone. At pH 4.2 – 5.1 the enol completely

ketonized to the unconjugated ketone i.e. 1,3-cyclohexadione.

Moorhoff and Paquette40 developed an ingenious method for the

conversion of cyclic ketones to oximino esters. In this process there is

simultaneous conversion of adjoining carbonyl and methylene groups

into two differently oxidized terminal carbon atoms. This has been

extended to a number of steroidal ketones for leaving the five

24

membered ring by the action of ethyl nitrite in tetra hydro furan

medium at low temperature.

Rothenberg and others41 investigated the oxidative cleavage of

cycloalkanones using sodium hypochlorite under phase transfer

catalysis conditions. The corresponding dicarboxylic acids are

reported to be the products.

Oxidative ring contraction of cycloalkanones is done by Giurg

and others42 to synthesize cycloalkane carboxylic acids using 30%

H2O2 in the presence of poly (bis-anthra ethyl) diselenide I as catalyst.

Panigrahi and Swain have investigated43 the kinetics and

mechanism of oxidation of 2-hydroxy cyclohexanone by V(V) in

aqueous acetic acid – perchloric acid mixtures. The reaction has been

found to be first order each with respect to [V(V)] and [H+]. The

dependence on [substrate] is two.

Radhakrishnamurti et al44 studied the kinetics of Mn+2 catalysed

oxidation of cyclic ketones by lead tetra acetate. Kinetics of oxidation

of a few cyclic ketones by lead tetra acetate in acetic acid medium with

perchloric acid catalysed by Mn+2 have been investigated. Reactions

are uniformly zero order in oxidant.

Radhakrishnamurti et al45 studied the kinetics of oxidation of

cyclic ketones by 1,3-dichloro-5,5-dimethyl hydantoin(DCDMH). The

results were interpreted as a dualistic pathway involving the oxidation

of keto form and the enolic forms simultaneously based upon

kinetic observations. The order of reactivity is

cyclohexanone> cyclpentanone>cyclooctanone> cycloheptanone.

25

Oxidation of Camphor:

Evans and Others46 investigated the oxidation of dl-Camphor by

SeO2. The product was found to dl-Camphor quinone.

Shiner and Others47 investigated the kinetics and mechanism of

oxidation of camphor quinone by periodate. The reaction rate is first

order in periodate and first order in -diketone over a fairly wide

range of concentration and the rate of attack of the periodate increases

as the degree of ionization increases.

Ronald studied48 the oxidation of camphor with peracetic acid

followed different courses depending on the acidity. In weakly acidic

solution corresponding lactone was formed in high yield.

Siyuye and others49 investigated the anodic oxidation of

norcamphor in aqueous electrolyte solutions. Norcamphor was

anodically oxidized at Pb/PbO2 anodes in 1M H2SO4, CH3CN / H2O

(v/v = 1/1). 3-oxo cyclopentane acetic acid and oxa bicyclo [3.3.0]

octan-3-one were obtained with yields upto 76% and 42%

respectively.

Sarah and others50 studied the microwave assisted oxidation by

SeO2 of camphor derivatives leading to -dicarbonyl compounds and

oxi imines. Compared to the classical reaction conditions, good yields

were obtained in much shorter reaction times.

Oxidation of deoxybenzoin:

Corey and Schaefer51 studied the mechanism of oxidation of

ketones by selenium dioxide. The oxidation of desoxy benzoin by

selenious acid in 70% acetic acid is catalysed both by acids and by the

26

base acetate ion. The rate expression for the acid catalysed reaction is

-d[SeO2]/dt=k[SeO2][desoxybenzoin][H+] and that for the acetate

catalysed process seems to be best represented as -d[SeO2]/dt=

k[SeO2][desoxybenzoin][OAc-]. For the acid catalysed reaction the

effect of a p-substituent on either ring of the ketone on rate is virtually

identical with that for acid catalysed enolisation, indicating that the

two reactions are of similar type. The presence of two ortho

substituents on the ring adjacent to the carbonyl group does not

depress the rate of either the acid or acetate catalysed oxidation, which

eliminates the possibility that carbonyl addition is involved.

