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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.
16. K.K.Sen Gupta and S.N.Basu, Carbohydr. Res, 72 (1979) 139.
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.
38. P.A.Prashanth, K.Mantelingu, A.S.Anandamurthy, N.Anitha Rangaswamy and K.S.Rangappa, J.Indian Chem. Soc., 78 (2001) 241.
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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44. Zaheer Khan, P.S.S.Babu and Kabir-ud-Din, Carbohydr.Res., 338 (2003) 133.
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,
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27. V.B.Agarwal, S.P.L.Agarwal and V.B.Agarwal. Chem. Abs. (1982) 65.
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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.
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30
39. G.D.Dlingeleski, G.Blotry and R.M.Pollock,
J.Org. Chem., 55 (1990) 1019.
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42. Giurg, Miroslaw and Mlochowski, Jacek,
Synth. Commun., 29(13) (1999) 2281.
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44. A.K.Sambasiva Rao,B.SyamaSundar and P.S.Radhakrishnamurti
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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.
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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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