Ogata and others52 investigated the kinetics of oxidation of

deoxybenzoin by nitric acid. Oxidation of deoxybenzoin with 1M HNO3

at 600C in 40% dioxin gives benzoin (25%), benzil (12%) and

degradation products (benzoic acid and para nitro benzoic acids and

benzaldehyde) (62%), while in 70% acetic acid benzoin (46%), benzil

(11%) and the degradation products (41%) are produced. Benzil and

benzoic acids are probably formed directly from deoxybenzoin, since

both benzoin and benzil are stable under these conditions.

Wiberg and others53 investigated the kinetics of the chromic

acid oxidation of deoxybenzoin in 91% acetic acid. The rate law is

given by

휐 =푘 푘 [푘푒푡표푛푒][퐶푟(푉퐼)][퐻 ]

푘 [퐻 ] + 푘 [퐶푟(푉퐼)]

Where k1 was found to be equal to the rate of enolisation. Benzoin was

shown to be the intermediate in the reaction, and the source of the

products: benzil, benzaldehyde and benzoic acid.

27

Khandual and Nayak studied54 the kinetics of oxidation of

deoxybenzoin by chromic acid in aqueous acetic acid medium. The

reaction rate is first order with respect to the oxidant as well as to the

organic substrate. The reaction rate decreases in the presence of

added Mn(II) ions. The presence of complexing agents like succinic

acid, piperidine etc decreases the rate. The thermodynamic

parameters for the oxidation have been computed. A mechanism

proceeding through an enol intermediate has been suggested.

Khandual55 reported the kinetics of OsO4 catalysed oxidation of

deoxybenzoin and p-nitro deoxybenzoin by alkaline hexa cyano

ferrate(III) in 30% t-butanol-water mixture at constant ionic strength.

The reaction is found to be first order each in [substrate], [Os(VIII)]

and [OH–], ([OH–]<0.04M) but independent of [hexacyano ferrate(III)].

The rates of reaction decrease with decrease in dielectric constant and

increase with increase in ionic strength of the medium. The entropy of

activation is found to be negative. The mechanism involves the

formation of an intermediate complex between enolate anion of

deoxybenzoin and Os(VIII) which rapidly decomposes followed by a

fast reaction between the reduced osmium species and hexacyano

ferrate(III).

Based on the survey it is felt worthwhile to investigate the

kinetics of oxidation of ketones in general with particular reference to

oxidation of camphor and deoxybenzoin by N-halo compounds like

trichloro isocyanuric acid (TCICA), chloramine-T (CAT) and metallic

oxidants like Ce(IV), V(V) and Mn(VII) in aqueous acetic acid-

perchloric acid mixtures.

28

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2. H.M.Dawson and J.S.Carter, J.Chem. Soc. (1926) 2282.

3. H.M.Dawson and N.C.Dean, J.Chem. Soc. (1926) 2872.

4. N.N.Mel’nikov and M.S.Rakitskaya,

J.Gen.chem(U.S.S.R)(1938)1369:Chem.Abs.,33(1933)4194

5. N.N.Mel’nikov and M.S.Rakitskaya,

J.Gen.chem(U.S.S.R)(1945)657:Chem.Abs.,40(1946)5702

6. Frederick R.Duke, J.Am.Chem.Soc., 70(1948)419

7. R.P.Bell and M.Spiro, J.Chem. Soc., (1953) 429.

8. R.P.Bell and K.Yates, J.Chem. Soc., (1962) 2285.

9. R.P.Bell and Dewis G.G., J.Chem. Soc., (1964) 902.

10. M.P.Singh, S.V.Singh and O.C.Saxena,

J.Am. Chem. Soc., 91 (1969) 2643.

11. J.Hine, J.C.Kaufmann and M.S.Cholod,

J.Am Chem. Soc., 94 (1972) 4590.

12. S.P.Mushran, R.Sanehi and A.K.Agarwal,

Z.Natur Forsch (B), 27(10)(1972) 1161.

13. S.P.Mushran, R.Sanehi and A.K.Bose,

J. Indian Chem. Soc., 1 (1973) 197.

14. J.Toullec and J.E.Dubois, Tetrahedron, 29 (1973) 2851.

15. J.Toullec and J.E.Dubois, J.Am. Chem. Soc. 96 (1974) 3524.

16. S.P.Mushran, A.Sharma and A.K.Bose,

Bulletin de la Acad. Polon. Des. Sci. Chem. (1974) 889.

17. S.P.Mushran, A.Sharma and A.K.Bose,

Ann. Soc. Sci., Brusselles Ser., 189(11) (1975) 567.

18. S.P.Mushran, R.Sanehi and A.K.Bose, Acta Chim. Acad. Sci.

Hungaricae, 84(2) (1975) 135.

19. R.I.Anantaraman and C.G.Nair, Indian J.Chem., 14A (1976) 45.

29

20. G.P.Panigrahi and P.K.Misro, Indian J. Chem., 16A (1978) 201.

21. T.Navaneeth Rao, M.Rajanna and P.K.Sai Prakash,

Indian J.Chem.,17A (1979) 297.

22. P.S.Radhakrishnamurti and D.K.Mahapatro,

Indian J.Chem. 18A (1979) 53.

23. J.P.Guthrie, Can. J. Chem, 57 (1979) 240.

24. K.J.Singh and S.N.Anand, J.Indian Chem.Soc.,LVI(1979)363

25. P.S.Radhakrishnamurti, D.K.Mahapatro,

Indian J.Chem., 19A (1980) 207.

26. J.E.Dubois, M.E.Alaoni and J.Toullec,

J.Am. Chem. Soc., 94 (1981) 103.

27. V.B.Agarwal, S.P.L.Agarwal and V.B.Agarwal. Chem. Abs. (1982) 65.

28. Raj Naraian Mehrotra, Indian J. Chem., 22A (1983) 169.

29. K.Vijaya Mohan,P.Raghunatha Rao and E.V.Sundaram

J.Indian.Chem.Soc.,LXI(1984)225.

30. D.P.Patnaik, Ph.D. Thesis, Berhampur University (India) (1985).

31. P.S.Radhakrishnamurti and N.K.Rath,

Indian J. Chem., 24A (1985) 300.

32. K.S.Vasudevan and N.Venkatasubramanian,

Indian J. Chem., 24A (1985) 304.

33. N.D.Valechha and Asha Pradhan, Indian J.Chem,24A(1985)773

34. J.P.Guthrie and J.Cossar, Can. J. Chem. 64 (1986) 2470.

35. S.P.Mushran, R.Sanehi and A.K.Agarwal,

Z.Phy. Chem., Leipzig, 255 (1987) 293.

36. P.S.Radhakrishnamurti, N.K.Rath and R.K.Panda,

J. Chem. Soc., Perkin Trans.II (1987) 517.

37. H.S.Singh, B.Singh, A.K. Jaiswal and P.KTandon,

J.Indian Chem. Soc., 67 (1990) 625.

38. L.N.Palo, Ph.D. Thesis, Berhampur University (India) (1990)

30

39. G.D.Dlingeleski, G.Blotry and R.M.Pollock,

J.Org. Chem., 55 (1990) 1019.

40. Cornelis M. Moorhoff and Leo A. Paquette,

J.Org. Chem., 56 (1991) 703.

41. Rothenburg, Gadi, Sasson, Yoel, Tetrahedron, 52 (1996) 1364.

42. Giurg, Miroslaw and Mlochowski, Jacek,

Synth. Commun., 29(13) (1999) 2281.

43. G.P.Panigrahi and Rajendra Swain, Indian J.Chem., 39A (2000) 775.

44. A.K.Sambasiva Rao,B.SyamaSundar and P.S.Radhakrishnamurti

J.Indain.Chem.Soc.,82(2005)310.

45. N.Madhavi,B.SyamaSundar and P.S.Radhakrishnamurti,

Oxid.Commun.,29(2)(2006)304

46. W.C.Evans, J.M.Ridgion and J.L.Simonsen., J.Chem. Soc., (1934) 137.

47. V.J.Shiner and C.R.Wasmuth., J. Am. Chem. Soc., 81 (1959) 37.

48. Ronald R. Sauers., J. Am. Chem. Soc., 81 (1959) 925.

49. Siyu Ye and F.Beck., Journal fur praktische chemie/

chemiker-zeitung., 334(1) (1991) 37.

50. Belsey Sarah, Danks Timothy and Wagner Gabriela,

Synthetic communications., 36(8) (2006) 1019.

51. J.E.Corey and J.P.Schaefer, J.Am.Chem.Soc.,82(1960) 918.

52. Y.Ogata and H.Tezuka, Tetrahedron., 25 (1969) 4797.

53. Kenneth B.Wiberg, Orm Aniline and Arnold Gatzke,

J. Org. Chem., 37(21) (1972) 3229.

54. P.L.Nayak and N.C.Khandual, Proc. Indian Acad. Sci., 1 (1974) 33.

55. N.C.Khandual, J.Indian Chem. Soc., 67 (1990) 621.

31

OXIDATION OF -DIKETONES

Shiner et al1 studied the oxidation of -diketones with periodate

in the pH range of (11-13) spectrophotometrically at 250 nm and

correlated with the changes in concentration of various ionized species

of periodate and separate rates of reaction of each of the periodate

species with -diketones were computed. The results were

interpreted in terms of nucleophilic attack of each of the six

coordinated periodate species on carbonyl carbon atoms of -

diketones to form a transient intermediate which undergoes

spontaneous decomposition to the products, iodate and two molecules

of carboxylic acid. The reaction rate is first order in periodate and first

order in -diketone over a fairly wide range of concentration and the

rate of attack of periodate increased as the degree of ionization

increased.

Peroxy acid oxidation of -diketones was studied by Panda

et al2. They postulated a Baeyer - Villiger type of oxidation of

-diketones by peroxy mono phosphoric acid and peroxy mono

sulphuric acid at different pH ranges. The reactions were second order

i.e. first order in peroxy acid and first order in -diketone. The

oxidation rates were strongly pH dependent, the rate increases with

increase in pH. From the pH rate data the reactivity of different peroxo

species in these oxidations were reported. A mechanism consistent

with rate determining nucleophilic attack of peroxo species on

carbonyl carbon of diketone molecule has been proposed. Acetic acid

and benzoic acids are respectively found to be the products of

oxidation of biacetyl and benzil.

32

Leffler3 reported the oxidation of -dicarbonyl compounds by

hydrogenperoxide. A cyclic mechanism was proposed for the oxidation

of these substrates based upon the migration of R–CO– group through

anionic mechanism. Furuya and Urasaki4 reported the oxidation of

benzil by peroxy acetic acid in sulphuric acid - acetic acid medium.

The rate of the consumption of peroxy acetic acid found to be the sum

of four terms shown in the following equation.

H][AcO][Bz]SO[Hk kH]][AcOSO[Hk{kdt

H]d[AcO2

242

122242td

2

It was considered that the first and second terms corresponds to

decomposition of peroxy acetic acid to acetyl peroxide, while the third

and fourth terms represent the uncatalysed and acid catalysed

oxidation of benzil with peroxy acetic acid respectively. A mechanism

passing through benzoic anhydride was postulated.

The reactivity of -diketones in particular 9,10-phenanthro

quinone, 1,2-naphtho quinone, acenaphthoquinone, benzil, 4,4|-

dimethyl benzil, 2,4,6-trimethyl benzil, 2,5,2|,5|-tetramethyl benzil,

2,3,5,6-tetramethyl benzil, 2,4,6,2|,4|,6|-hexamethyl benzil, phenyl

glyoxal was investigated in hydrogen donating solvents by Maruyama

and others5.

Sawaki et al6 studied the acyclic cleavage of benzils by alkaline

hydrogen peroxide. The reaction of benzil with alkaline hydrogen

peroxide in aqueous methanol gives benzoic acid and methyl benzoate,

and the yield of ester increases to 69% when H2O2 is added gradually.

The reaction in the presence of dimethyl sulfoxide gives the maximum

yield (81%) close to (89%) from the reaction of benzoic anhydride

33

under the same conditions but in the absence of H2O2. These results

rule out a dioxetane mechanism as the main pathway for the reaction

of benzil but suggest that an anhydride is intermediate.

The reaction of benzil with 18O labeled H2SO4 showed that one

oxygen atom is incorporated in the carboxylic acid product but none in

the ester. These results clearly eliminate epoxide mechanism but

supports Baeyer – Villiger type acyclic mechanism involving the

migration of an acyl group with formation of an intermediate

anhydride. A small portion of the anhydride appears to react with H2O2

to form peroxy acid which in turn oxidises more benzil in a path which

does not lead to ester.

Recently Mukherjee and others7 discussed the role of

conformeric triplets in the formation of hydrogen adduct radicals and

radical anions of benzil and 2,2|-dichloro benzil.

Anunziata and others8 studied the gas phase oxidation of methyl

ethyl ketone on V-ZSM-5 zeolite in the presence of molecular oxygen.

Two types of competitive partial oxidations, i.e., diacetyl formation and

oxidative scission reaction leading to acetaldehyde and acetic acid.

In a recent study from these laboratories9 kinetics of oxidation

of benzil in alkaline medium by periodate in tertiary butanol water

mixtures was investigated. The reaction was found to have varying

orders with the oxidant and benzil at different concentrations. The

mechanism suggested was complex formation between benzil

monohydrate anion and active periodate species. The rate law of the

following form explains the kinetic observations.

34

Rate = [Benzil]}K {1 ]}[IOK{1

][IO ][OH [Benzil] .Kk.K

142

-4T

-21

Benzil has been also subjected to oxidation by periodate with

Ru(III) as catalyst in aqueous acetic acid-perchloric acid mixtures. The

mechanism involves the complexation of catalyst and benzil and the

complexation breaks down through C-C cleavage leading to products.

The suggested rate law was

Rate = [ ( )][ ][ ( )]

Duvaux and others10 reported the oxidation of vicinal diketones

at room temperature to carboxylic acids by Mn(III) fluoride with 90%

yield. Investigation of the kinetics suggested that the reaction

proceeded via a cyclic intermediate. Franklin and others11 investigated the ion pair catalysis of the

auto oxidation of pyridine and benzil in emulsion system. It has been

found that the rates of absorption of oxygen by pyridine – aqueous

sodium hydroxide emulsions and the same emulsions containing

benzil were catalysed by the addition of quaternary salts and followed

the same rate law.

−d Pdt

= k P [OH ][benzil] + k P [OH ][Et NCl][benzil] Recently from our laboratories12 kinetics of oxidation of benzil

by Ce(IV) in aqueous acetic acid – perchloric acid mixtures was

reported. The reaction has been found to be first order in [oxidant],

first order in [substrate] and first order in [H+].

Hence a thorough investigation has been undertaken for the

oxidation of substrates like diacetyl and benzil with Ce(IV), V(V),

TCICA, CAT, KMnO4 as oxidants in order to establish the relative

oxidizing power of these oxidants.

35

References:

1. V.J.Shiner and C.R.Wasmuth, J.Am.Chem.Soc., 81 (1959) 37.

2. R.Panda, A.K.Panigrahi, C.Patnaik, S.K.Sahu and S.K.Mahapatra, Bull. Chem.Soc., Japan., 61 (1988) 1363.

3. J.E.Leffler, J.Org.Chem., 16 (1951), 1785.

4. Y.Furuya and I.Urasaki, Bull. Chem.Soc. Japan., 41 (1968), 660.

5. Kazuhiro Maruyama, Keiichi Ono and Jiro Osugi, Bull. Chem. Soc (Japan)., 45 (1972) 847.

6. Y.Sawaki and C.S.Foote , J.Am. Chem.Soc., 100 (1979) 6292.

7. Jaya Mukherjee, Dipankar Sen and Subhash Chandra Bera, Journal of Chemical Sciences., 104 (1992) 693.

8. Oscar A. Anunziata, Liliana B., Pierella and Andrea R. Beltramone, Catalysis letters, 75(12) (2001) 87.

9. Sham Rao K.Dhannure, Ph.D.Thesis, Acharya Nagarjuna University (2003).

10. Duvax Doroth’Ee, Ho Jason and Hughes D.E.Peter Journal of Chemical Research., 10 (2003) 635.

11. Thomas C.Franklin and Naoyuki Ogiya., International Journal of Chemical Kinetics., 12 (2004) 1055.

12. T.L.M.V.Subba Rao, B.Syama Sundar and P.S.Radhakrishnamurti, J. Inst. Chemists (India)., 80(5) (2008) 158.

36

OXIDATION OF SULPHUR CONTAINING -AMINO ACIDS

The earliest investigation of the oxidation of cystine is that of

Gerrit and Theodore1. Perbenzoic acid oxidation of cystine to cystine

dioxide which has disulphoxide structure was suggested. The

conformation was also reported on the basis of analysis and physical

properties like optical rotation.

Subsequently a number of publications have come on the role of

sulphide peptide functions in free radical transfer2.

Stephen and others3 studied the kinetics of oxidation of Cu(I)

complexes of cystine and penicillamine. The Cu(I).L2 complex with

cystine ligands of total Cu(I) concentrations of 10-30 µM was shown to

be oxidized by cystinyl radicals (RS) with a diffusion controlled rate

constants ka = 1.8 x 109Ms-1. The corresponding reaction with the

cystine disulphide anion (RS – SR) proceeded of a slower rate

k11b = 2.7 x 108Ms-1.

Reactivity of ferrate (VI) and ferrate (V) with sulphur containing

-amino acids like cysteine and cystine was investigated by Bielski and

others4. The results presented in this paper demonstrated that the

oxidation of organic and inorganic compounds by Fe(VI) are in

principle significantly accelerated by addition of an effective reducing

agent which converts Ferrate (VI) to ferrate(V), as Fe(V) is shown to

react 3-5 orders of magnitude faster than Fe(VI). Both hypervalant ion

species (Fe(VI)/Fe(V)) react preferentially with the protonated forms

of amino acids, and in the absence of dioxygen, the oxidation of amino

acids goes by chain reactions.

37

The kinetic study of aqua chromium (III) anation by L-cysteine,

a sulphur containing -amino acid was investigated by Kabir-ud-Din

and others5. All the kinetic measurements were made colorometrically

at a wave length of 545nm under pseudo first order conditions of

[cysteine]T > 10[Cr(III)]T.

Chansoria and Mishra6 reported the kinetics of Cu(II) catalysed

oxidation of cysteine hydrochloride under anaerobic conditions and in

presence of hydrochloric acid. The stoichiometry was found to be 2:1

and cystine is the oxidation product.

Cysteine kinetics and oxidation at different intakes of

methionine and cysteine in young adults were studied by Raguso and

others7. The studies supported the use of [1-13C] cysteine for studying

whole-body sulphur amino acids (SAA) oxidation and conclusions that

maintenance of SAA balance is best achieved by supplying methionine

at approximately the FAO/WHO/UNV recommendations for total SAA

intake (13mg kg-1d-1).

Raguso and others8 investigated the effect of cystine intake on

methionine. Kinetics and oxidation determined with oral traces of

methionine and cystine in healthy adults.

An elaborate study was made by Read and others9 to establish

the kinetics and mechanism of oxidation of a few cysteine and cystine

derivatives by potassium ferrate.

Grossi and Montevecchi10 found S-nitroso cysteine and cystine

from the reactions of cystine with nitrous acid during their kinetic

investigations in the pH range 0.5 – 7.0.

38

Steven and others11 studied the reactivity and oxidation

pathway of cysteine 232 in recombinant human 1-antitrypsin. The

oxidation pathway begins with a stable sulphenic acid (CySOH)

intermediate followed by the formation of sulphinic acid (CySO2H) and

cysteic acid (CySO3H) in successive steps.

Herszage et al12 used Hydrogen peroxide (H2O2) for the kinetics

of oxidation of cysteine (CySH) in aqueous buffers over a wide range of

pH (pH 4-13). By varying ratios of initial reactant concentrations to

explore the range of conditions they reported that a two step

nucleophilic model may function.

Luo et al13 reported the oxidation of cysteine by soluble

polymeric MnO2. The kinetics of reduction of soluble polymeric MnO2

by cysteine and glutathione was studied in the pH range of 4.0 – 9.0.

The concentration of thiols was varied between 1 and 2mM, with the

MnO2 concentration was varied between 2 and 12µM. In this pH range,

the reaction products were identified as Mn(II) and the corresponding

disulphides (cysteine sulphonic acid was formed only at pH<2).

Compton and others investigated14 the kinetics of oxidation of

cysteine by aqueous ferricyanide using boron doped electrode

voltammetry.

Vani and others15 studied the kinetics of oxidation of L-cysteine

in aqueous HClO4 medium using a one-equivalent oxidant,

hexachloro iridate(IV). The reaction exhibits second order dependence

with respect to oxidant and first order in cysteine. The rate decreases

with increase in H+ concentration indicating that the zwitterion form

of cysteine is more reactive. Cysteic acid was identified as the product

39

of oxidation. A suitable mechanism involving the formation of [IrCl6]2- -

sulphur bonded intermediate was proposed.

The kinetics of oxidation of cysteine(CySH) by 0.001M

hexacyano ferrate(III) ([Fe(CN)6]3-) in 0.06M HCl at 250C in aqueous

and / or in and 0.005M sodium dioctyl sulfosuccinate (AOT) was

reported by Hisham16 using optical absorption at 420nm. The reaction

was found to be first order in both [oxidant] and [substrate]. The rate

of the reaction in the presence of micelles was explained using a

pseudo-phase model of the kinetics.

The anaerobic oxidation of cysteine by Mn(III) in aqueous acetic

acid solutions was investigated by Salamon and others17 using stopped

flow spectrophotometric method at 200C. The formation and

disappearance of the [Mn(OAc)2Cys]- complex was monitored at

350nm. The rate depends strongly on the acetic acid concentrations.

The product was found to be cystine.

Fontana and others18 studied the oxidation of hypotaurine and

cysteine sulphinic acid by peroxy nitrite, producing the respective

sulphonates.

Stanbury and Hung have investigated19 the kinetics of oxidation

of cysteine by octacyano-molybdate(V) ([MO(CN)8]3-) in presence of

catalyst Cu2+. Cystine and cysteine sulphinate are the predominant

cysteine oxidation products.

The kinetics and mechanism of oxidation of L-cysteine by

Corey’s reagent (pyridiniumchlorochromate (PCC)) was investigated

by Vani and others20. The reaction exhibits first order dependence

40

with respect to PCC and fractional order in cysteine. The increase in

the oxidation rate with acidity suggests the involvement of a

protonated chromium (VI) species in the rate determining step. Cysteic

acid is identified as the product of oxidation.

Abdel-Halim and others21 reported the kinetics of the oxidation

of L-cysteine by trans-and cis-cobalt(III) and Iron(III) complexes in

aqueous solution. Kinetic measurements were run under pseudo first

order conditions.

The kinetics and mechanism of oxidation of cysteine with

hypohalous acid was reported22 by Nagy and Michael in aqueous

alkaline medium using stopped flow spectrophotometry. Two reaction

pathways are observed. The products are found to be cysteine

sulphinic acid and cystine.

Zahdeh Rana and others23 investigated the kinetics of oxidation

of cysteine and captopril via octacyanomolybdate(V) and

octacyanotungstate(V) in a buffered acid media using spectro-

photometry. The rate law for the oxidation is = k[RSH][Ox][H+]-1. The

results indicate that Cs3[MO(CN)8] is more reactive than Cs3[W(CN)8]

as an oxidizing agent. The products are corresponding disulphides.

Kinetics of oxidation of cysteine, cystine and cystine sulphoxides

by I in aqueous and aqueous acetic acid – perchloric acid mixtures

was recently reported from these laboratories. The Frost- Schwemer

treatment has been applied to evaluate the rate constants of

consecutive steps.

41

Vani and others25 studied the kinetics of oxidation of L-cystine

by Ce(IV) in sulphuric acid medium. The reaction exhibited first order

dependence with respect to [Ce(IV)] and [Cystine]. Ionic strength has

negligible effect on the rate. The rate of the reaction decreased with

increase in H+ concentration upto 0.5 – 1.0M and remains constant

thereafter. Cysteic acid was the main product.

Vani and others26 investigated the kinetics of oxidation of

L-cystine by pyridinium bromo chromate (PBC) using spectro

photometer. It was found that the reaction is first order in [PBC] and

fractional order in [cystine]. The reaction rate increases with increase

of [H+].

Mohanty and others27 studied the kinetics of oxidation of

L-cystine by 12-tungsto cobaltate(III) in aqueous perchorate medium

using spectrophotometer. The reaction showed first order dependence

on both [CO(III)W5-] and [L-cystine]T. The product of oxidation was

found to be cystine mono sulphoxide.

Vani and others28 investigated the kinetics of oxidation of

L-cystine by hexacyano ferrate (III) in alkaline medium using spectro

photometer at amax at 420nm. The reaction was found to be first

order dependent each on [oxidant] and [cystine]. The oxidation

product of the reaction was found to be cysteic acid.

In view of the above survey it was felt appropriate to investigate

the oxidation of cystine with iodine in acid medium to establish the

consecutive nature of the reaction and thereby to evaluate the rate

constants of consecutive steps by using Frost-Schwemer treatment.

42

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